Sedimentary Pyrite

Chapter 6

Sedimentary Pyrite

Chapter Outline 1. Introduction 2. Pyrite Structure 3. Pyrite Composition 4. Pyrite Solubility 5. Sedimentary Pyrite Formation 5.1. Syntheses of Pyrite under Low-temperature Conditions 5.2. Pyrite Nucleation and Crystal Growth 6. Kinetics and Mechanisms of Sedimentary Pyrite Formation 6.1. Mackinawite, FeSm, as a Reactant

233 234 235 235 240

241 246

247

6.2. Sulfide-controlled Reaction Kinetics 6.3. Polysulfide-controlled Pyrite Reaction Kinetics 6.4. Relative Rates of Pyrite Formation 6.5. Kinetics of Pyrite Nucleation and Crystal Growth 7. Sedimentary Pyrite Textures 7.1. Single Crystals 7.2. Framboids 7.3. Spheroidal and Nodular Pyrite References

255 257 260

263 265 265 268 275 278

249

Pyritologia: or a History of the Pyrites, the Principal Body of the Mineral Kingdom. Johann Friedrich Henckel, 1757. A. Milar, London (English translation), 376 pp.

1. INTRODUCTION As indicated by the quotation heading this chapter, the abundance of pyrite on the Earth’s surface has been long known and has led to pyrite being the target of many pioneering investigations. Thus, pyrite was the first mineral structure determined in 1914 by Bragg (Bragg, 1914) with his new X-ray diffraction system. Charles Hatchett (1804) showed that the composition of pyrite was FeS2 as part of a debate on the Law of Definite Proportions. William Pepys, Developments in Sedimentology. http://dx.doi.org/10.1016/B978-0-444-52989-3.00006-4 Copyright Ó 2012 Elsevier B.V. All rights reserved.

233

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Sedimentary Pyrite

a distant relation of the diarist, reported he had inadvertently synthesized pyrite serendipitously when a mouse got into a laboratory jar of ferrous sulfate (Pepys, 1811). He concluded that the pyrite could form by the reaction of the iron sulfate with “animal matter” and that the sulfate was “entirely deoxygenated” or reduced as we would say today. By 1838 this was recorded in a standard geology textbook (Bakewell, 1838). Allen et al. (1912) synthesized pyrite hydrothermally in 1912. Sugawara et al. (1953, 1954) first reported the importance of pyrite in marine sediments, and Berner’s (1964a, 1964b, 1970) and Rickard’s (1968a, 1968b) work in the 1960’s led to the key role of pyrite in the sedimentary system being fully appreciated.

2. PYRITE STRUCTURE The structure of pyrite (Fig. 1) is well-known. Pyrite is an iron (II) disulfide with a NaCl-type structure. The S
II 2 groups are situated at the cube center and the midpoints of cube edges, and the low-spin FeII atoms (d6, t62g ) are located at the corners and face centers. The arrangement of the disulfide dumbbells is such that the structure, although cubic, has a relatively low symmetry, space group Pa3. The structure has threefold axes along <111> directions and twofold axes along <100> directions. The twofold symmetry means that the [100], [010], and [001] zone axes (equivalent to the a, b, and c crystallographic axes) are not crystallographically interchangeable with each other by a simple

FIGURE 1 The structural elements of pyrite. (Reprinted with permission from Rickard, D. and Luther, G.W., 2007 Ó American Chemical Society).

4. Pyrite Solubility

235

90 rotation as in simple cubes. One result of this structure is that pyrite, along with several other minerals, exhibits chirality. Thus Guevremont et al. (1998) demonstrated that there are significant differences in the sensitivity of pyrite to oxidation of the (100) and (111) planes. This chirality of pyrite has been theoretically exploited in the involvement of pyrite in the adsorption of organic molecules and, consequently, in prebiotic syntheses implicated in the origins of life (Wa¨chtersha¨user, 1998). However, this idea was challenged by PontesBuarque et al. (2000) who argued that Stern-layer modulation of surface charge, acetate adsorptive behavior and the requirement for divalent cations in the attachment of organic key molecules, make a chiral-discriminator character of pyrite unlikely in this context. Bither et al. (1968) first presented a molecular orbital (MO) interpretation of the pyrite structure which has been further described by Tossell et al. (1981), Luther (1987) and Rickard et al. (1995). The MO and frontier molecular orbital (FMO) calculations have been shown to have significant implications for a fundamental explanation of pyrite properties and in predicting both bulk and surface reactions (Luther, 1990; Kornicker and Morse, 1991; Schoonen et al., 1999; Drzaic et al., 1984).

3. PYRITE COMPOSITION Kullerud and Yoder (1959) originally suggested that the composition of pure pyrite is stoichiometric FeS2. They concluded that deviations from stoichiometry were caused by analytical uncertainties or the presence of traces of other elements in the material. Ellmer and Ho¨pfner (1997) used theoretical arguments, and a critical review of stoichiometry measurements has shown that pyrite has a very narrow homogeneity range (<1&). Thomas et al. (1998) concluded that pyrite is a stoichiometric semiconductor with a homogeneity range <<0.5 atomic %. Pyrite displays both p-type and n-type semi-conductivity (Agaev and Emujazov, 1963) which results from trace amounts of other elements in the structure. For example, Oertel et al. (1999) were able to synthesis n-type pyrite by doping pyrite with 0.3 atomic% Co. Pyrite analyses commonly show the presence of trace and minor elements and this is to be expected in view of the solid solution ranges possible between pyrite and other disulfide minerals. The trace element spectra of pyrites in sediments have been the subject of innumerable investigations with a view to using these as environmental indicators. Unfortunately, it is a common PhD topic suggestion and students are advised to steer well clear of accepting this subject for their PhD studies. The problems are discussed in Chapter 13.7.8.

