Sedimentary Sulfides

Chapter 14

Sedimentary Sulfides

Chapter Outline 4.5. Distribution of S8 in 1. Background 544 2. Sulfide Phases in Sediments 546 Sediments 2.1. Acid Volatile Sulfide 547 5. Dissolved Species in 3. Analytical Chemistry of Sedimentary Sulfide Systems Sedimentary Sulfide 549 5.1. Total Dissolved Sulfide in 3.1. Analytical Precision and Sediment Porewaters Accuracy 549 5.2. Dissolved Sulfide in 3.2. Total Sulfur and Total Anoxic Sediments Reduced Sulfide 551 5.3. Dissolved Sulfide in 3.3. Analysis of Organic Sulfur 552 Suboxic Sedimentary 3.4. Analysis of AVS 554 Systems 3.5. Analysis of Porewater 5.4. Dissolved Fe Species Sulfide 556 5.5. Aqueous FeS Clusters 4. Solid Sulfide Phases in Sulfidic 5.6. Dissolved Organic Sediments 559 Complexes of Fe and S 4.1. Pyrite Distribution in 5.7. Dissolved Polysulfides Sediments 561 5.8. Sulfur Oxyanions 4.2. Organic Sulfur 567 5.9. Dissolved S8 4.3. The Origins of 6. Diagenetic Modeling of “Anomalous” Sedimentary Sulfide Systems Concentrations of AVS 570 References 4.4. Direct Observations of Mackinawite and Griegite in Sediments 573

Developments in Sedimentology. http://dx.doi.org/10.1016/B978-0-444-52989-3.00014-3 Copyright Ó 2012 Elsevier B.V. All rights reserved.

576 579 581 581

584 586 587 588 588 589 589 590 593

543

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Chapter | 14 Sedimentary Sulfides

In the Blue Muds the decomposition of this matter in the deeper layers leads to the reduction of the oxides in the red upper layer and to the formation of sulfides, which give a blue color to the deposit, but in the Red Clays and Red Muds the quantity of organic matter is insufficient to completely effect this change, and the deposit as a whole remains of a red color. John Murray & Rev. A.F. Renard, 1891. Report on the Deep-Sea Deposits Collected During the Voyage of H.M.S. Challenger. p. 254.

1. BACKGROUND The total sulfur content of normal marine sediments (i.e. excluding euxinic environments) varies up to 2.5% dry weight sediment (Fig. 1, recalculated from Goldhaber, 2004). The distribution is skewed toward 0 wt % total S suggesting that a majority of samples contain only minor amounts of S. However, the mean concentration is around 0.6 wt % (Goldhaber, 2004). Goldhaber pointed out that if the sulfur content of sediments was solely the result of trapping of seawater sulfate and its in situ reduction, the amount of reduced sulfur would be between 0.2 and 0.6 wt %. The reason it exceeds this figure is that the sedimentary sulfur system is not closed. In particular, it is open to sulfate addition through advection and bioturbation, and sulfide diffuses through the sediments along concentration gradients, particularly to the surface where it is mostly re-oxidized. The discrimination of the sulfur components of sediments is analytically difficult. The problem is that the analytical protocols available at present do not distinguish quantitatively either dissolved and particulate sulfur phases or the various solid sulfur phases. The analytical problems are confounded by the sensitivity of most of the sulfur phases to oxidation. In addition, the errors intrinsic to many of the sophisticated analytical protocols that have been developed are cumulative. FIGURE 1 Sulfur content of normal marine sediments, excluding euxinic samples. The equivalent pyrite concentrations are indicated (Recalculated from data collected by Goldhaber (2004)).

1. Background

545

For these reasons the sulfur and sulfide components of modern sediments are often simply reported in terms of the concentrations of those phases that release H2S gas on reaction of the sediments with non-oxidizing mineral acids (acid soluble sulfide or AVS) and those that require oxidizing acids to dissolve. This may appear a quaint division to outsiders, somewhat reminiscent of the odd classifications of the natural world that are reported in some ancient scientific texts and similarly derives from the limitations of contemporary technology. However, it is necessary to include this approach in describing the nature of sulfur in sediments in order to relate the text to the recent literature on the subject. The forms of sulfur with oxidation numbers
0 in sulfidic sediments are listed in Fig. 2. The distribution and chemistry of forms of sulfur with oxidation numbers >0 are discussed in Chapters 2 and 13.5. The true dissolved phasesdas against the operationally defined dissolved components discussed belowdinclude dissolved S(II) species which, as shown in Chapter 2.4, are mainly aqueous H2S and HS. Aqueous FeS clusters (Chapter 4.3) appear widely distributed. Organic sulfide species (Chapter 2) provide an important if variably significant potential reduced sulfur reservoir. Dissolved polysulfides and sulfur species may measurably contribute to the total dissolved sulfur budget of some sediments but are significant because of their roles in pyrite formation and the final sequestration of sulfur in sediments.

FIGURE 2 Summary listing of forms of sulfur in sulfidic sediments. The dissolved and solid phases are listed in order of their relative abundance. For a color version of this figure the reader is referred to the online version of this book.

546

Chapter | 14 Sedimentary Sulfides

Various Fe sulfide nanoparticles (Chapter 5.3) are probably the most abundant solid reduced sulfur phases in some sulfidic sediments. These include not only metastable phases such as mackinawite and greigite which appear to have a limited distribution in marine sediments, as discussed below, but also nanoparticulate pyrite. The problem here is mainly technical since the identification of nanoparticulate sulfides in sediments is difficult at present. Pyrite is the most abundant and widespread observed sulfide phase in sediments. Mackinawite and greigite are relatively uncommon. Other metal sulfides, which include iron sulfides such as marcasite as well as zinc, copper, nickel, cobalt and lead sulfides significant locally.

2. SULFIDE PHASES IN SEDIMENTS A zone is commonly observed in fine-grained marine sediments which includes bluish to grey colored material which was referred to as “blue mud” by early navigators (Chapter 1). In some sediments, especially in freshwater systems, wetlands and euxinic basins, the sulfide-rich zone includes a black-colored layer which has a finite thickness (Fig. 3). Metastable iron sulfides have been widely identified with this black sulfiderich zone because of their dark colors and the fact that they tend to partially dissolve with HCl to produce H2S gas. There have also been solubility

FIGURE 3 Sulfide-rich zone in a salt marsh sediment. A plug-in core is pictured as being taken from the sediment and the resultant core section is inset. The surface sediments are typically a brown-red color at least partially including iron (oxy)hydroxides. The black AVS zone emits a strong odor of H2S, indicating that much of the sulfide evolved in acid treatment is derived from dissolved S(II). Even cursory visual examination of the inset core shows a quite complex color structure within the black zone, suggesting a complex composition. (Reprinted with permission from Rickard and Morse, 2005). For a color version of this figure, the reader is referred to the online version of this book.

2. Sulfide Phases in Sediments

547

considerations which have attempted to demonstrate that the concentration of dissolved Fe(II) in this zone is controlled by the solubility of FeSm. This sulfide-rich zone at depth includes a variety of solid and dissolved iron and sulfur compounds (Fig. 4) and their concentration is represented by total sulfur measurements. As discussed below (3.2) total reduced sulfur (TRS) is not the same as total sulfur. Forms of reduced sulfur in sediments are operationally divided into filter passing (which are often described as dissolved) and solid phases (Fig. 4). Conventionally, 0.45 mm filters have been used for separating these compounds but a number of relatively large particles as well as nanoparticles may pass though these filters. The filters tend to give a better result than is obvious from the stated size because clogging reduces real filter passing sizes and many of the smaller particles flocculate and/or absorb to larger (e.g. clay) particles. The problem is that these effects are very empirical and not entirely reproducible. Much smaller filters (commonly 0.02 mm, but down to 0.01 mm) are available, but tend to be extremely slow in operation. A pre-filter is usually employed anyway. Filtration with small filters takes a longer time and is often used in conjunction with a pressure system both features may affect the composition of the filtrate.

2.1. Acid Volatile Sulfide One way around the filtration problem is to add HCl to the sulfidic material and measure the concentration of H2S gas that is evolved. The result is a measure that is

FIGURE 4 Operational classification of forms of reduced sulfur in sediments.

548

Chapter | 14 Sedimentary Sulfides

called the “acid volatile sulfide(s)” or AVS component of sedimentary sulfides. As is discussed below, acid reduced sulfide usually constitutes only a small fraction of the TRS in sediments. However, the ease with which it can be measureddas well as the occasional high concentrations in especially, freshwater and euxinic environmentsdmeans that it has attracted disproportionate interest in the community. The interpretation of the origins of AVS is not straightforward. When John Morse and I wrote the critical review on AVS in 2004, we discussed whether the subject was acid soluble sulfide or sulfides. The difference is not merely semantic but reflect a basic dichotomy of understanding between the original operational definition of AVS as simply the sulfide that is released with HCl addition and the mistaken assumption of many workers that this sulfide is composed of the metastable FeS minerals, mackinawite and greigite. That is that AVS is synonymous with mackinawite and greigite contents of the sediment. This is obviously wrong, as shown below, and neither phase has actually been widely observed in sediments. It may well be that using the acronym AVS encourages the belief that AVS is a defined phase or phases. In fact, it is probable that in many systems AVS often more closely approximates dissolved sulfide (H2S þ HS) in the porewater. This is obvious where sulfide-rich sediments emit H2S directly on samplingdwhich is a convoluted way of saying that you can often smell the H2S in sulfidic sediments. It is an interesting historical note that, in the sediments from the Gulf of California that Berner (1964a) studied in his original work on acid dissolved sulfide, he reported that most of the AVS was evolved from dissolved sulfide in the porewaters, a point lost on most ensuing studies by others. A variant of the basic AVS method is the simultaneously extractable metal (SEM-AVS) approach. In this technique, the metals, mainly iron, extracted by the HCl addition are also analyzed in an attempt to assess their potential toxicity. The measurement of the concentration of AVS has become routine in sedimentary investigations and SEM-AVS analyses have been used as a standard environmental indicator in environmental protection agency protocols. An earlier paradigm was that sulfide is first removed by reaction with Fe(II) to form metastable iron sulfides in this black or AVS zone which then transforms on aging to pyrite. This, as we have already seen in Chapters 5 and 6 for example, is wrong: 1. Sulfide is only removed to form metastable iron sulfides in those environments where the solubility product of the metastable forms is exceeded. 2. Metastable iron sulfides are not necessary precursors to pyrite formation and most sedimentary pyrite is formed in environments where metastable iron sulfides do not form. 3. Metastable iron sulfides do not transform to pyrite on aging. Where they are present they may represent one possible source of Fe(II) for pyrite formation from many other sources.

3. Analytical Chemistry of Sedimentary Sulfide

549

4. The formation of metastable iron sulfides may actually inhibit pyrite formation. Where mackinawite forms the amount of dissolved Fe and S available for pyrite formation is limited by the solubility product of mackinawite, FeSm, which is commonly around 106 M in natural sediments. Greigite is more stable that mackinawite and has an even more limited solubility (Chapter 5). It appears to be a dead end in sedimentary pyrite formation.

3. ANALYTICAL CHEMISTRY OF SEDIMENTARY SULFIDE Our knowledge of the biogeochemistry of sedimentary sulfides stems from chemical analyses of sulfidic sediments. The analyses of low temperature, experimental sulfide product in the laboratory is difficult, as described in Chapters 5 and 6, so the analyses of natural sediments are heroic enterprises, especially when including analyses on ships. An example of an analytical protocol for the solid phases of sulfidic sediments is shown in Fig. 5. The protocol is complex and the reported analyses may require four or more separate manipulations of the original sediment sample. The problems, as defined for the laboratory products, include: 1. The sampling itself. By definition, any sample removed from a sediment disturbs the sediment. Several different coring methods have been employed and, often, extreme measures have been taken to obtain pristine samples, such as sampling the interior of the core. 2. The sensitivity of sulfidic materials to oxidation. The solid iron sulfides oxidize rapidly in air. Indeed, as noted in Chapter 5, greigite and mackinawite are pyrophoric. Samples of sediments containing mackinawite or greigite should visibly oxidize when exposed to the atmosphere and this might be used as an indicator of the presence of these metastable iron sulfides. Immediate freezing of samples and handling under an inert gas atmosphere are highly recommended (e.g. Lasorsa and Casas, 1996). 3. The nanoparticulate nature of much of the sulfides and the problem of discriminating dissolved and particulate phases. The presence of nanoparticulate sulfides is a particular problem. 4. The difficulty of quantitative dissolution of solid sulfide phases (see below). 5. The lack of any robust protocol for discriminating between different iron sulfide solids in mixtures, as discussed in Chapter 5.

3.1. Analytical Precision and Accuracy A major problem is precision. The analytical protocols that have been developed and, as indicated in Fig. 5, are complicated and often involve several experimental steps in order to obtain the final analytical result. Although the errors on each step can be measured, the total error on the analytical result is the cumulative errors of each step in the process. Specifically, the total uncertainty in the

550 Chapter | 14 Sedimentary Sulfides

FIGURE 5 A sophisticated protocol for analysis of solid phases in sulfidic sediments (Reprinted with permission from Werne et al., 2003).