4. PYRITE SOLUBILITY The solubility of pyrite in water at ambient temperatures is very low as expected for a Fe(II) low-spin t62g electron configuration. Indeed, it is not

236

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Sedimentary Pyrite

directly measurable experimentally. Pyrite solubility data are derived from heat capacity measurements of the formation of pyrite from its elements at higher temperatures. The equilibrium solubility product of pyrite, Klsp;pyrite in aqueous solutions is usually given by the Eqns (1), (2) and (3): FeS2 þ Hþ ¼ Fe2þ þ HS
þ S0 þ




H þ HS

¼ H2 S

FeS2 þ 2Hþ ¼ Fe2þ þ H2 S þ S0

K ¼ 10
14:2

(1)

K ¼ 10

(2)

7

K ¼ 10
7:2

(3)

Rickard and Luther (2007) revisited published values for the solubility of pyrite and recalculated its solubility in the light of more recent thermodynamic data. They based their calculations on Ksp (pyrite) (25  C, I ¼ 0) ¼ 10
16.4 (Emerson et al., 1983), which was recalculated from data collated by Robie et al. (1978) which, in turn, were based on original measurements by Grønvold and Westrum (1962) and Toulmin and Barton (1964). These measurements are independent of the errors in the earlier National Bureau of Standards (NBS) value of DGof (Fe2þ) (see Chapter 3.1.1) since they were derived for reaction (4): Fe0 þ S02 ðgÞ ¼ FeS2

(4)

Toulmin and Barton had used the Stull and Sinke (1956) value for DGof of 19.13 kcal mol
1 which compares closely with the Cox et al. (1989) value of 79.7  0.3 kJ mol
1 recommended by Nordstrom and Munoz (1994). However, the Emerson et al. value for Klsp;pyrite is affected by the erroneous NBS DGof (Fe2þ) value. Correcting for this leads to Klsp;pyrite ¼ 10
14:3 for reaction (1). At pH ¼ 7, the {H2S}{Fe2þ} product in the S0 stability field is 10
21.3. The congruent dissolution of pyrite involves the disulfide ion, S
2 2 . Previously, the stabilities of S2(
II) species were only imperfectly understood, and therefore the congruent dissolution of pyrite: (S02 (g))

FeS2p ¼ Fe2þ þ S2
2

(5)

has not generally been used to describe pyrite solubility. However, Harmandas et al. (1998) suggested Ksp (pyrite) ¼ 8.51  10
26 for this reaction. This assumes an equilibrium constant for the reaction þ Sð0Þ þ H2 S ¼ S2
2 þ 2H

(6)

of 10
15.6, which compares with 10
18.5 according to the Kamyshny et al. (2004) polysulfide data set. Recalculation according to the Kamyshny polysulfide data set gives 10
24 for the congruent dissolution of pyrite Eqn (5). The relative stability of S2(
II) species according to the Kamyshny data set has interesting implications regarding the thermodynamics of pyrite formation in low-temperature aqueous systems like marine sediments and anoxic basins.

237

4. Pyrite Solubility

The dominant polysulfide species in the pH range of most natural waters is HS
2 and this is the third most abundant sulfide species in this region after HS
and H2S. We can therefore consider the reaction FeS2 þ Hþ ¼ Fe2þ þ HS
2

(7)

for which log K ¼
13.84. This suggests that at pH ¼ 7, log {Fe } 2þ
{HS
2 } ¼
20.84 and that at log {Fe } ¼
9, log {HS2 } ¼
11.84. That is, 2þ
in nanomolar Fe concentrations, just picomolar HS2 is required to precipi
tate pyrite. We can relate {HS
2 } to {HS }, for example, directly through 2þ

HS
þ S0 ¼ HS
2

(8)


for which log K ¼
1.76. This suggests that log {HS } in equilibrium with picomolar HS
2 is
10.24. That is less than 1 nM of dissolved sulfide is necessary to precipitate pyrite in solutions with nanomolar dissolved Fe. The implications of these very small quantities of sulfur species required to precipitate pyrite are considerable in both experimental and natural systems. In an experimental system with high total sulfide, at least at the millimolar level, 0 the purity of the reagent with respect to HS
2 specifically, but S species generally, would need to be greater than analytical grade usually available to investigators. The sulfide reagent would need to contain less than 6.5 ppm S(0) species. Another way of looking at it is to consider a 100% pure sulfide reagent and consider how much of an electron acceptor, like Fe3þ or O2, would be required to produced the minimum amount of HS
2 required to precipitate pyrite. In the case of a Fe2þ reagent, the limiting concentration of Fe3þ required would be determined by the stoichiometry of equilibria like
(9) 2Fe3þ þ H2 S ¼ 2Fe2þ þ S 0 þ 2H That is, the Fe3þ:S(0) ratio is 2:1. This implies that the Fe2þ reagent used would need to contain less than 10 ppm in order for the disulfide not to be present in sufficient quantities for pyrite to precipitate. As a codicil it might be noted that the Fe3þ/Fe2þ ratio in any solution at equilibrium is determined by: Fe3þ þ e
¼ Fe2þ

(10)

for which Nordstrom and Munoz (1994) recommend log K ¼ 13. Simple inspection suggests that, at all Eh values above water breakdown, enough Fe3þ will be present at equilibrium in an Fe2þ solution to accept sufficient electrons from S(
II) to produce enough S(0) to precipitate pyrite. A similar calculation can be made for oxygen. Experimental oxygen control is limited by the lower limit of quantitative analysis of O2(g), since measurements of dissolved O2 in sulfide solutions are currently constrained by technical problems. The most precise O2(g) controls in sulfide experimentation have been achieved in the Cardiff laboratory, with O2(g) levels maintained at less than 1 ppmv, the lower limit of measurement. As can be seen from the

238

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Sedimentary Pyrite

calculations with respect to Fe3þ, this level of O2(g) is easily sufficient to produce enough S(0) to precipitate pyrite under all reported experimental Fe and S(
II) concentrations. Unpublished work in the Cardiff laboratory examined pyrite formation in the presence of <1 ppmv to 100 ppmv O2 and found no increase in pyrite formation. The role of O2 may be limited to (a) high concentrations in conditions where S(0) forms from S(
II) oxidation, and that it is the presence of S(0) that enhances pyrite syntheses through kinetic, rather than equilibrium thermodynamic, effects and (b) creating a surface O2 layer on sulfide minerals, such as pyrite, greigite and mackinawite, which produces sufficient hS
2 2 to surpass the pyrite solubility product at the surface. In the natural environment, of course, there are adequate quantities of electron acceptors to produce the minimum HS
2 concentrations required for pyrite formation, even in the low Fe contents of normal marine systems. From the equilibrium thermodynamic viewpoint, there is no problem with the observation that pyrite forms in most Earth surface environments, including sedimentary systems. Pyrite is the stable phase in systems with even submicromolar concentrations of S(
II) since (a) its solubility product is so low and (b) the relatively high stability of S2(
II) means that the pyrite solubility product is exceeded in all reasonable experimental and natural Fe- and S(
II)containing environments. In fact, pyrite has a significant stability even in more oxic environments where {SO2
4 } > {S(
II)}, since the boundary between oxic and anoxic (or suboxic) systems is usually taken to be where {SO2
4 } ¼ {S(
II)}. S(
II) species, and therefore pyrite, continue to have a significant activity into the SO2
4 dominated region. The resultant pH-Eh diagram for pyrite stability in sediments, is shown in Fig. 2. Pyrite has an extensive stability region over pH ¼ 2–10, both above and below the SO2
4 /S(
II) boundary which is often taken as marking the upper limit of “reduced” systems. Including the measured stability data for Fe(II)polysulfide complexes (Chadwell et al., 1999, 2001) suggests that the lower boundary of pyrite stability is marked by the FeS04 or FeS05 complexes. If both these complexes are suppressed, then the region is occupied by pyrite. That is, pyrrhotite and troilite have no stability regions at 25  C in aqueous solution relative to pyrite. This is consistent with these minerals being rare to absent in marine sediments. Rickard and Luther (2007) considered the relative stability of pyrite in the presence of the Green Rusts, mixed FeIIFeIII oxyhydroxides including a variety of anions such as sulfate. At seawater sulfate concentrations they showed that upper boundary of pyrite stability could be limited by the stability of sulfate Green Rust II. However, the thermodynamic data for the Green Rusts is uncertain and the lack of robust stability data for some Fe hydroxide complexes makes the Green Rust stability zone in sediments with respect to pyrite less certain. In soils, where the Green Rusts have been observed, they may have an effect on the upper stability limit for pyrite. The pyrite stability field extends into the area where {SO2
4 }  {S(
II)}, which is conventionally described as “oxidized” by environmental geochemists.

0.2 {ΣS}=10–3

{ΣS}=10–6 0.1

0.1 rhombic sulfur

0

SO4–

–0.1 –0.2

Eh (volts)

Eh (volts)

0

H2S 10–15 Pyrite 10–12

–0.3 –0.4

10–9 –0.5 4

4. Pyrite Solubility

0.2

6

8 pH

rhombic sulfur SO4–

–0.1 –0.2

H2S(aq) Pyrite 10–2

–0.3 –0.4

HS–

HS– 10

–0.5 4

6

8

10

pH

FIGURE 2 Pyrite stability in seawater at 25  C, 100 kPa total pressure and {S}T ¼ 10
3 and 10
6, {Fe}T between 10
15 and 10
9 (Calculated with thermodynamic data listed in Rickard and Luther, 2006).

239

240

Chapter | 6

Sedimentary Pyrite

Near to this redox boundary, the calculated ionic activity product (IAP) in equilibrium with pyrite changes extremely rapidly. Butler and Rickard (2000) noted that the change could be as much as 1014 within 50 mV. The effect is significant in that supersaturations with respect to pyrite vary over very small regions of pH-Eh space in this region. Since, as discussed below, spontaneous pyrite nucleation requires extremely large supersaturations, this IAP gradient may affect whether pyrite nucleates rapidly or not in a particular sedimentary environment. The results are that (1) pyrite formation varies over extremely small distances in a sedimentary or aquatic environment due to local spatial heterogeneities and (2) that metastable phases, such as mackinawite and greigite, may be preserved for geologically significant time periods, especially within the lower part of the pyrite stability field where pyrite IAPs are lower and vary less in pH-Eh space.

5. SEDIMENTARY PYRITE FORMATION It is instructive to explore what is meant exactly by the “formation” of pyrite, or of any other mineral for that matter. Formation is divided into two processes: nucleation and crystal growth. In practice the overwhelming mass of the material is normally a result of crystal growth; crystal nuclei are exceedingly small and contribute little to the bulk of the material produced. In other words, when a sedimentary sample or experimental reaction product is analyzed, the pyrite process which is being interrogated is normally crystal growth. Berner set the ball rolling with his 1970 paper entitled Sedimentary Pyrite Formation. Berner was well aware of the importance of the separate processes of nucleation and crystal growth and provided some of the original chemical sedimentological theory for these processes in his 1971 textbook. Berner (1970) proposed that sedimentary pyrite formed by the reaction FeSm þ Sð0Þ ¼ FeS2p

(11)

Rickard (1975) pointed out that reaction (11) could not describe the molecular mechanism of pyrite formation since S(0) is in the form of S8, which would make this an impossible multimolecular reaction step. Berner himself was quite aware of the distinction and viewed his equation as summarizing an overall process. However, it did give rise to the mistaken view by some investigators that sedimentary pyrite was formed through a solid-state transformation since, conventionally, FeSm and S(0) are solids. As noted in Chapter 14.5.4, the concentration of dissolved S(0), as a discrete aqueous molecule, is extremely low in sedimentary environments. By contrast, polysulfides consist of mixtures of S(
II) and S(0) (see Chapter 2) and the sulfur in pyrite is in the form of S2
2 , which is a mixture of 50% S(
II) and 50% S(0). Butler et al. (2004) demonstrated isotopically that the sulfur in pyrite was derived entirely

5. Sedimentary Pyrite Formation

241

from the reactant polysulfide, even with FeSm as a reactant. So the overall reaction (11) is better written as: FeSm þ Sn ð
IIÞ ¼ FeS2 þ Sðn
1Þ ð
IIÞ

(12)

This suggests that, in terms of mass balance, the net process involves the apparent addition of S(0) to FeSm as Berner (1970) had observed. However, in actuality, the pyrite-S is entirely acquired from the polysulfide reactant and not from the FeSm. In its most extreme form, reinterpretations of the Berner process led some investigators to suppose that sedimentary-pyrite formation proceeds through a series of solid-state transformations involving mackinawite and greigite successively. In fact there is no evidence for this process and it is highly improbable chemically considering the chemistry and structure of the mackinawite and greigite. Likewise, the balanced reaction (Wilkin and Barnes, 1996): 2FeSm þ 1=2H2 O þ 3=4O2 ¼ FeS2p þ FeOOH

(13)

is not a mechanism because it involves 3.25 molecules which makes it statistically improbable and, anyway, the fractions of molecules as written are, of course, mechanistically impossible. Wilkin and Barnes (1996) were pointing out that, in the absence of excess sulfur, FeSm will dissolve and, since the solubility of FeS2p is far lower than that of FeSm, FeS2p will be produced. The involvement of phases such as O2 and FeOOH are merely there as suggested electron donors and acceptors in order to provide an overall reaction balance. Again the reaction describes a possible mass balance at equilibrium in the presence of FeSm as a reactant.