3. Analytical Chemistry of Sedimentary Sulfide

551

analyses is the square root of the sum of the uncertainties of the individual processes in the analytical protocol. Precision is not discussed in detail in most reports of sedimentary sulfide analyses. However, the experimental experienceddiscussed particularly in Chapter 5dis that the best precision for a simple analysis of a pure solid sulfide phase under ideal conditions may approach a one sigma value of 3%. The error in this case (Rickard et al., 2006) appeared mainly in the sampling and digestion procedures. This degree of precision was not easily attained and it is probable that, routinely, the precision of any individual measurement involving, say, four procedural steps, is likely to produce a one sigma error of >25%. This scale of precision is difficult to reduce through repeat analyses in natural samples: there are only a limited number of splits that can be taken from a single sample of natural sediment considering the size restrictions on duplication imposed by natural variations in the measurements. Few attempts have been reported which assess the accuracy of sedimentary sulfide analyses since the original work of Kaplan et al. (1963). The approach is to measure the total sulfur in a sedimentdthrough the combustion techniques discussed belowdand see how near the totals of the various extractions approach the total. Kaplan et al. (1963) found that their analyses of individual sulfur extracts from sulfidic sediments were usually less than the total sulfur concentrations collected by combustion techniques. However, the sum of their extractions was within 70–90% of the total which suggests a good, first order accuracy in the analytical protocol. Unfortunately, these data appear less frequently in more recent published reports.

3.2. Total Sulfur and Total Reduced Sulfide Total sulfur contents of sediments are usually measured by combustion of the dried sediment in air and analyzing the evolved sulfur oxide gases. The total sulfur includes organic-S which may include significant quantities of reduced sulfur. TRS compositions are routinely reported as the results of sulfur analyses after reducing all inorganic sulfur to S(II) with Cr(II), for example. The TRS concentrations appear to be relatively robust measurements of sediment pyrite contents (cf. Canfield et al., 1986). One remarkable feature (Raiswell and Canfield, 1998) is that all the TRS content of marine sediments of all types closely correlated with the pyrite–S content (Fig. 6). The samples were collected from deep-sea, continental margin and continental shelf environments as well as oxic, suboxic, anoxic and euxinic systems. The actual correlation found had a correlation coefficient of 0.89. In view of the probable relative precisions of the analytical protocols employed this suggests that the observed variation constitutes a direct, 1:1 relationship between TRS and pyrite concentrations and that the correlation approaches 100% of the observed variance. The relationship is used in Fig. 1 to convert S measurements to wt % pyrite. Weight % pyrite–S in Fig. 6 can be converted to weight % pyrite by multiplying by 1.875 or approximately doubling the pyrite–S value. The results

552

Chapter | 14 Sedimentary Sulfides FIGURE 6 Relationship between TRS and pyrite–S in marine sediments. The slope of the line is 1.0. (From data collated by Raiswell and Canfield, 1998).

show that most of the reduced sulfur in marine sediments of all types is present as pyrite. The correlation is surprising since the chromium reduction method also releases part of the organic-S. Cr-reduction liberates some organic polysulfides (Ferdelman, 1994) as well as sulfur from humic extracts (Francois, 1987b) although it does not release significant assimilatory sulfur from living organisms (Canfield et al., 1986). The use of voltammetry and chromatographic methods to determine the products of the reduction distinguishes the organic sulfur products from the inorganic products. However, iodine titrationsdwhich have been widely usedddo not. The reason for the correlation between pyrite–S and TRS is probably related to the relative small fraction of organic-S relative to organic-S in most sediments. Sediments or sediment samples with relatively high organic-S fractions compared with pyrite–S fractions are contained in the >10% variance in the 1:1 correlation shown in Fig. 6.

3.3. Analysis of Organic Sulfur More than 1000 organic-S compounds have been identified chromatographically and spectroscopically from sulfidic sediments (Russell et al., 2000). Quantitative analyses of organic sulfur in sulfidic sediments depend on the extraction of organic sulfur compounds. The major classes include thiols, organic sulfides, disulfides, polysulfides, thiophene derivatives (Tissot and Welte, 1984) and sulfonates (Vairavamurthy et al., 1994). Because of the complexity of organic-S compounds in anoxic sediments, it is often difficult to quantitatively analyze individual organic-S compounds. However, the compound specific sulfur isotopic composition of chromatographic separates can be determined and this provides a powerful probe into the organic sulfur system in sulfidic sediments (e.g. Werne et al., 2008; Passier et al., 1997).

3. Analytical Chemistry of Sedimentary Sulfide

553

Humic material was a term originally used to describe the dark brown material extracted from soils by alkaline solutions and is currently a general term describing the macromolecular organic compounds that can be extracted from sediments in alkaline solutions. Some marine sediments can contain as much as 70% of their organic C in the form of humic material. Although freshwater and near continental humic material derives from familiar terrestrial sources, much of this material in marine sediments represents the degradation products of plankton (Nissenbaum and Kaplan, 1972). Francois (1987b) showed that relatively quantitative amounts of humic substance can be extracted from sulfidic sediments if free sulfide and elemental S are removed before the alkali extraction. Even so, pyrite–S can contribute to the organic-S contents of the extracted humics and organic-S measurements need to be corrected for pyrite–S contamination. By contrast, the reverse does not seem to be the case: Bru¨chert (1998) noted that the NaOH extraction had little measurable effect on the concentration of TRS. Analytical protocols for extracting the major classes of organic sulfur compounds in sulfidic sediments (humic acids, fulvic acids and protokerogen) have been developed (Bru¨chert, 1998; Ferdelman et al., 1991; Francois, 1987b; Zaback and Pratt, 1992). Acidification of the humic extracts precipitates humic acids. The soluble fraction includes smaller macromolecular humic compounds classed as fulvic acids. The residue, very large organic polymers, is protokerogen (Bru¨chert, 1998) which condenses to kerogen during diagenesis. Canfield et al. (1998) pointed out that studies of operationally-defined classes of organic extracts, such as fulvic and humic acids, may only address a subset of the total organic sulfur. Analytical protocols for low-molecular weight organic polysulfides have been developed by Henneke et al. (1997) and Yu¨cel et al. (2010). The process involves the initial extraction of elemental sulfur and low molecular weight, nonpolar organic sulfur compounds with a 3:1 mixture of methanol and toluene. This extract contains low molecular weight cyclic R–Sn–R (n > 3) and linear R–Sn (n > 1) organic polysulfides (Henneke et al., 1997) which can be broken down with HCl, whilst at the same time releasing AVS as H2S. The product of the reaction of the low molecular weight polysulfides with HCl is elemental sulfur, which can be analyzed by standard methods. These methods tended to extract organic polysulfides and sulfides. Vairavamurthy et al. (1994) showed that X-ray adsorption near edge spectroscopy (XANES) could be used to probe the composition of organically bound sulfur in sulfidic sediments. The method showed that oxidized sulfur species, sulfonates, constituted a significant fraction of the organic sulfur (Fig. 11). Although the XANES approach provides simultaneous qualitative and quantitative information about the organic sulfur species of whole sediment samples, it has not been widely applied to sulfidic sediments since these early investigations. The problems probably include the time between sampling and analysis, the integration of this technique into conventional analytical protocols, the

554

Chapter | 14 Sedimentary Sulfides

difficulty of discriminating detrital and authigenic organic-S, and getting time on the synchrotron.

3.4. Analysis of AVS The primary differences in extraction techniques have been centered on the measurement of the AVS component of sulfidic sediments. This has been used not only as a measure of the concentration of possible non-pyritic iron sulfide phases but also as a correction factor for the pyrite concentration. Conventionally, the pyrite concentration is the difference between the TRS concentration and the concentration of AVS. Although there was some earlier work (e.g. Byrne and Emery, 1960; Kaplan et al., 1963) the seminal paper for the concept sedimentary AVS was that of Berner (1964a). He defined AVS in terms of the H2S evolved from sediments using essentially the method for sulfide evolution of Kolthoff and Sandell (1952) and was inspired by the earlier work of Ostoumov (1953) on Black Sea sediments. This method consists of adding enough 1 N HCl to cover a weighed sediment sample in a flask and heating the flask until all H2S is expelled by boiling. Gas (usually N2 or Ar) is allowed to flow through the flask during this process and the H2S collected as ZnS in a Zn acetate trap. The sulfide is subsequently analyzed upon reacidification of the ZnS. Various workers have used hot (Berner, 1964b; Kolthoff and Sandell, 1952; Westrich, 1983; Wieder et al., 1985; Zhabina and Volkov, 1978; Berner et al., 1979) or cold HCl (Albert, 1984; Aller, 1977, Aller, 1980a; Brumbaugh and Arms, 1996; Brumbaugh et al., 1994; Goldhaber et al., 1977; Jørgensen, 1977; Westrich, 1983; Allen et al., 1993; Chanton and Martens, 1985; Cornwell and Morse, 1987) and included SnCl2 (Berner et al., 1979; Cornwell and Morse, 1987; Westrich, 1983) or TiCl2 (Albert, 1984) to reduce interference from Fe(III) by keeping the solution reducing. Both 1–3 N and 6 N HCl concentrations are also in common usage, but 6 N HCl is probably more reliable for sediments that may contain substantial other acid soluble components such as carbonate minerals because it offers a greater dissolving capacity per unit volume. Albert (1984) used 1 N H2SO4 instead of HCl with TiCl2. Although within a single report the analytical results may be comparable, it is difficult to compare analytical results from different laboratories at different times. The extraction efficiency or precision of the digestions of iron sulfide minerals was addressed directly by Cornwell and Morse (1987). Table 1 lists results for digestions of laboratory materials. It should be noted that it is probable that, as discussed in Chapter 5, the synthetic greigite used probably included mackinawite (see Chapter 5.4.2) and therefore the extraction efficiency listed for this material is a maximum value. Allen and Parkes (1995) and Polushkina and Sidorenko (1963) reported similar results. They found that only 81  3% of the mackinawite was recovered in hot 6 M HCl digestions over 1 h and 104  14% recovered in cold

555

3. Analytical Chemistry of Sedimentary Sulfide

TABLE 1 The Extraction Efficiency of Different Techniques used for Synthetic Mackinawite and Greigite Presented In Order of Increasing Strength. Synthetic and mineral (<50 mesh) pyrite results are also included (data from Cornwell and Morse, 1987) % Mineral S Extracted Method

Mackinawite Greigite Synthetic Pyrite Mineral Pyrite

1 N HCl

92

40e67

0

0

6N HCl þ SnCl2

100e102

63e75

4e10

0

1 N H2SO4 þ Ti

101

96e100

48e82

0

Hot 6 N HCl þ SnCl2

100e101

93e100

97e100

2

6 M HCl digestions over 1 h. The cold 1 N HCl leach does not quantitatively extract either mackinawite or greigite sulfur and does not release sulfide from pyrite. The reasons were demonstrated by Rickard et al. (2006) and are described in Chapter 5. Basically the problem centers on the formation of solid sulfur in acid dissolutions at normal oxidation potentials. However, the evidence suggests that there is a potential considerable uncertainty (e.g. 28% for hot 6 M HCl) in the measurements of AVS in sediments. This uncertainty is compounded by the precision of the other processes in the analytical protocol, such as sample preparation. For example, drying the sediment sample may affect recovery of AVS. Cornwell and Morse (1987) noted that the recovery efficiency of AVS from dried FeS was less than that of wet FeS. This seems to be a general observation. The reasons are unknown. I have noted that dried FeS has a strong static charge and it may be that this contributes to a difficulty in wetting the sample surface. Dried FeS also tends to flocculate into hard cakes with limited pore space (cf. Watson et al., 2000) this may reduce the surface area to such a degree that dissolution is less efficient. At the other end of the extraction efficiency range is the hot 6 N HCl þ SnCl2 which, while putatively extracting mackinawite and greigite, also results in close to quantitative production of sulfide from synthetic pyrite and minor (2%) production of sulfide from mineral pyrite. It is noteworthy that both the cold 6 N HCl þ SnCl2 and hot 6 N HCl, while quantitatively extracting mackinawite and a major portion of greigite, also extract portions of synthetic pyrite similar to fractions of total reduced sulfide often observed as AVS in sediments (respectively 4–10% and 5–6%), but insignificant macroscopic mineral pyrite is extracted. The 1 N H2SO4 þ Ti method quantitatively extracts both mackinawite and greigite, and extracts up to 82% of the synthetic pyrite but none of the mineral pyrite. Thus, no method is capable of quantitatively

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Chapter | 14 Sedimentary Sulfides

TABLE 2 The Percentage of TRS that is Extracted as AVS by Different Methods for a Variety of Different Sampling Sites (Based on Morse and Cornwell, 1987)

Method

Skan Bay

Tidal Creek

Cape Lookout Bight

Louisiana Shelf

Mississippi Delta

FOAM

1 N HCl

43

21

31

0

9

5

6 N HCl þ SnCl2 57

27

33

0

14

7

1 N H2SO þ Ti

52

31

38

0

13

12

84 Hot 6 N HCl þ SnCl2

55

67

0

83

56

511

246

85

132

142

TRS Conc. (mmol gdw1)

246

extracting both mackinawite and greigite S without also significant portions of S from pyrite as well. Consequently, even in the absence of dissolved sulfidic constituents in a sediment, there is no reliable method for the quantitative extraction of the metastable iron sulfides, mackinawite and greigite, from sediments. All this suggests that the overall precision of AVS measurements may be very poor indeed. This is of considerable concern not least because this leaching method is the basis of most SEM-AVS studies which are widely used as standard protocols for environmental assessments. An attempt to determine the accuracy of AVS analyses in sulfidic sediments by Morse and Cornwell (1987) is listed in Table 2. Results for sediments from a wide variety of locations (Table 2) demonstrate that the hot 6 N HCl þ SnCl2 consistently result in substantially higher AVS concentrations than the other methods and the cold 1 N HCl AVS concentrations are always lowest. This is consistent with the mineralogical studies. The relative amounts of AVS extracted by the different methods vary considerably among sites. A likely explanation for this is that the dissolved to mineral ratio of AVS components can vary greatly. It is difficult to conclude that AVS concentrations represent anything very specific in sediments.