5.1. Syntheses of Pyrite under Low-temperature Conditions Laboratory syntheses of pyrite have proved a powerful tool for investigating pyrite properties, since Allen et al. (1912) reported the first systematic study. However, interpretations of the results in terms of the mechanism of sedimentary-pyrite formation are limited by simple apparent mass balance, the lack of definition of reactants and products and the difficulties in controlling and characterizing the experimental conditions. Table 1 summarizes experimental syntheses of pyrite. Work prior to 1935 was reviewed in some detail by Mellor (1935) and further listings were reported by Schoonen and Barnes (1991c) and Wei and Osseo-Asare (1995). This summary is restricted to experiments typically <100  C and in aqueous solutions. Because of the interest in pyrite in material science, as a possible solar-cell material for example, there are a large number of published reports of syntheses in nonaqueous solvents and/or at high temperatures. The reactions described in Table 1 need to be treated with caution and the original reports should be read in detail. One problem is that FeS2 also

242

Chapter | 6

Sedimentary Pyrite

TABLE 1 Pyrite Recipes: Reported Abiologic Syntheses of Pyrite in Aqueous Solutions Typically at <100  C and 100 kPa Total Pressure Fe S reactants reactants a

Fe(II)

H2S

T ( C) i

20 e95

Comment

Reference

3e5

Open to air

Berner, 1964b

w 25

7e8

Farrand, 1970

65

6.9, 7.9

Berner, 1970

25ie60

?.

Sweeney and Kaplan, 1973

NaHS þ S8

65i

1.5e8.8

Schoonen and Barnes, 1991c

Na2Sng

25i

4.4e9.5

Rickard, 1968a

H2S þ Na2S4

75i

2.4e7.2

Murowchick and Barnes, 1986

NaHS þ Na2S4 25i

7.3e7.6

Schoonen and Barnes, 1991b

5.5e8

Luther, 1991

3.4e6.9

Murowchick and Barnes, 1986

5.5

Luther, 1991

i

H2S þ S8

Na2S2, Na2S4, Na2S5g

25i,100

H2S þ Na2S2O3 75i Fe IIIb

pH

Na2S5h

25i

NaHS, Na2S

3.6, 6.5

H2S

25i

5.5

FeCO3

Cystine

100

?

FeOOHc

H2S

20e25i

4

Berner, 1964b

25i

3.8e6.5

Roberts et al., 1969

i

25

4.4e7.0

Rickard, 1968a

25i

6e8

Rickard, 1974

?

7e8.5

Pyzik and Sommer, 1981

H2S þ S8

w 25i

6.5e7.5

Kribek, 1975

S8

65

7

60e85

?

25ie50

6e8

Rickard, 1975

6e8

Wilkin and Barnes, 1996

Na2S

FeSm

d

H2S þ S8 H2S

i

70

25e125

Roberts et al., 1969 þ Humic acids

Lambert, 1973

Berner, 1969 Open to air

Sweeney and Kaplan, 1973

Rickard, 1997

243

5. Sedimentary Pyrite Formation

TABLE 1 Pyrite Recipes: Reported Abiologic Syntheses of Pyrite in Aqueous Solutions Typically at <100  C and 100 kPa Total Pressuredcont’d Fe S reactants reactants

T ( C)

pH

Comment

100

ox FeSme

Roberts et al., 1969, Taylor et al., 1979

60e100

6

Wilkin and Barnes, 1996

70i

6e8

70i H2S þ RSH, RSSR, sulfonate

6e8

Wilkin and Barnes, 1996

H2S þ Na2SO3, 70i Na2S2O3

6e8

Wilkin and Barnes, 1996

Open to air

g

Fe1
xS Fe3S4g

FeS2p

a

Wilkin and Barnes, 1996

Na2Sn

25i

7

Rickard, 1968a

H2S

35e160

?

Taylor et al., 1979

70 FeStf

Reference

Wilkin and Barnes, 1996

H2S

100

?

Taylor et al., 1979

H2S

100

w7

Drobner et al., 1990

i

70

6e8

Wilkin and Barnes, 1996

70i

6e8

H2S þ Na2SO3, 70i Na2S2O3

6e8

Wilkin and Barnes, 1996

H2S

70i

7.71

Wilkin and Barnes, 1996

NaHS

25i

6.5

RSH, RSSR, sulfonate

Open to air

Wilkin and Barnes, 1996

Undersaturated Harmandas et al., with respect to 1998 FeSm

As FeSO4, FeCl2 or Fe(NH4)2(SO4)2. As FeCl3 or Fe(NO3)3. c As a-FeOOH or unspecified Fe(III) (oxy)hydroxide. d Precipitated FeS is assumed to be FeSm. e FeSm oxidized in air before reaction with unspecified composition. f Troilite. g Commercial unspecified pyrrhotite. h S(
II) present, see text. i Reaction temperature. In all other cases the reactants were mixed at room temperature and heated to the noted temperature. Modified from Rickard and Luther (2007). b