3.5. Analysis of Porewater Sulfide As described in Chapter 2, the chemistry of S(II) in aqueous environmental systems has been relatively well-constrained. pK1, H2S is close to 7 which means that H2S dominates the system at acid pH values and HS is the dominant species in alkaline solutions. pK2, H2S is less precisely constrained but is

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557

estimated to be around 18. This means that the aqueous sulfide ion, S2, has no significant activity in natural aqueous systems. The concentrations of dissolved sulfide in sediment porewaters are reported in terms of total dissolved sulfide which is usually approaches the sum of the H2S, HS and aqueous Fe sulfide complex and cluster concentrations. Byrne and Emery (1960) and Kaplan et al. (1963) had earlier attempted to separate dissolved and solid components in sulfidic sediments, but Berner (1964a) reported unacceptably high losses of sulfide by their methods. The standard method for determining dissolved S(II) in sediments is through taking a sample, usually in the form of a core, expressing the porewater by centrifugation or pressure, filtering and analyzing the filtrate (Fig. 7). The results are imprecise and may include a number of particulate as well as dissolved species, as discussed above (Fig. 3). In the scheme shown in Fig. 7 there is no information as to how the porewater was obtained in the first place. This is usually obtained by suction, squeezing (Bender et al., 1987; Reeburgh, 1967) or centrifuging (Lyons et al., 1979) the sediment over a 0.45 mm filter often on board ship. These processes are basically destructive and introduce inaccuracies into the analyses mainly through dissolution and the effects of pressure/temperature changes but also through the proportion of the fluid collected and how far that represents the porewater composition as a whole (e.g. Hammond, 2001). A further problem is the resolution of the method. A sizeable sample is needed to extract sufficient porewater for analysis and this requires relatively large sediment samples. The resulting analytical data are the result of spatially integrated measurements over significant vertical or horizontal sediment intervals. Alternative methods for porewater extraction from sediments include direct sampling with syringes or suction probes. These are less suitable for fine siliciclastic sediments. Passive, in situ, sampling methods called peepers, involve the insertion of passive diffusion cells into the sediment which equilibrate with the porewaters

FIGURE 7 Protocol for analyzing porewater from sulfidic sediments (Reprinted with permission from Werne et al., 2003).

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Chapter | 14 Sedimentary Sulfides

(Mayer, 1976; Sayles et al., 1973). These peepers usually consist of isolated volumes of water surrounded by a dialysis membrane. The problem of knowing when equilibration has been reached between the peeper and the porewater has been addressed by Thomas and Arthur (2010). The peepers are then extracted and the sulfide measured in the extracted fluid by conventional methods. An alternative approach is to fix the porewater components in thin films, usually a polyacylamide hydrogel, which equilibrates with the pore solutions (Davison et al., 2000). Two basic arrangements are used: diffusive gradients in thin films (DGT) and diffusive equilibration in thin films (DET). The DET gel techniques suffer from diffusional relaxation, as the components sequestered in the gel diffuse to areas with lower concentrations. The DGT technique involves accumulation of the analytes within a second gel layer that includes a specific agent to bind the analyte. DGT, often in combination with DET, has been used to determine the sulfide distribution in sediment porewaters (Motelica-Heino et al., 2003; Teasdale et al., 1999; Widerlund and Davison, 2007). The DGT and DET techniques are elegant and can be used to report sulfide and dissolved metal distributions in sediment porewaters at a fine, micrometer scale. However, they are mostly semi-quantitative (although a neat reverse tracer diffusion technique has been developed by Thomas and Arthur, 2010), their insertion into the sediment may cause disturbance at a scale which is larger than the reported measurement discrimination and they do not readily distinguish different chemical species of the same ion. More accurate measurements of dissolved sulfide species in sediments are obtained by in situ electrode measurements. Commonly the cores are probed directly with solid state microelectrodes pioneered by Luther and his coworkers (e.g. Luther et al., 1998, 2001, 1999), with typical tip sizes of 100 mm. The differences between this method and conventional analytical methods are difficult to determine because of natural variations even at the same sample site. However, Konovalov et al. (2006) reported sulfide concentrations of 1600 mM using microelectrodes which compares with the highest measurement of 400 mM using conventional methods from similar sulfidic sediments in similar locations in the Black Sea (Lyons and Berner, 1992). It may well be that this is partly due to the superior depth resolution of the microelectrodes compared to the larger integrated sediment volumes used in conventional extractions. When we first deployed microelectrodes in a sulfidic esturarine mud (Rickard et al., 1999), we were surprised that clear signals for the dissolved sulfide species were obtained. We had expected a complex spectrum reflecting the complicated composition of natural sediments. In fact, the species that the microelectrode sees in these systems is relatively limited, both by concentration, the potentials of the voltammetric scans and the lability of the species with respect to the mercury electrode. Voltammetric analyses provide information on speciation and thus the dissolved sulfide species, such as H2S, HS, FexHSy and the aqueous clusters

559

4. Solid Sulfide Phases in Sulfidic Sediments

FIGURE 8 Cyclic voltammetry of an iron sulfide system. Scanning the potential in the negative direction causes three reactions at the Hg electrode: H2S, FeS(aq) and Fe2þ are removed by reaction with the Hg. Scanning back in the positive direction releases these components from the Hg. (Reprinted with permission from Rickard and Morse, 2005).

can be identified. Analysis of the current provides information on the concentration of the dissolved species present, although this needs to be quantified for individual matrices. In the example shown in Fig. 8, the size of the voltammetric peak (#1) in the positive scan should equal the size in the negative scan since the peak height records the current and this is a function of the concentration. However, FeSaq is not released in the negative scan as a discrete entity, as predicted by theory. Instead a larger peak (#6) is returned, being the cumulative S(II) that was fixed as FeSaq and H2S. This larger peak (#6) could be interpreted as being a measure of dissolved AVS. However, the use of the electrochemical approach in combination with conventional analyses of the acid–leachate and porewater S(II) may provide a better insight into the nature of sedimentary sulfide although the errors inherent in this process discussed above still mean that the precision of the results remains poor.

4. SOLID SULFIDE PHASES IN SULFIDIC SEDIMENTS Analytically, the solid inorganic phases in sulfidic sediments are often divided into acid dissolved sulfide and pyrite, which is reported as the difference between the concentration of TRS in the sediment and the acid dissolved sulfide concentration. However, as we have seen above, both the analytical methodology and the interpretation of the results of AVS measurements are extremely variable and uncertain. Although various schemes have been proposed for analyzing mackinawite and greigite in pyritic sediments it would seem that all these schemes suffer intrinsically from considerable uncertainties (cf. Chapter 5). The problems

560

Chapter | 14 Sedimentary Sulfides

encountered include considerable potential deviations from 100% recovery, lack of reproducibility through large standard deviations in analyses and considerable overlaps between the contributions of each phase to different stages in the digestion procedure. AVS is an operationally-defined term that in itself contains no intrinsic information as to the sources of the evolved H2S. It should never, therefore, be used interchangeably, as is common practice, with sedimentary FeS or metastable iron sulfides because this implies specific mineral sources. This is not the case for all sediments and, indeed, as previously noted, some studies have assumed that AVS is entirely associated with iron sulfide minerals and the terms AVS and FeS have been used interchangeably (Morse, 1999). However, this is incorrect, as can be demonstrated using a simple formula    f 3 (1) AVSM ¼ AVST  cH2 S  10 ð1  frM Þ where the subscripts T ¼ total, M ¼ mineral; c ¼ dissolved (including clusters and nanoparticles) sulfide concentration in mmol L1, f ¼ porosity, and rM ¼ mineral density. AVS is in mmol gdw1 (note that although extractions are usually performed on wet sediment concentration, they are generally reported in dry weight units). The fraction of total AVS present in the mineral fraction (FM) can then be calculated    cH2 S  103 f (2) FM ¼ 1  AVST 1  frM For an AVST concentration of 10 mmol gdw1, f ¼ 0.7, and rM ¼ 2.6 g cm3, dissolved sulfide does not become important until it reaches a concentration of about 1000 mM (FM ¼ 0.91), but above 11 mM (FM ¼ 0.09) AVS is almost entirely in the dissolved phase. This may provide an insight as to why, in highly sulfidic sediments, where iron is limiting for sulfide mineral formation, there is a persistence of AVS with depth in the sediment. However, AVS may be almost all dissolved sulfide species, as originally shown by Berner (1964a), and have no relationship to the content of metastable iron sulfide minerals. A similar result was found by Canfield et al. (1998). As shown below, porewater sulfide is part of a dynamic system where it may be differentially generated at different levels within the sediment and move mainly along diffusion gradients within the sediment. So the discovery of high concentrations of AVS at various depths in the sediment should be no surprise. In the extreme example, of course, AVS is trapped in ancient sedimentary rocks, such as limestones known as “stinkstones” (Chapter 1). The concern of some earlier workers about apparently anomalous concentrations of AVS within sediments may partly reflect the mistaken identification of AVS with FeS. Against this

4. Solid Sulfide Phases in Sulfidic Sediments

561

conceptual background, the apparent preservation of metastable mackinawite over long time periods was a major problem.

4.1. Pyrite Distribution in Sediments Pyrite is commonly observed in sulfidic sediments. However, measurements of its concentration are not straightforward. Few if any point-counting studies have been reported and, anyway, they would probably be relatively imprecise. Pyrite concentrations are measured as the difference between total reduced sulfide and AVS and, as we have seen above, the AVS methods may collect sulfur from the dissolution of fine-grained pyrite. The result is that the reported concentrations of pyrite in unconsolidated sediments are probably minimum values. This needs to be kept in mind when considering the uncertainties in measurements of the degree of pyritization (DOP) in recent sediments, described in Chapter 13.7.3. A typical coastal sediment contains 0.5–1.0 wt % pyrite–S (Canfield et al., 1998) which is about 1–2 wt % pyrite. The amount of pyrite in Raiswell and Canfield’s (1998) compilation for sulfidic sediments from all types of marine environments (Fig. 6) ranges up to over 4 wt % or about 1.6 vol%. This low volume percentage is caused by the relative density of pyrite compared with the sediment. In terms of the wet sediment with a nominal 80% porosity, 4 wt % pyrite–S is equivalent to about 0.3 vol %. The point is that even in sulfide-rich sediments, pyrite is a relatively minor phase. Four wt % sulfur sequestered in the pyrite in sulfidic, marine sediments is equivalent to a concentration of over 600 mmol g1 (or 3 mM in a wet sediment), underlining once again the requirement for reduced sulfide to be transported to the site of pyrite formation in an open system. In Chapter 6 it was shown that experimental data proved that mackinawite and greigite does not transform to pyrite through a solid-state reaction. The formation of pyrite with mackinawite and greigite as reactants involved the dissolution of these phases and the reaction of the dissolved Fe(II) and sulfide species to form pyrite. In other words, there is a complete disconnect between the formation of pyrite and that of mackinawite and greigite. These phases behave in the same way as other iron reactants. It is interesting then to examine if the same lack of any connection is observed in natural sulfidic sediments. This is more difficult against the background of the problem of chemically defining the nature of the sulfide phases in sulfidic sediments and the uncertainties in the analyses. In earlier studies, the mistaken assumption that sedimentary pyrite formed from metastable iron sulfides, such as mackinawite and greigite, and that these were identical to AVS, caused some confusion when real sedimentary sulfide systems were analyzed. In this paradigm, the concentration of pyrite should increase with sediment depth and the concentration of AVS should increase with depth. The consternation was caused by the common observations that

562

Chapter | 14 Sedimentary Sulfides

pyrite was abundant in the surface layers, that pyrite formed in sediments where AVS was below analytical detection limits and that even in the rare cases where an inverse relationship was observed, the mass balances between pyrite and AVS did not add up. For example, at classic study sites like FOAM in Long Island Sound almost no AVS is observed in the uppermost zone of sediment where most pyrite is produced Fig. 9. These observations suggested that there was no simple, direct relationship between AVS and pyrite. The proof of the disconnect between metastable iron sulfides and pyrite formation in natural sediments was first demonstrated in an elegant experiment by Bob Howarth in 1979. Howarth (1979) used 35S as a tracer in salt marsh sediments and showed that 70% of the 35S ended up in pyrite. The remainder was in soluble sulfides and metastable iron sulfides, mackinawite and greigite, were absent. Howarth wrote (ibid p. 50) This finding is consistent with the hypothesis that pyrite forms rapidly at low temperatures only when iron monosulfides are undersaturated.

Howarth’s study was important since he first demonstrated that pyrite formation was not related to the formation or presence or other authigenic iron sulfide

FIGURE 9 Dissolved Fe and H2S profiles at the FOAM site (Aller, 1977). (Reprinted with permission from Rickard and Morse, 2005).

4. Solid Sulfide Phases in Sulfidic Sediments

563

minerals such as mackinawite and greigite. As noted above, these minerals are often mistakenly identified with AVS (AVS) in sediments and Howarth (1979) explicitly stated that pyrite formation in sediments occurred in the absence of measurable AVS concentrations and was actually more rapid when AVS was absent. The counterintuitive conclusion is that the formation or presence of potential solid phases in the AVS fraction, such as mackinawite and greigite, actually inhibits pyrite formation. The reason, as mentioned in Chapter 6, is probably that the dissolution of these mineralsdwhich is necessary to provide the dissolved Fe and S for pyrite formationdis relatively slow and acts as a kinetic bottleneck. Unfortunately, his results were ignored by a section of the sedimentary sulfide community and the AVS “myth”, as John Morse called it, remained potent into the 21st century. Although there are relatively few reports of detailed analyses of sedimentary sulfides which include a breakdown of both dissolved and solid phases, many investigations of the concentrations of the simpler AVS and pyrite have been reported. Most studies have focused on the relative abundance of AVS to that of pyrite–S (based on the presumption that no pyrite–S is contributed to the AVS fraction) and the factors that control this relationship in different environments. Goldhaber and Kaplan (1974) summarized much of what was known in the mid-1970s about sedimentary sulfides. They observed that AVS tends to occur in greatest abundance relative to pyrite in rapidly deposited sediments and that, in slowly accumulating sediments, AVS is a minor component of the total reduced sulfide content of sediments or even absent despite an abundance of pyrite. Table 3 presents a selection of many investigations into the pyrite–S: AVS ratio. Very often these reports do not contain the necessary ancillary information to make an interpretation of controlling processes possible (e.g. extent of iron pyritization, porewater chemistry of iron, sulfide, pH, and sulfate, and sediment accumulation rates). The results in Table 3 suggest that in most marine sediments most of the solid sulfur is contained in pyrite. There are exceptions to this where the AVS fraction reaches extremely high values; for example, in the euxinic sediments of the Black Sea, Orca Basin and Saguenay Fjord. However, it would be a mistake to assume that these sediments with high AVS are the norm: in fact, they are reported because they represent exceptional systems. Robert Aller’s dissertation (Aller, 1977), which includes results of Goldhaber et al. (1977) for the FOAM site (Aller, 1980a,b), was the most comprehensive early study of AVS in marine sediments. It should be noted that Aller referred to AVS as FeS, which is an unfortunate use of terminology that has been followed in many subsequent papers. The “classic” profile is shown for the FOAM site in Long Island Sound in Fig. 10A. The TRS increases rapidly down to about 4 cm and then is close to constant. Acid volatile sufide is always a small fraction of the total reduced sulfide and has a maximum value at about 4–6 cm. It then decreases with depth to close to undetectable values. The AVS maximum is close to coincident with the depth at which dissolved Fe decreases to low values and detectable dissolved sulfide rapidly increases (Fig. 10A).

564

Chapter | 14 Sedimentary Sulfides

TABLE 3 Fractions of PyriteeSulfur and AVS in a Selection of Marine and Estuarine Environments. PyriteeS is assumed to be the difference between TRS and AVS. Based on Rickard and Morse (2005) Relative sulfide concentration

Location

Water Depth (m)

Sediment Depth (cm) %FeS2eS %AVS Reference

Aarhus Bay

18

20e240

98e100

0e2

Thode-Andersen and Jørgensen (1989)

Arabian Basin 3566

3e18

83e99

1e17

Passier et al. (1997)

Bannock Basin

3350

7e811

92e 97

3e8

Henneke et al. (1997)

Black Sea

2094

0e40

13e67

33e88 Lyons and Berner (1992)

Bornholm Deep

70e90

0e350

79e100

0e21

Boesen and Postma (1988)

Cabot Straight 495

1e35

71e98

2e29

Mucci, unpublished cited in Gagnon et al. (1995).

Carmen Basin 2772

0e100

30e98

2e70

Goldhaber and Kaplan (1980)

East China Sea

708e1277

0e160

33e100

0e67

Lin et al. (2000)

Effingham Inlet

100e200

0e28

45e91

8e55

Hurtgen et al. (1999)

FOAM

9

3e6

92e100

0e8

Morse and Cornwell (1987)

Gotland Deep

165

0e180

80e100

0e20

Boesen and Postma (1988)

Gulf of Mexico Shelf

161

30e35

98e100

0e2

Morse and Cornwell (1987)

Gulf of St. Lawrence

260e330

1e35

71e95

5e29

Mucci, unpublished cited in Gagnon et al. (1995).

Kau Bay

457

2e756

38e97

3e62

Middelburg (1991)

Kattegat

42

1e12

71e100

0e29

Sørensen and Jørgensen (1987)

Laguna Madre, Texas

1.5

0e25

40e100

0e60

Morse (1999)

565

4. Solid Sulfide Phases in Sulfidic Sediments

TABLE 3 Fractions of PyriteeSulfur and AVS in a Selection of Marine and Estuarine Environments. PyriteeS is assumed to be the difference between TRS and AVS. Based on Rickard and Morse (2005)dCont’d Relative sulfide concentration

Location

Water Depth (m)

Sediment Depth (cm) %FeS2eS %AVS Reference

Kysing Fjord

e

1e8

77e94

6e23

Fossing and Jørgensen (1990)

Limfjorden

10

6e14

57e94

6e43

Howarth and Jørgensen (1984)

Mississippi Delta

509

97e99

74e100

0e26

Morse and Cornwell (1987)

North Sea

26

0e15

53e80

20e47 Sørensen and Jørgensen (1987)

Oman Continental Slope

527

16

99e100

0e1

Orca Basin

2478

0e32

3e20

80e97 Hurtgen et al. (1999)

Pescadero Basin

3361

0e100

87e95

5e13

Goldhaber and Kaplan (1980)

Skagerrak

520

0e15

100

0.00

Sørensen and Jørgensen (1987)

Skan Bay

65

5e10

16e100

0e84

Morse and Cornwell (1987)

Saguenay Fjord (basin)

265

3e42

16e93

7e84

Gagnon et al. (1995)

Saguenay Fjord (head)

90

3e42

12e83

17e88 Gagnon et al. (1995)

St. Lawrence Estuary

335

3e22

19e79

21e81 Gagnon et al. (1995)

Santa Barbara w620 Basin

25e100

89e94

6e11

Kaplan et al. (1963)

Tomales Bay, California

?

0e350

77e98

2e23

Chambers et al. (2000)

Tyro Basin

3370

12e529

93e98

2e7

Henneke et al. (1997)

Passier et al. (1997)

566

Chapter | 14 Sedimentary Sulfides

AVS and TRS (µmol/gdw)

(A) 0

0

50

100

150

200

250

Zone of major sulfide formation

AVS Maximum

Depth (cm)

5

10 TRS

AVS

15

(B)

0

2

AVS and TRS (µmol/gdw) 4 6 8 10

12

14

0

TRS

5

Depth (cm)

AVS

10

15

20

25 AVS and TRS (µmol/gdw)

(C)

0

0

50

100

150

20 TRS

Depth (cm)

40

60 AVS

80

100

120

200

250

FIGURE 10 AVS and TRS depth profiles for (A) the FOAM site in Long Island Sound (Goldhaber et al., 1977), (B) the Louisiana shelf west of the Mississippi River delta (Lin and Morse, 1991) and (C) Laguna Madre, Texas (Morse, 1999). (Reprinted with permission from Rickard and Morse, 2005).

4. Solid Sulfide Phases in Sulfidic Sediments

567

Lin and Morse (1991) conducted a study of sulfide geochemistry throughout much of the Gulf of Mexico and some of its adjacent estuaries. A wide variety of patterns of AVS distributions was observed, few of which resembled those in Long Island Sound. In many cores no detectable AVS was observed, but in others it was a major component of the total reduced sulfide that increased with depth (Fig. 10B). After considering these results and other even more unusual observations in the area (e.g. Morse, 1999 and Fig. 10C) it became clear that there is a very wide variety of patterns of AVS and total reduced sulfide distributions in marine sediments.

4.2. Organic Sulfur Organic-S constitutes the second largest sulfur reservoir in sulfidic sediments and may dominate, containing up to 80% of the total sulfur in sediments (e.g. Zaback and Pratt, 1992). However, the fraction of total sulfur associated with organic compounds in recent sediments usually varies between 10 and 20%, which suggests that 80–90% of the total sulfur is usually in the form of inorganic reduced sulfur, mostly pyrite. Much of the sulfur in organic compounds is in the form of sulfides and polysulfide linkages between compounds in organic macromolecules. The sulfur in organic matter was originally considered to be mainly derived from reactions of organic materials with porewater sulfide (Nissenbaum and Kaplan, 1972; Vairavamurthy and Mopper, 1987). The identification of trisulfide linkages in kerogens (Kohnen et al., 1989) originally demonstrated that reactions with polysulfides were important. Vairavamurthy et al. (1994) showed that up to 20% of organic-S near the surface of sulfidic sediments was in the form of sulfonates, formed by the reaction between organic materials and intermediate oxidized S forms such as sulfite and thiosulfate. The relative concentrations of organic polysulfides and sulfonates are shown in Fig. 11. Sulfur is incorporated into organic matter by two main processes: (1) assimilatory sulfate reduction in organisms, especially microorganisms (see Chapter 8.1) and (2) abiotic reactions between porewater sulfur species and organic compounds. Assimilatory sulfur constitutes <1% of living biomass (Zehnder and Zinder, 1980) and may contribute 10–25% of the organic-S in sediments (Anderson and Pratt, 1995; Wakeham et al., 1997). Most of the organic-S is formed through reaction of, especially reduced sulfur, with organic compounds during diagenesis. Sulfide builds large three-dimensional network of altered biomolecules connected by mono-, di- and tri- sulfide bridges. The result is an increase in the content of macromolecular organic polymers during diagenesis (Eglinton et al., 1994). These large organic polymers with extensive sulfur cross-linking are protected from enzymatic decomposition and microbial attack and this may contribute to organic matter preservation (de Leeuw and Sinninghe Damste´, 1990; Russell et al., 2000; Lepot et al., 2009).

568

Chapter | 14 Sedimentary Sulfides FIGURE 11 Sulfur species in sediments from the Peru Margin showing the proportion of organic sulfur (as sulfonates and polysulfides) and inorganic sulfur (as pyrite and sulfate). (Modifed with permission from Vairavamurthy et al., 1994).

The timing of the incorporation of sulfur into organic matter in sulfidic sediments has been generally accepted to be early diagenetic since the original suggestion by Valisolalao et al. (1984) (e.g. Wakeham et al., 1997; Kohnen et al., 1989; Russell et al., 2000; Schouten et al., 1995; Hartgers et al., 1997). Werne et al. (2000) dated the process to
10,000 a, Kok et al. (2000) found 1000–3000 a for the process in lake sediments and Urban et al. (1999) reported significant sulfurization in less than 60 a in lake sediments. The nature of the processes involved in the formation of organic-S compounds in sulfidic sediments can be seen in the depth profiles reported from the Peru margin by Vairavamurthy et al. (1994). Fig. 11 shows that the organicS content exceeds the pyrite–S content in the top 10 cm of the sediment. However, the organic-S contents decrease with depth and are only intermittently significant, probably related to organic-C enriched layers. In the top 10 cm up to 20% of the total organic-S is in the form of sulfonates which are more oxidized than the more abundant polysulfides. Both forms of organic-S reach maxima in the top 5 cm of the sediments with sulfonates being more abundant at the sediment–water interface. The observations are consistent with the idea that organic-S forms rapidly in sulfidic sediments and that reactions between biomolecules and more oxidized S forms, such as polysulfides, sulfite and thiosulfate, are more rapid than reactions with sulfides. As we have seen, most of the microbiologically produced sulfide is oxidized at or near the sediment–water interface. The consequent redox cycling involves the

4. Solid Sulfide Phases in Sulfidic Sediments

569 FIGURE 12 Distribution of inorganic and organic sulfur phases in organic-rich estuarine sediments of St Andrews Bay, Florida (Modifed with permission from Bru¨chert, 1998).