244

Chapter | 6

Sedimentary Pyrite

describes the pyrite dimorph, marcasite, and the synthesis of pyrite needs to include XRD confirmation of the product. Furthermore, the product is rarely 100 wt% pyrite, but also contains various amounts of reactants and other Fe sulfide products. In reality the various Fe salts used appear to be largely irrelevant: Fe(II) and Fe(III) are mostly dissociated in these systems. Experimentally, the use of Mohr’s salt, Fe(NH4)2(SO4)2, is recommended since this is more resistant to oxidation than simple Fe(II) sulfates, chlorides and nitrates. The reasons for this are unknown. Except for the case of Harmandas et al. (1998), all the reactions with Fe salts were supersaturated with respect to FeSm and, because of the kinetics, FeSm can be assumed to have been precipitateddeven when this is not reported. The experimental systems used are overwhelmingly of the batch reactor type and the reactant concentrations are therefore necessarily in the millimolar to molar rangedoften much greater than would normally be encountered in natural systems. The reaction processes therefore often involve the reaction of a dissolved FeS species and/or nanoparticulate FeSm with the solution. In order to overcome this, a series of reactions have been carried out using variously defined FeS precipitates as reactantsdrather than the less welldefined in situ precipitate. Obviously, for well-defined experiments, especially kinetic studies, the FeS reactant needs to be both well-characterized and reproducible. Freeze-dried FeSm has been used as a defined reactant (Rickard, 1997), for example, and the process of freeze-drying has been shown to arrest the development of FeSm (Wolthers et al., 2003). The development of this material continues when it is placed back in solution. As discussed in Chapter 5, the reactant FeSm is sensitive to oxidation and changes with time. Opening systems to air or using apparently oxidized FeSm also makes the reactant materials difficult to define. Indeed, it appears from Chapter 5.4.4 that all FeSm precipitates in aqueous solutions have at least a monolayer of oxygen on the surface with concentrations up to ca 15 atom %. However, the bulk oxygen concentration in the FeSm precipitate is below detection limits (Rickard et al., 2006). We actually synthesized FeSm in nonaqueous solvents and found no difference in the composition of this material. Thus oxygen cannot contribute to the overall redox mass balance of pyrite formation with FeSm as a reactant, although it may affect the local kinetics of pyrite nucleation. One of the problems in the literature has been the use of less well-defined FeSm reactants which were basically partially oxidized, and then drawing generic, rather than empirical, conclusions about the reaction. The Cardiff laboratory has contributed to overcoming this problem by sending synthetic samples of well-defined reactants to other laboratories worldwide on request. One of the effects of the prior precipitation of FeSm is that the actual reactant concentrations of dissolved Fe and S(
II) are not easily controlled. A similar discussion is applicable to the sulfur reagents. H2S is defined in Table 1, where this is introduced as a gas, but it should be noted that this normally

5. Sedimentary Pyrite Formation

245

required a counterion to maintain solution pH and this is generally Naþ. Thus there is little chemical difference between this and Na2S or NaHS, which are basically solutions of H2S in NaOH and are mainly dissociated. In the case of polysulfide reactants, rapid equilibration occurs in solution providing a spectrum of polysulfide stoichiometry, even if a well-defined polysulfide reagent is used in the first place. The best that can be attained is dominance of a particular species under specific physicochemical conditions, as noted above. In each case, the reaction between Fe(II) and polysulfides appears to have been oversaturated with respect to FeSm and this is therefore precipitated rapidly in the experiments. The minimum S(
II) concentrations in rapidly equilibrated aqueous polysulfide systems occur in the presence of excess S8. However, even here the dominant dissolved S species are S(
II), as discussed above. Although there are several reports that Fe2þ reacts directly with polysulfides to form FeSm and S(0) (e.g. Allen et al., 1912; Berner, 1970; Luther, 1991; Schoonen and Barnes, 1991a) all these experimental systems were supersaturated with respect to FeSm and thus this material was necessarily produced as a product. Harmandas et al. (1998) approached the reaction in systems which were undersaturated with respect to FeSm, like most sedimentary systems, and showed that pyrite was the product of the reaction. Mechanistically it has been demonstrated that pyrite formation in lowtemperature aqueous solution involves the formation of a dissolved [FeS] transition intermediate (Rickard, 1975; Luther, 1991; Rickard and Luther, 1997) where the {Fe(II)}{S(
II)} IAP is greater or equal to the stability constant for the complex or cluster. The suppression of such intermediates can inhibit pyrite formation (Rickard et al., 2001). This raises an immediate experimental problem in that Rickard et al. (2001) showed that trace amounts of aldehydic carbonyl were sufficient to suppress pyrite formation. This means that the reaction systems need to be very clean with respect to aldehydic carbonyls, which are widespread in a number of key biochemical pathways; it also implies that other trace organics might also have similar effects. The possible effects of trace contaminants, including microorganisms, in low temperature, experimental sulfide systems have not been investigated. It might help to explain some of the contradictory and often irreproducible results that have been reported in pyrite syntheses. It would seem certain to have an effect on the distribution of pyrite in natural systems. Even so, experimentally the main problem appears to be the control of oxidation. The maximum O2 content measured in the Cardiff laboratory for commercial analytical-grade O2-free N2 gas was 17 ppmv. Obviously, this was from a bad batch, but even blowing N2 gas with lower ppmv-level O2 into sulfidic reaction vessels can provide sufficient oxidant to change the system. In the Cardiff laboratory all reactions have been carried out in anoxic chambers under an inert gas atmosphere which had been scrubbed for O2 removal on Zr powder in a hightemperature furnace. A further problem, as mentioned above, is the analysis of O2 at low levels; conventional methods even in the absence of S compounds have

246

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Sedimentary Pyrite

detection limits of >1 ppm which can constitute a substantial amount of O2 in these systems. It is possible that the O2 contents in natural marine and sedimentary anoxic systems are lower than that routinely attainable in the laboratory, since microorganisms are likely to remove all O2 to an extremely low level. The listed pH of the experiments is merely a guide. In most cases it is the final pH and this is often reached within a few hours. The problem is that the reactions generally tend to produce relatively large quantities of acid, and buffering in batch systems is normally impractical since at the concentrations required the buffer itself reacts with the Fe and S reactants. In most cases the reactants were added at room temperature and then heated to the temperature listed. Since the precipitation of FeSm is fast, this means that a FeSm precipitate was initially formed in the reaction vessel at room temperature, and the subsequent, ill defined, mixture was the one involved in the reaction. Of course, the initial reactant could include Fe3S4g or FeS2p as well as FeSm. Where the reaction as a whole was carried out at the listed temperature, this is noted. The reaction products have been quantitatively analyzed in very few cases indeed, mostly in the kinetic investigations (e.g. Rickard, 1975, 1997; Harmandas et al., 1998). The solid products are usually identified qualitatively by powder X-ray diffraction, for example, and the product solutions are rarely analyzed at all. Unfortunately, these poorly defined experimental syntheses have formed the basis for widespread discussions of pyrite formation in sedimentary environments. Even if the reactants and products are analyzed, measurements of their concentrations provide very little information about the processes involved in pyrite formation. In order to do this, it is necessary to observe the products and reactants over time and probe the reaction kinetics.