formation of sulfur redox intermediates like polysulfides, sulfite and thiosulfate. It is in this zone that much of the organic-S appears to form. The distribution of sulfur phases in the estuarine sediments of St Andrews Bay, Florida, (Fig. 12) also shows a maximum in organic-S in the top 10 cm of the sediments. In this case, the organic-S is more conventionally classed as humic and fulvic acid sulfur and protokerogen sulfur. The soluble fulvic acid–S fraction decreases with depth and Bru¨chert (1998) found that part of the soluble fulvic–S was recycled back into the sediment porewater. A further part was sequestered in humic acids and protokerogens. Protokerogen–S increase with depth as the sulfur linkages between organic molecules form increasingly larger macromolecules. Organic-S dominates the sediment sulfur budget at or near the sediment–water interface, suggesting rapid formation of organic-S compounds. Pyrite–S begins to dominate at around 5 cm depth, but the amount of pyrite appears to be relatively constant with depth after this interval. Interestingly, AVS forms a very minor part of the sulfur budget and is less significant than the organic sulfur phases. These results are consistent with the general idea that inorganic sulfur species, such as H2S and polysulfides, may be incorporated into organic matter by reaction with functionalized lipids (especially unsaturated lipids with more than one double-bonded moiety) during the early stages of diagenesis (e.g. Francois, 1987a; Sinninge Damste´ et al., 1989; Kohnen et al., 1989; Eglinton et al., 1994; Wakeham et al., 1997; Hartgers et al., 1997; Kok et al., 2000; Werne et al., 2000, Werne et al., 2008; Russell et al., 2000; Wakeham et al., 1995; van Kaam-Peters and Sinninghe Damste´, 1997). The process involves the inter- or intra- molecular incorporation of the reduced sulfur species into low-

570

Chapter | 14 Sedimentary Sulfides

molecular weight functionalized lipids and the formation of high molecular weight abiogenic geopolymers (de Leeuw and Sinninghe Damste´, 1990). Canfield et al. (1998) suggested that the organic-S addition to normal coastal marine sediments would be 11–29 mmol g1 by extrapolation of the results of analyses of an inshore, marine sapropel. They pointed out that this is a small quantity of S compared with the typical 150–300 mmol g1 pyrite–S. Some authors have suggested that sulfurization of organic matter occurs mainly after the bulk of pyrite precipitation (e.g. Francois, 1987a; Zaback and Pratt, 1992), but others (e.g. Bru¨chert and Pratt, 1996) have shown evidence for the simultaneous formation of pyrite and organic-S, as shown in Fig. 12. It has been argued that the formation of organic-S is limited to sulfide-rich environments where the generation of sulfide is greater than the supply of reactiveFe (e.g. Berner, 1984; Raiswell et al., 1988, Raiswell et al., 1993a; Suits and Arthur, 2000; Sinninghe Damste´ and De Leeuw, 1990; Mongenot et al., 2000, Mongenot et al., 1996; Baudin et al., 1999). However, Yu¨cel et al. (2010) found that organic-S formed wherever reduced S is available in Black Sea sediments even in the presence of high concentrations of reactive-Fe. It seems from the above data that the idea of reactive Fe limiting the formation of organic-S does not appear to be valid as a generalization. OrganicS appears to form rapidly during diagenesis and reaches maximum concentrations before or coincident with substantial pyrite formation. If one of the major processes is dependent on the concentration of intermediate sulfuroxidized forms, this is perhaps not surprising. Petroleum maturation of sulfur-rich kerogens occurs at lower temperatures than sulfur-poor kerogens and sulfurization of sedimentary organic matter impacts on the kinetics of petroleum formation (Vandenbroucke and Largeau, 2007; Krein and Aizenshtat, 1994; Sinninghe Damste´ et al., 1998). Although the incorporation of polysulfide linkages into biomolecules results in polymerization and initial preservation of organic matter, the polysulfide linkages are thermally labile (Aizenshtat et al., 1995). Lewan (1998) showed that it is the formation of sulfur radicals, rather than the weakness of the S–S bonds, that causes early maturation of sulfur-rich petroleum precursors. The oxygen content of the atmosphere is determined ultimately by the burial of organic carbon (Chapter 17). The preservation of organic carbon by sulfurization thus contributes to maintaining the oxygen content of the atmosphere. It may be no coincidence that the rise in the oxygen composition of the Earth’s atmosphere c. 2.3 Ga, was coincident with the apparent increase in microbiological sulfate reduction on a global scale.

4.3. The Origins of “Anomalous” Concentrations of AVS It is quite clear that sulfide collected from sediments upon the simple addition of HCl (AVS) may be generated from a variety of mixtures of dissolved and particulate sulfide phases. However, the earlier mistaken idea that AVS was

4. Solid Sulfide Phases in Sulfidic Sediments

571

identical with metastable iron sulfides, particularly mackinawite but also greigite, which were necessary “precursors” to pyrite formation in sediments, led to many apparent contradictions. There was especially much confusion when the concentrations of AVS and pyrite were measured in real sedimentary sections (cf. Table 3). These analyses showed that the concentration of AVS was highly variable and occasionally surprisingly high. The formation of these so-called “anomalous” concentrations of AVS was an apparent problem since it should have changed to pyrite according to the paradigm current at the time. For example, in New England and Sapelo Island, Georgia, salt marshes Howarth and his associates demonstrated that even for short-term (a few hours) incubations that the majority of the reduced 35S(II) was found in the pyrite fraction not in AVS (Howarth, 1979; Howarth and Teal, 1979; Howarth and Giblin, 1983; Howarth and Merkel, 1984). This is not surprising in view of the disconnect between AVS and pyrite, of course, but at the time these observations were counter to the prevalent theory. Thode-Andersen and Jørgensen (1989) used 35SO2 4 to trace the formation of the products of sulfate reduction in coastal sediments of Denmark. The relative amounts of AVS, pyrite and elemental sulfur formed depended both on the overall sulfur chemistry and the rate of sulfate reduction. At low metabolic rates, undetectable H2S and low AVS:FeS2 (<1:20) ratios were observed and only 32–55% of the radio-labeled sulfur was recovered in the AVS extract. In sediments with high metabolic rates and high H2S concentrations, high AVS:FeS2 (>1:10) ratios occurred and a major fraction (63– 92%) of the reduced 35S was recovered in the AVS. These observations are also predictable if it is understood that AVS is not exclusively FeS. Rapid sulfide generation is likely to result in high AVS concentrations where the rate of sulfide production is greater than the rate of reaction with available Fe. However, Thode-Andersen and Jørgensen noted that isotope exchange made it impossible to directly calculate the differential rates of formation of the various reduced-S products. A related study by Fossing and Jørgensen (1990) added further details by demonstrating isotope exchange for elemental sulfur, polysulfides, S(II)T and FeS, but not with dissolved sulfide or pyrite. The isotope exchange between elemental sulfur and S(II)T was mediated by polysulfides. Exchange between elemental sulfur and FeS took place only via polysulfides and/or S(II)T. Another study using the 35SO2 4 -tracer technique to investigate AVS behavior was conducted in sediments from the Ballastplaat mudflat of the Scheldt estuary in Belgium by Panutrakul et al. (2001). They found that the predominant (70–80%) end product 35S was generally AVS, with the exception of surface sediments where 35S–pyrite and 35S–So were more dominant. The thickness of the surface layer varied seasonally. However, although most was found in the AVS fraction, pyrite was always the reduced 35SO2 4 dominant reduced (nonradiogenic) inorganic sulfide pool. In near-surface sediments (0–5 cm), the concentration of AVS was about 13% of TRS, but

572

Chapter | 14 Sedimentary Sulfides

below this interval it increased to about 30%. This is opposite to what might be expected if FeS is a precursor to pyrite formation. Interestingly, in these sediments the reactive iron concentrations were high which meant that FeSm could form. Most of the sulfide formed in the surface layers of sediments is oxidized and part of this may be reduced again leading to a cyclical process, with consequent effects on the isotopic composition of authigenic pyrite. The role of oxidants has been found to influence AVS to pyrite ratios (e.g. Aller, 1977; Thode-Andersen and Jørgensen, 1989; Oenema, 1990; Middelburg, 1991; Gagnon et al., 1995). The possible oxidants include not only oxygen, but also nitrate and iron and manganese oxides (e.g. Aller and Rude, 1988; Schippers and Jørgensen, 2002). The oxidants have two different basic functions. The first is oxidative destruction of AVS which occurs more rapidly than the oxidation of pyrite, greatly favoring pyrite accumulation relative to that of AVS. If sulfate reduction occurs at depths in the sediment, then AVS may build up leading to anomalously high AVS: pyrite ratios (e.g. Gagnon et al., 1995). Two major factors that can favor high AVS concentrations are rapid deposition rates (Goldhaber and Kaplan, 1974; Oenema, 1990; Middelburg, 1991; Gagnon et al., 1995; Wijsman et al., 2001) and low oxygen to euxinic conditions in the overlying water (e.g. Hurtgen, 1999) and for seasonally and permanent anaerobic lakes by (Marnette et al., 1993). Rapid deposition rates favor high retention efficiencies for sulfide by more quickly removing sediments from the zone where high rates of oxygen transport into the sediments by diffusion or bio-irrigation occur (Middelburg, 1991; Kostka and Luther, 1994; Gagnon et al., 1995; Lyons, 1997). Boesen and Postma (1988) and Middelburg (1991) suggested that occurrence of AVS in modern sediments may be attributed to lack of elemental sulfur or polysulfides. However, this does not seem to be the case since, in rapidly accumulating near-shore sediments where elemental sulfur could be found throughout the sediment, AVS was extracted from throughout the sediment (e.g. Wijsman et al., 2001). Mackinawite dissolves rapidly where it is undersaturated with respect to its solubility product (5.3.5) and therefore it is probable that FeSm is a dynamic component which is continually nucleating and dissolving as the relative saturation changes. The observation of this phase at any point in a sediment is not necessarily a result of its long-term preservation but may result from neoformation. In addition, the development of high dissolved sulfide concentrations in sediment reflects mainly the local reduction of sulfate and/or the diffusion of sulfide from underlying organic rich layers, as is seen in the Black Sea (2.1). AVS does not represent any specific phase or combination of phases. Thus, any general theory for its presence is likely to be inapplicable to any specific situation since the concentrations do not describe any defined or reproducible material. The concentrations of AVS in sediments are highly

4. Solid Sulfide Phases in Sulfidic Sediments

573

variable in space and time and there is no concentration which can be regarded as “anomalous”.

4.4. Direct Observations of Mackinawite and Griegite in Sediments The primary authigenic iron sulfide minerals that have been directly observed to occur in sediments undergoing early diagenesis are euhedral and framboidal pyrite, greigite, mackinawite and, rarely, marcasite and hexagonal pyrrhotite. However, except for pyrite and marcasite, these minerals are not usually observable by traditional techniques such as XRPD and scanning electron microscopy (Morse and Cornwell, 1987). A common, but highly questionable practice, has been to infer the presence of various “metastable iron sulfides” (¼mackinawite and greigite) by the occurrence of AVS and sediment color (e.g. Berner, 1964a) and by comparing ion activity products to the equilibrium solubility products of the minerals (e.g. Ba˚gander and Carman, 1994; Doyle, 1968). Other indirect techniques have included changes in sediment magnetic properties (e.g. Hilton, 1990; Hilton et al., 1986a; Neretin et al., 2004; Roberts and Turner, 1993; Blanchet et al., 2009) and Mo¨ssbauer spectroscopy (Hilton et al., 1986b; Manning and Ash, 1979; Manning et al., 1979). Direct observations, primarily by XRPD, of mackinawite and greigite in unconsolidated sediments are relatively uncommon and can be listed in a single Table (Table 4). In this listing, I have specifically excluded are observations from more ancient sediments because, as discussed in Chapter 15, the occurrence of these minerals in lithified sediments is often due to nondiagenetic processes. I have also excluded observations of intracellular iron sulfides in bacteria since, as shown in Chapter 8; its potential for substantial contributions to the sedimentary reduced sulfur pool is likely to be minor. As shown in Table 4, only a few observations of mackinawite and greigite have been reported in sediments, with many of these in freshwater bodies and marshes or in sediments overlain by anoxic to sulfidic waters such as those in the Black and Baltic seas (Table 4). Roberts et al. (2011) listed 39 occurrences of greigite in sedimentary rocks, reported during previous 30 years. Except for those listed in Table 4, all these occurrences were in older sedimentary rocks, mostly of Cenozoic age, and the presence of greigite is generally suggested by the rock magnetic characteristics. Roberts et al. (2011) point out, however, that greigite has no unique magnetic properties and that the identification of greigite by magnetic properties alone is not robust. However, even with this caveat, it is probable that many of these reported occurrences of greigite do reflect the occurrence of greigite. Interestingly most of the listed occurrences are from lake sediments and reported occurrences of greigite in marine sediments are relatively rare. Pseudohexagonal pyrrhotite has only reported once in sediments from the Sea of Japan by Kobayashi and Nomura (1972), where it appears to have

574

Chapter | 14 Sedimentary Sulfides

TABLE 4 Summary of Direct (XRD, SEM) Observations of Mackinawite, Greigite and Pyrrhotite in Unconsolidated Sediments (from Rickard and Morse, 2005) Mineral

Environment Comments

Mackinawite

River

Product of reaction with iron Berner (1962) trash in Mystic River, Boston, MA

Marsh

In sideritic marsh rock

Pye (1981)

Lake

Lake Superior

Dell (1972)

Tertiary California lakes

Skinner et al. (1964)

Greigite

Reference

East Twin Lakes near Kent, OH in Nuhfer and association with pyrite Pavlocic (1979) English Lake District

Hilton (1990)

Mid Pleistocene Southern Italy

Porreca et al. (2009)

Lake Qinghai, China

Ai et al. (2011)

Estuarine

Loch Lomond

Snowball and Thompson (1988)

Marine

Tertiary sediments

Polushkina and Sidorenko (1963)

In microfossils of recent sediments

Jedwab (1967)

In plant vacuole

Morse and Cornwell(1987)

Landsort Deep in the west-central Bo¨ttcher and Baltic Sea Lepland (2000) Mixed mackinawiteþ greigite

Marsh

Small amounts of greigite and monosulfides in siderite-rich sediment; Norfolk, England

Pye (1984)

Marine

Sulfide concretion, Black sea

Volkov (1961)

Polymineralic aggregates in Lein et al. (1980) weakly reducing sediments with little pyrite (Pacific Ocean)

formed as an overgrowth on magnetite, perhaps during earlier more reducing conditions in this sea. Mackinawite has also only been reported once in marine sediments (Lein et al., 1980) in association with greigite in mildly reducing pelagic sediments from the Pacific Ocean. The only additional marine

4. Solid Sulfide Phases in Sulfidic Sediments

575

occurrences of greigite were a single observation in a plant vacuole by Morse and Cornwell (1987) and by Bo¨ttcher and Lepland (2000) in the Landsort Deep of the brackish west-central Baltic Sea. Consequently, there appears to be little direct evidence for the common formation of mackinawite and greigite in “normal” (see Chapters 3.2 and 15.4) marine sediments. The lack of direct observations of mackinawite and greigite in sediments does not necessarily mean that they do not occur. Few published reports on sedimentary sulfides actually report any attempt to observe the solid phases. The nanoparticulate nature of precipitated mackinawite also means that direct observations by conventional means, such as XRPD and optical microscopy, are likely to be unsuccessful. What it does mean is that mackinawite or greigite cannot be assumed to be present, and basing geochemical or isotopic mass balance estimations involving AVS on this assumption introduces considerable uncertainty. It is probable that high proportions of TRS as acid volatile sulfide in sediments suggests the presence of solid sulfide phases such as mackinawite and greigite. Indeed, in several cases, these are the sedimentary systems where mackinawite and greigite are actually observed. Mackinawite and greigite are metastable phases with high solubilities relative to pyrite. Greigite formation is dependent on the prior formation of mackinawite, since greigite is formed at low temperatures in aqueous solution by the solid-state transformation of mackinawite. Mackinawite solubility (Chapter 5.3.3) is therefore the key parameter in the formation of these metastable iron sulfides in sediments. In order for mackinawite to form [Fe(II)]T and [S(II)]T must c.106 M in neutral to alkaline solutions. Note that these concentrations need to be preserved in order that mackinawite does not dissolve, since mackinawite dissolution is relatively rapid (Chapter 5.3.5). These concentrations occur commonly in inshore sediments, salt marshes and some lakes and estuaries and mackinawite has been observed in these environments. Neither of these phases changes to pyrite at low temperatures in aqueous solution. Greigite is more stable and its preservation in sediments for longer periods of time is to be expected. Mackinawite dissolves quite rapidly when its solubility product is not equaled or exceeded. This has been observed experimentally by Grimes et al. (2002) where the mackinawite precipitated in plant cells was seen to dissolve and pyrite to form (Chapter 6.6.1, Fig. 6.4). The longer-term preservation and accumulation of mackinawite in sediments is therefore dependent on the maintenance of relatively high concentrations of sulfide and dissolved iron in the interstitial waters. Such conditions may occur where sulfide is continuously being generated microbiologically at depth or where sulfide is diffusing or advecting into the mackinawite layer. Whenever the porewater total dissolved iron or sulfide concentration dips below c. 2 mM, mackinawite dissolves. Mackinawite also nucleates rapidly (Chapter 5.3.4). It is therefore probable that sedimentary FeSm is part of a dynamic system of precipitation and dissolution rather than a fixed phase, as mentioned above (14.1.9).

576

Chapter | 14 Sedimentary Sulfides

There are several reported occurrences of mackinawite and greigite in soils (e.g. Stanjek et al., 1994; Stanjek and Murad, 1994) and systems affected by human activities. These latter environments have resulted in some interesting, quasi-experimental investigations. Indeed the earliest occurrence of tetragonal FeS outside the laboratory was in the corrosion products of a sour oil pipeline in Kansas (Meyer et al., 1958). It was called “kansite”, but the International Mineralogical Association rejected the name since the environment of formation was deemed to be artificial. Even Berner’s original identification of Fe3S4 in a natural environment was on iron trash in the Mystic River where sulfide production was powered by industrial and sewage waste (Berner, 1962). More recently, Burton et al. (2011) described greigite formation in a remediated wetland environment. In this system, they were able to demonstrate that greigite formation was entirely decoupled from pyrite formation.

4.5. Distribution of S8 in Sediments The nature of solid and dissolved sulfur in sediments is variable but whether it occurs as elemental sulfur, colloidal sulfur, rhombic or crystalline sulfur it is mostly S8. Although elemental sulfur, So, is commonly referred to as the solid or dissolved species, the monomer does not occur in natural aqueous systems. One consequence of this is that when converting measured sulfur concentrations to a molar scale you need to divide by 8. The analytical problem is that sulfur with an oxidation number of 0 occurs in a number of species, including polysulfides, thiosulfate and polythionates as well as dissolved S8 and various solid forms. Knowledge of the distribution of S8 in sediments has been constrained by the need to develop protocols that can distinguish S8 from S(0) in other species, as well as more robust analytical methods that might be deployed on ships or in the field. The introduction of the original cyanolysis protocol, where it is used to extract S(0) and the reactant complex determined spectrophotometrically (Bartlett and Skoog, 1954) provided the basis for a potential environmental method. However, it was subsequently found that the original protocol did not extract microbiologically produced colloidal sulfur (Janssen et al., 1999). Kamyshny et al. (2009) stated that cyanolysis reported less than 1% of crystalline S8 but all of the colloidal S8. The accuracy of the listing in Table 5 of sedimentary S8 concentrations collected by this method is therefore uncertain. Reaction of S8 and polysulfides with sulfite leads to the formation of thiosulfate which can then be analyzed by a variety of methods: iodine titration (Szekeres, 1974), voltammetry (Luther et al., 1985) or HPLC (Fahey and Newton, 1987). The sulfur contents reported with this method in Table 5 are likely to be maximum concentrations. The introduction of reversed phase high performance liquid chromatography (RP-HPLC, Mo¨ckel, 1984a,b,c) has led to a robust and popular analytical

4. Solid Sulfide Phases in Sulfidic Sediments

577

protocol for analyzing sedimentary S8. The sulfur is extracted with an organic solvent, ideally methanol (Zopfi et al., 2004), before being analyzed. Kamyshny et al. (2009) showed that the method was relatively accurate with recoveries of colloidal and crystalline S8 usually 80% depending on the solvent used. The problem with the method is that it does not permit accurate estimations if S8 in the presence of other S(0)-bearing species. The S8 concentrations listed in Table 2 using RP-HPLC are therefore likely to be maximum values. The methodology has been subsequently improved (Kamyshny, 2009) and integrated into a general protocol for the quantitative analysis of zero valent sulfur compounds in natural systems (Kamyshny et al., 2009). This protocol also may distinguish between colloidal and crystalline sulfur. However, analysis of Kamyshyny et al.’s (2009) survey of analytical methods for sulfur shows that precisions of 30% relative would be excellent but, because of the cumulative nature of the errors in extraction and analyses, 50% is more probable with many analyses at lower precision than this. Although use of the same methodology in any given analytical campaign will improve the relative values of the analyses, the random error on the extractions can still be very large. Table 5 lists the solid sulfur content of some sediments. Remember that the concentrations are given in mmol (g dry wt)1 and need to be divided by 8 to get the molar concentration of S8. The highest S8-S concentrations are reported from salt marsh sediments where visible sulfur is accumulating. Similar sediments in brackish water systems have low S8-S contents, as expected. The concentration of S8-S in marine sediments varies up to 100 mmol (g dry wt)1 in the suboxic and euxinic sediments of the Black Sea. In coastal waters subject to variable oxygenated conditions in the overlying water the S8-S concentration is normally of the order of 10 mmol (g dry wt)1. In marine sediments overlain by normal oxygenated seawater, the concentration is usually around a magnitude lower. Even with the large analytical uncertainties indicated above, the differences in the average S8-S contents of these environments appear to be robust. There appear to be very few generic conclusions to be gleaned from S8-S profiles through sediments. Fig. 13 shows data from similar environments but different locations, analytical protocols and sampling expeditions in the Black Sea. The Black Sea is an interesting test site since different environments can be sampled in the same time interval. In Fig. 13, sediment profiles under oxic, suboxic and euxinic water are compared. The oxic water contained 213 mmol L1 O2 (Zopfi et al., 2004). Yu¨cel et al. (2010) reported that O2 penetrated 8 mm into the top 1 cm, fluff layer of their core. Both profiles show small amounts of S8-S variably distributed with depth. The suboxic water contained no detectable O2 (Yu¨cel et al., 2010) or <5 mmol L1 (Zopfi et al., 2004) and no detectable S(II). Zopfi et al’s profile showed virtually no S8-S whereas Yu¨cel et al. found large amountsdup to 52 mmol (g dry wt)1. The varying results probably reflect the range of environments described as suboxic. The euxinic environment is the classic and much studied Unit 1 of the central Black Sea basin. Various

578

Chapter | 14 Sedimentary Sulfides

TABLE 5 Sulfur Content of Marine Sediments. Table Modified and Extended from Thamdrup et al. (1993) and Zopfi et al. (2004). Units of mmol cm3 are equated to mmol (g dry wt)1 on the basis that the porosity w 50% (cf. Yu¨cel et al., 2010, p. 365) and the density of the dried sediment w 2

Environment Location

S8eS concentration Method mmol g1

Transitional

6.2

Normal

Texel, Netherlands

Reference

UV Visscher and Gemerden, spectroscopy 1993

Kalo Lagoon, 6.8 Denmark

Cyanolysis

Thode-Andersen and Jørgensen 1989

Aarhus Bay, Denmark

10e17

Cyanolysis

Troelsen and Jørgensen, 1982

Black Sea


2

Weser Estuary, 0.8 Germany

Wijsman et al., 2001 RP-HPLC

Canfield and Thamdrup, 1996

Aarhus Bay, Denmark


4.6

Cyanolysis

Thode-Andersen and Jørgensen 1989; Thamdrup et al., 1994; Moeslund et al.?

North Sea

2.2

Cyanolysis

Sørensen and Jørgensen, 1987

Kattegat

0.5

Cyanolysis

Sørensen and Jørgensen, 1987

Arabian Sea


4

RP-HPLC

Passier et al., 1997

Skagerrak

0.2

Cyanolysis

Sørensen and Jørgensen, 1987

Saguenay Fjord Eastern Pacific, Peru upwelling

UV Gagnon et al., 1996 spectroscopy
1.6

Eastern 0.5e15 Pacific, Chile upwelling Black Sea


1.2

Cyanolysis

Fossing, 1990

RP-HPLC

Ferdelman et al., 1997; Thamdrup and Canfield, 1996; Zopfi et al., 2008

RP-HPLC

Zopfi et al., 2004; Yu¨cel et al., 2010

579

5. Dissolved Species in Sedimentary Sulfide Systems

TABLE 5 Sulfur Content of Marine Sediments. Table Modified and Extended from Thamdrup et al. (1993) and Zopfi et al. (2004). Units of mmol cm3 are equated to mmol (g dry wt)1 on the basis that the porosity w 50% (cf. Yu¨cel et al., 2010, p. 365) and the density of the dried sediment w 2dCont’d S8eS concentration Method mmol g1

Reference

Mid-Atlantic Bight

0.04e13.7

RP-HPLC/ sulfitosis

Ferdelman, 2004

Suboxic

Black Sea

20e100

RP-HPLC

Yu¨cel et al., 2010

Euxinic

Gotland Deep 6.4

Cyanolysis

Podgorsek and Imhoff, 1999

Arkona Deep 0.34

Cyanolysis

Podgorsek and Imhoff, 1999

Black Sea

1.1

RP-HPLC

Zopfi et al., 2004

Black Sea


100

RP-HPLC

Yu¨cel et al., 2010

Tyro Basin

2.5

RP-HPLC

Henneke et al., 1997

Bannock Basin 4.8

RP-HPLC

Henneke et al., 1997

Environment Location

measurements of the sulfide content of the overlying waters have been reported but Zopfi et al. (2004) reported 75 mmol S(II) L1 and Yu¨cel et al. (2010) 400 mmol S(II) L1 overlying the sediment profiles in Fig. 13. Both profiles show that S8-S steadily decreases from the sediment–water interface. The S8-S values on the profiles reflect the sulfide contents of the overlying water. Yu¨cel et al. (2010) sampled down to 60 cm depth in the anoxic basin where Unit 1 was replaced by a turbidite layer. The high S8-S contents of this layer (
82.8 mmol (g dry wt)1) settled down to a low constant value (within analytical error) in the underlying sediments at <2 mmol (g dry wt)1. It may be that this sort of level represents the background level S8-S content of sulfidic fine-grained sediments.