5.2. Pyrite Nucleation and Crystal Growth Pyrite formation involves two distinct physical processes, nucleation and crystal growth. Whereas pyrite crystal growth appears relatively fast, nucleation of pyrite can be slow and potentially rate-limiting (Schoonen and Barnes, 1991b). Harmandas et al. (1998) investigated this in experiments which approached pyrite formation from undersaturation with respect to FeSm and below the FeS0 stability region. The supersaturation limit, U*, is defined as the supersaturation level up to which a phase can be expected not to precipitate spontaneously. The supersaturation limit can be regarded as the supersaturation level at which pyrite nucleation is relatively fast and not rate-limiting. Harmandas et al. measured Upyrite in the presence of pyrite seeds to be 5.7  1014 for reaction 5. This result is important since it means that at Upyrite < 1014 the rate of pyrite nucleation determines the rate of pyrite formation. At Upyrite > 1014 the rate of pyrite crystal growth is the rate-limiting process and this feature was exploited by Harmandas et al. in their experimental investigation of pyrite crystal-growth kinetics.

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As with many insoluble salts, pyrite nucleation is kinetically hindered and high supersaturations are required for nucleation to occur. The presence of pyrite seeds in the system may enhance the nucleation kinetics and thus the supersaturation limit proposed by Harmandas et al. probably approaches the minimum value. At 25  C, the solubility of FeSm written in terms of {Fe2þ} {H2S} at pH ¼ 7 is 10
10.5. The solubility of pyrite written in terms of {Fe2þ} {H2S} at pH ¼ 7 is 10
21.2. The supersaturation with respect to pyrite at pH ¼ 7 is thus 1011.7. However, in neutral to alkaline systems, the dominant Fe species in equilibrium with FeSm is FeS0. At pH ¼ 7, in the presence of FeSm, {Fe2þ} {H2S} ¼ 10
14.9, which is close to Upyrite ¼ 5.7  1014 required for pyrite to nucleate spontaneously. The result suggests that systems which contain FeSm in pure water are close to the pyrite supersaturation limit, and small supersaturations with respect to FeSm will thus cause pyrite to nucleate spontaneously. In other words, the presence of FeSm in experimental and environmental pyriteforming systems is neither happenstance nor because of some requirement for FeSm as a precursor in pyrite formation. The reason is that at the supersaturations required to initiate rapid pyrite nucleation, the system will tend to be saturated with respect to FeSm.

6. KINETICS AND MECHANISMS OF SEDIMENTARY PYRITE FORMATION An interesting and important question in sedimentary geochemistry is: how does pyrite form in sediments? The above discussion has demonstrated that pyrite can be synthesized through a number of routes but many of these do not seem to be likely in most sedimentary systems because, for example, the reactants are scarce in natural environments or, more commonly, their experimental concentrations are far higher than those usually observed naturally. The equilibrium thermodynamics of pyrite formation are consistent with the observations of the widespread distribution of pyrite in sediments. However, the fact that thermodynamically there is enough Sn(
II) present in all sulfidic sediments to account for pyrite formation says nothing about the mechanism of pyrite formationdnor, indeed, of how and if it will occur. The confusion in the geochemical literature is often expressed in terms of writing balanced equations for pyrite formation, often based on partial analyses of laboratory syntheses, and then claiming that these represent alternative pyrite-forming processes or mechanisms. None of these balanced reactions that have been proposed are wrong in the sense that, ultimately, at equilibrium, the product pyrite will be formed. And all of them could approximate to mechanisms (if unlikely) if the rates of reaction were such that equilibrium was attained rapidly. However, they do not actually describe molecular mechanisms. There is no a priori reason to suggest that all sedimentary pyrite forms through a single mechanism. This is obvious if we consider high-temperature anhydrous syntheses of pyrite through the reaction of FeS in its various forms

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with S8 vapor or the relatively rapid, anhydrous, solid-state equilibration of Fe3S4g to pyrite and pyrrhotite at temperatures above c. 200  C. The mechanism of a reaction is a molecular process normally acquired from studies of the kinetics of pyrite formation and the derived rate law. The rate law itself describes the slowest molecular step in the process. The rate-controlling molecular mechanism in a chemically controlled reaction formally involves a transition state complex. However, in experimental studies of heterogeneous kinetics such transition state complexes may not be formally accessible or defined. In the case of pyrite formation, the kinetic evidence described below demonstrates that an aqueous [FeS] reaction intermediate may be involved in a rate-controlling reaction. However, as has been repeatedly demonstrated (Luther, 1991; Rickard, 1997; Rickard and Luther, 1997) that this [FeS] reaction intermediate is not equivalent to FeSm. An alternate way of looking at the problem of sedimentary-pyrite formation is to consider pyrite crystal growth directly. Pyrite crystals grow quite readily in environments where reactants like FeSm are not present and the growth mechanism is likely to be a more direct process. Basically, the principle of microscopic reversibility suggests that the fundamental reaction for pyrite nucleation and growth is simply the opposite of the congruent dissolution reaction. That is: Fe2þ þ S2
2 ¼ FeS2p

(14)

The problem then devolves into defining rate-controlling processes that produce the reactants Fe2þ and S2
2 . Much of the discussion in the literature centers on the origin of the disulfide. This may be formed in solution, through the reaction between S(
II) species and S(0) for example, or at a solid surface. As shown above, sulfide solutions will have a significant concentration of S2(
II) species through equilibrium reactions and this can be increased in the presence of S(0). Furthermore, mixtures of aqueous polysulfides equilibrate rapidly (Kamyshny et al., 2003) and therefore any S2(
II) reacting to form pyrite is rapidly reformed. The problem with the origin of the Fe2þ ion in pyrite is more subtle. Fe2þ may be produced from the dissolution of FeSm or Fe3S4g but, as indicated below, there is a large number of other iron compounds which can give rise to 2þ has rarely been implicated Fe2þ. By contrast with S2
2 , the production of Fe in chemical rate-limiting steps in sedimentary-pyrite formation although, as discussed in Chapter 3, the availability of reactive iron in sediments may limit pyrite production. Three mechanisms for pyrite formation have been established, one involving Fe(II) and polysulfide (Rickard, 1975; Luther, 1991), one involving FeSaq and S(
II) (Rickard, 1997; Rickard and Luther, 1997) and a third mechanism for pyrite crystal growth (Harmandas et al., 1998). As pointed out by Rickard and Morse (2005), a virtually infinite number of theoretically balanced reactions could be proposed, depending on the reactants and

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products chosen. All these proposed reactions involve the end member molecular mechanisms.