5. DISSOLVED SPECIES IN SEDIMENTARY SULFIDE SYSTEMS Porewater may constitute over 90% of the volume of a surface sediment, decreasing downward to less than 20% at the onset of lithification. Much of the porewater is therefore expressed upward during compaction through fractures, burrows and the sediment pore spaces. This upward flow is accompanied by a series of reactions as the dissolved components of the pore continually react

580

Chapter | 14 Sedimentary Sulfides

FIGURE 13 Sediment S8-S profiles from different environments in the Black Sea. Top row: measurements by Zopfi et al. (2004). Bottom row measurements by Yu¨cel et al. (2010). The Zopfi et al. measurements have been recalculated from mmol cm3 (see Table 2).

and re-react with solid phases in the sediment and respond to changes in the redox environment. Porewater composition is extremely sensitive to these reactions compared with solid phases. For example, as we have seen above, the disappearance of FeSm with depth in anoxic sediments is mainly caused by dissolution where the concentrations of dissolved sulfide and iron become less than the IAP for FeSm. In a sediment with 0.1 wt % FeSm and 80% porosity, c. 80 mmol L1 sulfide is released to the porewater per percentage FeSm dissolved. This is a large amount relative to the porewater sulfide concentrations in sulfidic sediments but the dissolution of a few percent FeSm might be difficult to detect. The problems with the interpretations of porewater sulfide concentrations are not only those of analysis described above but also intrinsic to a system where sulfide is also being produced through microbial sulfate reduction in the sediment. Even in euxinic systems, the concentrations of sulfide in the porewater of surface sediments are usually greater than those in the overlying

5. Dissolved Species in Sedimentary Sulfide Systems

581

waters and a net flux of dissolved sulfide diffuses from the sediment into the overlying water column. In oxic environments, the same process is fundamental to the flux of sulfide to the surface sediment layers where it is oxidized by microorganisms. These results in most of the sulfide produced in sediments being re-oxidized and partially recycled (Jørgensen, 1977). Porewater sulfide profiles have been widely used as a basis for quantitatively measuring reaction kinetics and stoichiometries in sulfide sediments. However, these models are usually one-dimensional and assume steady state conditions. More recent investigations of sulfidic sediments, described in this volume, have shown that steady state conditions are not the normal situation. Aller et al. (2010) has termed this feature of sulfidic sediments as unsteady diagenesis and this appears to be the normal situation. Aller himself in his PhD thesis (1977) initially introduced nonsteady state processes into the sedimentary sulfide system when he documented the importance of bioturbation and this was further developed by Berner and Westrich (1985).

5.1. Total Dissolved Sulfide in Sediment Porewaters As noted in Chapter 13 the concentration of dissolved S(II) species in sulfidic sediments can be considerable and constitute an environmental hazard. For example, the concentrations of dissolved S(II) in the surface layer of the Walvis Bay sediments commonly reaches millimeter levels and Bru¨chert et al. (2009) reported 20 mM. The classic depth profiles of porewater sulfide in sediments in oxic and anoxic environments are shown in Fig. 14. These profiles were collected using conventional filtrations and wet chemical analyses and thus may not only report dissolved H2S and HS concentrations. Even so, they are likely to be comparable with similar relative errors. In normal oxic environments the concentration of sulfide is in the micromolar range. This means that dissolved sulfate concentrations are not significantly affected by sulfide formation and are relatively constant with depth. The sulfide concentrations increase with depth suggesting a net upward diffusion gradient (Fig. 14 inset). In the anoxic environment, the dissolved sulfate rapidly disappears in the sediment porewaters and the dissolved sulfide is some 1000 times more concentrated. Again there is a marked depth gradient reflecting upward diffusion.

5.2. Dissolved Sulfide in Anoxic Sediments Depth profiles of dissolved sulfide in anoxic sediment porewaters vary considerably even within the same sedimentary unit. The classic area is the Black Sea which has been studied intensively although the number of porewater sulfide profiles is still limited compared to the area involved. The original description of Unit 1 undisturbed, laminated sediments by Degens and Ross (1972) has been shown to be limited to deep central areas of the basin (Mitropolsky et al., 1982

582

Chapter | 14 Sedimentary Sulfides

FIGURE 14 Classic profiles of total dissolved sulfide (H2S) and sulfate (SO2 4 ) in porewater of sediments beneath oxic and anoxic waters. Inset is the expanded profile of the H2S contents in oxic water environments with a mM scale (i.e. 1000  the mM scale used in the main plots). These data are from Effingham Inlet, a fjord off Vancouver Island, and were collected by Hurtgen et al. (1999) using conventional analytical methods. (Reprinted with permission from Hurtgen et al., 1999 Ó American Journal of Science).

cited in Yu¨cel et al., 2010). All other anoxic sediments in the Black Sea contain turbidites. This is important since the turbidite layers vary considerably in porewater chemistry and this in turn leads to various porewater sulfide profiles even within the same sedimentary units of this single basin (Fig. 15). Of the six profiles shown in Fig. 15, only one (Fig. 15f), shows the increasing downward curve which might be assumed to be a normal sediment distribution. The porewater sulfide concentrations are affected by sediment provenance and composition as well as the sulfide production in underlying layers. For example Fig. 14b and f suggest that sulfide is being produced in an underlying deeper, organic-rich layer. The implications of this variability in porewater sulfide profiles in nearsurface sediments are considerable. They show that porewater sulfide is part of a dynamic system rather than a single stage, constant system which may be modeled by simple mathematical algorithms. The profiles record both differential production (through microbial sulfate reduction) and destruction (through microbial sulfide oxidation and abiologic Fe and Mn oxidation) as well as variable transport of sulfide within the sediment, mainly by diffusion. Bioturbation in anoxic Black Sea sediments appears to be less significant than in oxic environments and this adds a further dimension to the dynamics of sulfide biogeochemistry in normal, oxic marine and freshwater environments.

583

5. Dissolved Species in Sedimentary Sulfide Systems

(a)

0

100

Sulfide, µM 200 300

400

(b)

0

–20 Water

20 Blackish homogeneous 40

Fluff Dark Grey to Black Soupy

80

Depth, mm

120

60 80

160 200 240

100 120

Greyish homogeneous

140

Dark Grey Sandy

280

“Old” Fluff

320 Light Grey Clayish

360

160

400 Sulfide, µM 400 500 600

(c)

300 –20

(d) 100

Fluff

Fluff

Depth, mm

Depth, mm

120

“Old” Fluff

140 Grey

“Old” Fluff

180 220 Grey

40

“Old” Fluff Light Grey Sandy “Old” Fluff

260 300

380 0

100

80 Light Grey

Sulfide, µM 200 300 400

500

(f)

60

Water

Sulfide, µM 200 400

500

Water

Fluff

0

Fluff

Blackish homogeneous

20 40 Depth, mm

80 Depth, mm

0

–20

20 40

Dark-to-Black

60

Light Grey Clayish

340

0

500

0

100

–20

Sulfide, µM 300 400

Water

Dark Grey

60

200

–20

Water

20

(e)

400

40

Fluff

Depth, mm

Sulfide, µM 200 300

Water

0

0

100

100 120 140

60

Dark Grey to Black Soupy

80

160 100

180 200 220 240

“Old” Fluff Black homogeneous

120 140

Light Grey Clayish

FIGURE 15 Variations in porewater dissolved sulfide profiles in anoxic Black Sea sediments. (Reprinted with permission from Konavalov et al., 2007).

584

Chapter | 14 Sedimentary Sulfides

From the geochemical point-of-view, the observations suggest that the sulfide molecules in any ancient sedimentary rock reflect a complex history which may not be directly related to the biogeochemistry of contemporaneous overlying water column.

5.3. Dissolved Sulfide in Suboxic Sedimentary Systems Revsbech et al. (1983) used pH, O2 and S(II) microelectrodes in a laboratory study of incubated cores from a microbial mat from Solar Lake, Sinai. They found that the O2 and S(II) zones overlapped in the mats. The deployment of voltammetric solid state electrodes resolved the problem of the simultaneous measurement of O2 and S(II) in natural aqueous and sedimentary systems. Brendel and Luther (1995) originally tested the voltammetric approach on sediment cores and showed an overlap of the S(II) and O2 zones. However, Luther et al. (1998) showed a distinct separation of the sulfidic and oxygenated zones from sediments from the Canadian continental shelf and slope. The development of a suboxic zone in sediments appears to be a widespread phenomenon. It must be recalled, however, that the definition of suboxic is entirely operational. It refers to a zone where dissolved O2 and S(II) are not detected. In the best case, in situ measurements with voltammetric microelectrodes, this means an O2 concentration of less than 3 mmol L1 and an S(II) concentration of less than 0.03 mmol L1. Although this means that chemically significant amounts of dissolved O2 may still be present in the suboxic zone, it is likely that dissolved O2 is not the dominant oxidant in these environments. A further caveat is the microbiological perspective described in Chapter 10, where sulfate-reducing microorganisms are widely detected in or very near aerobic parts of the system. There are three alternative situations in sedimentary environments (1) the sulfidic zone overlaps the oxic zone (e.g. Fig. 16) (2) the sulfidic zone is separated from the oxic zone by a suboxic region where neither S(II) nor O2 is detected (e.g. Fig. 17) and (3) the sulfide zone is in contact with the oxic zone. The relative distribution of the three alternative situations is currently unknown through insufficient analyses. In addition, the limited data that are available suggest that the three situations are not mutually exclusive but may develop at different times in the same location. Jørgensen and Nelson (2004) described the suboxic zone existing in finegrained marine sediments where Fe and Mn oxy(hydro)oxides form a reactive barrier which binds H2S (Fig. 18). The kinetics of the reaction between S(II) an FeIII oxyhydroxides are relatively rapid. They lead to the direct formation of S(0) and Fe(II) species, which may react with dissolved S(II) to produce iron sulfides. Manganese sulfides are relatively soluble, so the analogous reaction with MnIV oxyhydroxides produces S(0) and Mn(II) species. In modern sediments, deeper penetration of O2 can occur in coarse grained clastic deposits. For example, in the Wash in the UK, the sulfide zone may be 1 m below the

5. Dissolved Species in Sedimentary Sulfide Systems

585

FIGURE 16 Overlapping S(II) and O2 zones in a microbial mat sediment from the Guaymas Basin. The overlapping zone moves up and down through the sediment diurnally (From data in Gundersen et al., 1992).

FIGURE 17 Distinct separation of S(II) (C) and O2 (B) zones in Canadian shelf sediments showing a welldeveloped suboxic zone (From data in Luther et al., 1998).

586

Chapter | 14 Sedimentary Sulfides

FIGURE 18 Mn, Fe and S(II) and O2 profiles in the top 10 cm of a typical fine-grained marine sediment. The shaded area indicates deeper black layers in the sediment where S(II) dominates below the uppermost brownish layers with little H2S or O2. Note that O2 disappears at some distance from the H2S zone and Fe(III) oxyhydroxides continue into this zone. In this study, particulate Mn was assumed to be Mn(IV) (From Thamdrup et al. (1994) and Jørgensen and Nelson (2004)).

sediment–water interface. This phenomenon has been partly credited to current-induced advective porewater transport (Huettel et al., 1998). In modern sediments, bioturbation and sea grass roots (Blaabjerg et al., 1998) can introduce O2 into deeper levels. However, for most of geologic history, burrowing organisms were not present and bioturbation was restricted. The implication is that banded or layered rocks, analogous to modern undisturbed varved sediments, should be more abundant in ancient (e.g. Precambrian) rock sequences.

5.4. Dissolved Fe Species The composition of the dissolved iron species associated with sulfidic sediments is not as well-constrained. However, it is of considerable importance if calculation of saturation states of porewaters with respect to iron sulfide minerals is to be attempted. In low sulfidation environments (i.e. environments with either abundant reactive Fe or limited reducible S) [Feaq]T > [Saq]T, and aqueous Fe species are likely to be important. Fe(II) forms a series of weak ion pairs or complexes with most anions. It is improbable that Fe(II) hydroxyl complexes make any substantial contribution to AVS. As shown in Chapter 3, FeOHþ is a weak complex which, even if there is some uncertainty on the value of pK(FeOHþ), it is unlikely to contribute significantly to the total dissolved Fe(II) activity. Fe(OH) 3 is

587

5. Dissolved Species in Sedimentary Sulfide Systems

a stronger complex but unlikely to be significant in the pH range of AVS. The FeII carbonate, chloride and sulfate complexes are also relatively weak and will make little contribution to the total dissolved Fe(II) of AVS. In sulfidic environments, inorganic dissolved Fe(III) species are not significant. I conclude that the dominant nonsulfide, inorganic Fe species in sulfidic sediments is hexaaqua Fe(II) (e.g. Turner et al., 1981). As discussed in Chapter 4.3, aqueous iron–sulfide complexes play a potentially important role in the sulfidic sediments. The independent measurements for the stability of [Fe(SH)]þ show a degree of congruency which suggests some confidence that [Fe(SH)]þ may be a significant component of porewaters in contact with FeSm. The stability of further bisulfide complexes, such as [Fe(SH)2]0 and [Fe(SH)3], is controversial. However, in high sulfidation environments (i.e. environments with either abundant reducible S or limited reactive Fe) [S(II)]T  [Fe(II)]T, and [Fe(SH)]þ is the dominant Fe(II) bisulfide complex.

5.5. Aqueous FeS Clusters Aqueous FeS clusters make up a significant fraction of the AVS of many natural aqueous and sedimentary environments, as shown in Chapter 4 (Table 6). The significance of FeSaq in sulfidic sediments is not limited to its potential relationship with mackinawite. The presence of FeSaq provides

TABLE 6 “Dissolved” S8eS Contents of Sediment Porewaters Environment Marine

Salt Marsh

Concentration Location mmol S L1 Method

Reference

Rehoboth 1.45e15.8 Bay

Solid state electrode voltammetry

Saguenay fjord

5.5e21

UV spectroscopy Gagnon et al. (1996)

Eastern Pacific, Chile

4e59

RP-HPLC

MidAtlantic Bight


19


555

Rozan et al. (2000)

Ferdelman et al. (1997); Thamdrup and Canfield, (1996); Zopfi (2004) Ferdelman (1994)

RP-HPLC/ sulfitosis

Ferdelman (1994)

588

Chapter | 14 Sedimentary Sulfides

a means for transporting Fe(II) in S(II)-rich systems. The final fixation of Fe and S in pyrite in sedimentary systems requires that both Fe and S are transported to the site of pyrite formation (Raiswell et al., 1993b). In many sulfidic sediments, FeSaq is the only iron sulfide phase that is actually observed apart from pyrite. Work in the Cardiff laboratory has managed to obtain FeSaq solutions without any visible precipitated FeSm. These solutions are clear and not black. So it would seem that FeSaq does not cause the black color in sulfidic sediments. Obviously, in view of its key role in transport, solid phase formation and ecology, FeSaq constitutes a key research target.

5.6. Dissolved Organic Complexes of Fe and S The role of organic complexes and colloids in the bioinorganic chemistry of sulfidic environments is not well understood. As noted in Chapter 3.1.2, much progress has been made on identifying the organic ligands responsible for complexing Fe in oxic environments. However, the significance and chemistry of complexation of Fe(II) in reduced natural systems remains unknown. Luther et al. (1992) examined organic ligands in salt marsh sediments and concluded that organic ligands were involved in seasonal iron cycling. They suggested that multidentate organic chelates containing oxygen, such as carboxylate ions, complex Fe(II).

5.7. Dissolved Polysulfides Understanding the distribution of polysulfides in aqueous environmental systems has been limited by sampling methods. The S(II) system is highly sensitive to oxidation and the absolute exclusion of air during the sampling and analytical procedures is not entirely possible. The result is that most ex situ analyses of environmental polysulfides are probably on the high side, due to artifactual polysulfide production during sampling and analysis. The development of in situ analytical methods (Brendel and Luther, 1995; Rozan et al., 2000) has provided a more accurate insight into the distribution of polysulfides in natural aqueous systems (e.g. Luther et al., 2001). The presence of polysulfides in estuarine sediments was limited to a thin transition zone between sediments with S(0) dominant and the deeper AVS zone (Rozan et al., 2000). The Sx(II) concentration was 13.9 mM or about 20% of the S(II) concentration. Some more general implications were derived from examining deep hydrothermal vent systems (Rozan et al., 2000). Here polysulfides occurred in diffuse flow regions where trace O2 was present. The Sx(II) concentration was estimated to be 0.27 mM or around 5% of the S(II). The implication of these observations is that polysulfides are likely to be present in more oxidized zones in sediments where trace O2 is a possible constituent. It would imply that

5. Dissolved Species in Sedimentary Sulfide Systems

589

polysulfides are unlikely to make up a substantial proportion of the total dissolved sulfur source for AVS.

5.8. Sulfur Oxyanions Although sulfur oxyanions are not a direct source of sulfide in sediments, they are discussed briefly here because a large number of sulfur oxyanions in which sulfur has oxidation numbers between S(0) and S(VI) have been reported from the natural environment. The sulfur microbial ecosystem associated with sulfidic natural aqueous systems is complex and produces a spectrum of these intermediate oxidized sulfur compounds. The most important of these is probably thiosulfate which disproportionates to S(0) and S(VI). Various schemes have been devised to identify these species (El Dein et al., 1989; Leck and Ba˚gander, 1988) in natural systems. However, as with polysulfides, ex situ methods risk artifactual formation of oxidized sulfur species. In situ electrochemical analyses have demonstrated that they do not normally contribute significantly to the total AVS concentration. In situ measurements by Luther et al. (2001) showed the presence of thiosulfate, elemental sulfur and tetrathionate in salt marsh microbial mats, subtidal sediments and hydrothermal vent diffuse flow waters. In microbial mats they related S(II) oxidation to microbiologically-mediated processes: in sediments and diffuse vent flow they concluded that Fe(III) phases were the direct oxidants.

5.9. Dissolved S8 As noted above solid S8 is in some degree of equilibrium with aqueous S8 with dissolved concentrations expected to be
6 nmol S8 L1 in water at 4  C. Electrochemical methods have been used to measure aqueous S8. The advantage of these methods is that the aqueous S8 content of sediment porewater can be probed in situ. This reduces the problem of filtration methods not removing nanoparticulate S8 or S8 being formed artifactually during the process of porewater extraction. Boulegue et al. (1982) originally used Ag/Ag2S electrodes to detect polysulfides in sediments. Although Luther et al. (1985) and Wang et al. (1998) demonstrated that polarographic and voltammetric methods could be used to detect polysulfides and analyse elemental sulfur in porewaters, the operation of a hanging mercury drop electrode on board ship was not always straightforward. Wang and Tessier (2009) used laboratory-based voltammetric methods to determine total S(0) in sediment porewaters from lakes and then calculated the amount of dissolved S8 using Kamnyshy’s (2004) polysulfide equilibrium data (see Chapter 2.5). The equilibrium between S8, S(II) and Sn(II) species is extremely rapid (Kamyshny et al., 2003) and the use of the equilibrium assumption is justifiable. The introduction of the solid state Au–Hg electrode into environmental

590

Chapter | 14 Sedimentary Sulfides

studies enabled the direct measurement of sulfur species in sediment porewaters (Luther et al., 2001a,b; Rozan et al., 2000). The protocol distinguishes total Sn(II), S8 and HS. Table 6 shows that far more apparent aqueous S8 is measured in sediment porewaters than is permitted by S8 solubility in inorganic systems. Although it is possible that this is a disequilibrium system, the presence of solid S8 means that the kinetic inhibitor affects the crystal growth part of the precipitation process rather than the nucleation stage (where such gross supersaturations are commonplace). The other possibility is that the particle size of sulfur sols can be extremely small. This has been alluded to by Kamyshny et al. (2009) and practically demonstrated by Wang and Tessier (2009) where they showed that the amount of apparent aqueous S8 is lower when the solution is passed through a 0.01 mm filter compared with a 0.2 mm filter. Most likely the bulk of the dissolved S(0) measured is actually contained in polysulfides.

6. DIAGENETIC MODELING OF SEDIMENTARY SULFIDE SYSTEMS Extensive attempts to model of the sedimentary sulfur system have been made in order to further understand how the system works and to make quantitative assessments which permit predictions (and retrodictions for ancient sedimentary systems) to be made. Most of these models have been based on the one-dimensional steady model for sedimentary sulfide of Berner (1964b) which was adapted for sedimentary metal sulfides by Rickard (1973) and is summarized in Fig. 19. The key to this model was Berner’s insight that if the S redox boundary in the sediment rises at the same rate as the sedimentation, the dynamics of the system could be

FIGURE 19 Elements of the original one-dimensional, steady state model for sulfidic sediments (Berner, 1964; Rickard, 1973).

6. Diagenetic Modeling of Sedimentary Sulfide Systems

591

described by steady state equations. That is, the rate of change of any property, p, of sediment with time, t, and depth, x, may be described by vp dp vp  ¼ s (3) vt x dt vx t where s is the sedimentation rate. This equation could be used to solve for metabolizable C for the boundary conditions assuming that organic matter metabolization was independent of the oxidant and only dependent on the concentration of organic C. X ¼ 0 C ¼ C0 and X/N C/0 For steady state conditions, the solution to Eqn (3) is an exponential function of the rate of utilization of the metabolizable C and depth. The change in metabolizable carbon is then related to sulfide production as a simple first order reaction. The model had a number of intrinsic assumptions: (1) sulfide is formed entirely as a result of microbial sulfate reduction within a reduced S zone; (2) sulfide production is nowhere limited by sulfate concentration; (3) the rate of utilization of metabolizable C is first order and is all used by sulfate-reducers; (4) chemical reactions involved in the model are fast and do not perturb the steady state model (5) In this model the transport of sulfide is through diffusion and in described by Fick’s first and second laws. The result of the steady state assumption is a differential equation which can solved analytically. These models give a good description of quiescent sedimentary environments with minimal bioturbation such as in the Californian basins where the model was developed. The reason that these models work is that the assumptions made are relatively good approximations to the real system. For example, the assumption that sulfate-reduction can be related to organic C is reasonable where sulfate reduction is a major microbial process and that the organisms utilize many of the products of other microbes in the complex ecology. Since the early studies by Berner and Rickard, further work on sedimentary sulfides revealed a number of other processes at work in the sulfur cycle. An attempt to synthesize current knowledge about processes in the sedimentary sulfur cycle is shown in Fig. 20. These processes have been described and discussed at some detail in this book. They include the influence of methanogenesis, microbial iron reduction, aqueous iron sulfide clusters, the role of FeSm, redox recycling of sulfur, the role of Mn in sulfide oxidation, anoxic sulfide oxidation, organic S chemistry, biochemistry of sulfur reduction and oxidation, microbial ecology and advances in understanding sulfur isotope systematic. In tandem with these, there have been considerable

592

Chapter | 14 Sedimentary Sulfides

FIGURE 20 A modern version of the sedimentary sulfide cycle (Reprinted with permission From Rickard and Luther, 2007 Ó American Chemical Society). For a color version of this figure, the reader is referred to the online version of this book.

advances in analytical methods of analyzing sulfidic sediments, such as multiple sulfur isotopes, nontraditional stable isotopes and microelectrode voltammetry. Time dependence was introduced to the original Berner model by Lasaga and Holland (1976) and more complex treatments of organic matter mineralization by Berner (1974) and Jørgensen (1978). Boudreau (1984) introduced oxidant dependence and this was further developed by Rabouille and Gaillard (1991a, 1991b) and Boudreau and Canfield (1993). More dynamic, unsteady diagenetic models, including seasonality, bioturbation and porewater irrigation have been developed by several workers (e.g. Aller et al., 2010; Goldhaber, 2004; Jørgensen et al., 2004; Aller, 1980b). Temperature dependence was originally included by Westrich and Berner (1988) using a simple Arrheniustype temperature relationship. The original models assumed steady state conditions. This leads to elegant mathematics since, in steady state conditions, the simultaneous differential equations that describe individual processes can be solved analytically for various boundary conditions. Indeed, it would be possible to include all the processes included in Fig. 20 as a vast array of simultaneous differential equations and obtain solutions. However, it has become more apparent that steady state diagenesis, even as an approximation to the real world, is rare or at least a special case. For example, in the Black Sea, a model environment for euxinic

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systems, the extent of areas of quiescent sedimentation is limited and the sedimentary sequence is frequently interrupted by turbidites (see Chapter 13). A large number of biogeochemical cycles, if not all of them in a holistic Earth System approach, are coupled and interdependent. The sedimentary sulfur cycle is related to the global cycles of the key biologic elements C, N, O, P and the important redox elements, Fe and Mn. Boudreau (1996), Wang and Van Capellen (1996), Van Capellen and Wang (1996) and Soetaert et al. (1996) independently published numerical, public–domain, coupled reaction-transport models aimed primarily at solving nutrient and organic matter during early diagenesis. These models necessarily include a number of simplifying assumptions and depend on the parameterization of transport and reaction processes. However, variants of these models have been applied, with varying success, to a wide variety of diagenetic environments. Boudreau (1997, p. 361) ended his classic textbook on diagenetic modeling with the admonition to diagenetic modelers states that they should “not take the results your models too seriously”. The models, by definition, are approximations to reality and, as we have seen, the reality is complicated and by no means entirely understood. They are powerful tools to investigate specific processes during diagenesis but do not necessarily entirely describe these systems.

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Chapter | 14 Sedimentary Sulfides

Zopfi, J., Ferdelman, T.G., Fossing, H., 2004. Distribution and fate of sulfur intermediates – sulfite, tetrathionate, thiosulfate and elemental sulfur – in marine sediments. In: Amend, J.P., Edwards, J.E., Lyons, T.W. (Eds.), 2004. Sulfur Biogeochemistry – Past and Present, vol. 379. The Geological Society of America, Boulder, CO, pp. 97–116. Zopfi, 2000. deleted. Zopfi, J., Bo¨ttcher, M.E., Jørgensen, B.B., 2008. Biogeochemistry of sulfur and iron in Thioplocacolonized surface sediments in the upwelling area off central chile. Geochim. Cosmochim. Acta 72, 827–843.