6.1. Mackinawite, FeSm, as a Reactant As shown in Table 1, a number of Fe reactants have been proposed in pyrite formation. However, FeSm has been singled out for special mention because (a) it is normally produced initially in experimental syntheses of pyrite and (b) it was formerly believed to be widespread in sediments (Volkov, 1961; Berner, 1970; Jørgensen, 1977; Boesen and Postma, 1988; Schoonen and Barnes, 1991a; Canfield et al., 1992; Wilkin and Barnes, 1996; Hurtgen et al., 1999; Benning et al., 2000; Neretin et al., 2004). Rickard and Morse (2005) discussed this in some detail and noted that the idea may have partly arisen because of the assumption that HCl-soluble sulfide in modern sediments is equivalent to FeSm (Aller, 1977). As shown above, FeSm is involved as a reactant in most experimental syntheses of pyrite because of the demands of the experimental design. However, FeSm is not generally observed to form in normal marine sediments because the Fe concentration in seawater is usually so low. In many publications (Volkov, 1961; Roberts et al., 1969; Berner, 1969, 1970; Farrand, 1970; Lambert, 1973; Kribek, 1975; Jørgensen, 1977; Taylor et al., 1979; Boesen and Postma, 1988; Luther, 1991; Schoonen and Barnes, 1991c; Canfield et al., 1992; Wilkin and Barnes, 1996; Lyons, 1997; Hurtgen et al., 1999; Benning et al., 2000; Neretin et al., 2004) mackinawite is referred to as a “precursor” mineral. Of course, FeSm could be a “precursor” in the sense that it forms far more rapidly than pyrite. However, it is not a “precursor” in the sense that it is a necessary prerequisite for pyrite formation. In fact, there is no reason to suppose that FeSm is any more of a “precursor” mineral than FeOOH, Fe2O3, Fe(OH)3, Green Rust 2, fougerite or any other Fe phase which may react with sulfide to form pyrite. Certainly, it has been unequivocally demonstrated experimentally and in natural systems that FeSm does not “transform to pyrite” in the sense of a solidstate transformation. This is shown in Fig. 3, where the pyrite product in a reaction where FeSm is the source of the reactant Fe2þ, is situated at a distance from the FeSm. That is, the Fe(II) component of the pyrite must have traveled in solution from the FeSm solid reactant source as a dissolved species before reacting with S2(
II) to form pyrite. The dissolution of FeSm in the formation of pyrite has also been neatly demonstrated in a series of experiments on the pyritization of plant cells (Rickard et al., 2007). The celery used contains both well-defined, open ended, tubular xylem cells and closed, more equidimensional parenchyma cells. Initially, FeSm is precipitated in the open xylem cells but is unable to penetrate the closed parenchyma cells (Fig. 4a, b). After reaction with H2S, the FeSm in the xylem cells dissolves and pyrite is formed in the parenchyma cells (Fig. 4b). That is, FeSm is not transformed to pyrite via some solid-state reactants. Rather,

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FIGURE 3 Pyrite crystals formed from the reaction of aqueous Fe(II) with H2S. Note that the crystals form well away from FeSm and nucleate on each other in a cascade arrangement (arrowed). The arrowed pyrite crystals clearly nucleate and grow in the solution and do not result from the transformation of iron–sulfide precursors. This is a rare view of the spatial relationships between synthetic iron-sulfide products. This SEM (scanning electron microscope) image shows the in situ arrangement of pyrite and mackinawite on a pyrite crystal surface: obviously the spatial relationships in filtered products are mixed and not in situ.

FeSm dissolves and the Fe and S components diffuse into the closed parenchyma cells through pores, and pyrite is precipitated on the interior cell walls. The result is unsurprising since the FeSm precipitate is nanoparticulate with a size as small as 2 nm. By contrast, the pyrite forms as crystals around 2 mm in size or 1000 times larger. Put in another way, it would take 109 FeSm nanoparticles to produce one pyrite crystal. The idea of some sort of solid-state transformation is physically improbable. FeSm, where it occurs, dissolves and pyrite forms from the reaction between dissolved iron and sulfur species to which the products of the FeSm dissolution reaction contribute. Thus FeSm is no different from any other Fe reactant. In the case of a closed system with no added dissolved Fe or S, the solubility product of FeSm is far greater than that of pyrite, and pyrite will ultimately precipitate from the products of the dissolution reaction. That is: FeSm þ Hþ /FeðIIÞ þ HS


(15)

FeðIIÞþS2 ð
IIÞ/FeS2p

(16)

Interestingly the sum of reactions 15 and 16 is FeSm þ Hþ þ S2 ð
IIÞ/FeS2p þ HS


(17)

where HS
reacts with S(0), for example, to produce more S2(
II) and the net effect is that the system becomes more alkaline. The distinction between an imagined solid-state process and the proven dissolution process is more than

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FIGURE 4 Pyrite formation with FeSm as a reactant involves initial dissolutiondrather than transformationdof FeSm. These SEM 4 quadrant backscatter images show the location of FeSm and pyrite in experimental pyritization of Apium sp. at low temperatures. The images demonstrate that the original FeSm in the xylem elements dissolves and pyrite is formed in the previously FeSmfree parenchyma cells (from Rickard et al., 2007). (a) Image of xylem elements from the experiments involving whole celery petioles. FeSm can be clearly seen infilling open xylem elements. The closed parenchyma cells in contrast do not contain FeSm. (b) Section of xylem elements showing the feather-like nature of the infilling FeSm aggregates. (c) Section of Apium petiole after reaction with H2S. The xylem elements are virtually free of FeSm, whereas the parenchyma cells are partially filled with pyrite.

mere sophistry. Sulfur-isotopic compositions of sedimentary sulfides are used as key proxies for investigations of the modern sulfur cycle and its geologic evolution. Since the sulfur cycle is intimately related to the global oxygen and carbon cycles, these measurements are central to understanding how the Earth system works. The conventional sedimentary cross-over plots (Fig. 5) showing the decrease of acid volatile sulfur with sediment depth and the increase of pyrite-S with depth do not balance, even if it is assumed that the acid volatile sulfur

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FIGURE 5 A cross-over plot compiled from data in Rickard et al. (1999) showing the mass balance problem in iron sulfides in sulfidic sediments from an Fe-rich estuarine environment. The increase in the concentration of FeS2p–S with depth is not balanced by the decrease in the concentration of acid volatile sulfide (AVS-S) in micromoles per gram dry weight of sediment. Note that the concentration of aqueous FeS clusters (FeSaq), shown as the current (in nanoamps) of the FeSaq voltammetric peak, increases in concert with pyrite consistent with FeSaq being a reaction intermediate in sedimentary-pyrite formation in systems containing FeSm. For a color version of this figure, the reader is referred to the online version of this book.

derives mostly from FeSm. This dissolution stage in the formation of pyrite in FeSm–containing sedimentary environments means that trace elements in FeSm are not incorporated directly into pyrite, as might be the case if a solid-state transformation occurred. As the FeSm dissolves, its trace element load is released back into solution. The importance of this is that trace elements removed by FeSm precipitation are not permanently fixed in pyrite. This, of course, has consequences for pollution studies and the use of pyrite trace elements as paleoenvironmental indicators. Since pyrite forms from dissolved species, which may include the products of FeSm dissolution, the specific location of pyrite formation is not necessarily the same as the site of any original FeSm precipitate. This means that vertical profiles of sedimentary sulfides are dynamic and cannot be treated in terms of static mass balances of sulfide phases, either chemically or isotopically. Mackinawite solubility in neutral to alkaline solutions is discussed above. Previously, it was supposed that the pH-dependent reaction Eqn (5.4) controlled mackinawite solubility in all pH values. Log Ksp;2 (FeSm) for this reaction is around
3.5, which means that {Fe2þ} in equilibrium with FeSm at pH ¼ 7 in the presence of millimolar S(
II) is around 10
14.5, which would imply that FeSm would not precipitate from normal seawater in the presence of millimolar sulfide. In fact, at the neutral to alkaline pH values characterizing many natural aqueous environments, including marine environments, FeSm is far more soluble and its stability is constrained by its intrinsic solubility Eqn (5.6) with log Kint (FeSm) ¼ 10
5.7. The total dissolved sulfide concentration, [S(
II)]T, and the total dissolved iron concentration, [Fe]T, both include the concentration of the aqueous iron–sulfide cluster, [FeSaq]. This means that, in the limiting case where most of the [S(
II)]T and [Fe]T are in the form of FeSaq, the concentrations of [S(
II)]T and [Fe]T must both equal at least 10
5.7 M for mackinawite to form. These concentrations must also be maintained for mackinawite to be preserved and not dissolve.

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The required minimum [S(
II)]T concentration of c. 2 mM occurs widely in sulfidic systems. However, the similar minimum [Fe]T concentration of c. 2 mM is three magnitudes greater that that of normal seawater (see Chapter 3). This means that the sulfidation of normal seawater, as described in Chapter 13, will not lead to mackinawite precipitation. As discussed in Chapter 14.1.1, the concentration of dissolved Fe in sediment pore waters is currently poorly constrained. The problem is analytical, particularly the distinction between dissolved Fe and nanoparticulate and colloidal Fe in pore waters. Thus although dissolved Fe concentrations of the order of 1–10 mM are often reported for sediment pore waters, these must be regarded as maximum values. By comparison, it took analysts many years to obtain a robust measurement of the concentration of dissolved Fe in the comparatively simple medium of openocean water and, before that time, dissolved Fe concentrations were similarly reported to be in the micromolar range (see Chapter 3). These considerations suggest that the [Fe]T concentration in marine sediment pore waters is normally insufficient to precipitate FeSm and that any FeSm that may form locally will become undersaturated and rapidly dissolve. This result is in accord with the data in Table 14.6 which shows that FeSm has only once been reported from apparently normal marine sediments. Many of the reports of FeSm as a precursor to pyrite actually stem from studies of inshore marine and fresh water environments where [Fe]T is high compared with normal marine systems. Such environments include the Mystic river site, originally studied by Berner, which had a scrap iron dump, the Black Sea, certain fjords and the Baltic Sea. Even the open-ocean site where mackinawite was reported (Lein et al., 1980) may have been affected by hydrothermal vent activity. The consequence of the contrast between the somewhat erratic and often short-lived occurrence of mackinawite in sediments and the widespread and persistent distribution of sedimentary pyrite is that mackinawite is not a necessary “precursor” to pyrite formation and pyrite forms in sediments in the absence of mackinawite. In systems, such as inshore and fresh waters where [Fe]T is high and FeSm does form, then it is in equilibrium with relatively high [Fe]T and [S]T. As reported above, thermodynamically such concentrations are adequate to precipitate pyrite. This is another way of saying that FeSm is unstable with respect to pyrite. As shown below, the actual mechanism of the reaction to form pyrite where FeSm is a reactant involves the dissolution of FeSm and the precipitation of pyrite. Note also that it is obvious that the processes describing pyrite formation from FeSm reactants can only occur in sediments where FeSm is present in the first place. In the majority of sediments, where FeSm does not occur, pyrite cannot of course be formed from this material. The apparent evidence for greigite, Fe3S4g, being involved in pyrite formation appears to have been (1) that FeSm transforms through a solid-state reaction to greigite, as discussed above, and (2) the observations of “magnetic pyrite”, from Doss’s (1912) original report, which appear to derive from intimate mixtures of

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greigite and pyrite. Bonev et al. (1989), for example, described pyrite rimming greigite concretions from Black Sea sediments. It has not been demonstrated either experimentally or mechanistically and it is structurally and chemically improbable. The origin of the “magnetic pyrite” and the intimate mixtures of greigite and pyrite in natural systems is simply that greigite dissolves and pyrite precipitates. In geology and mineralogy, this type of process is usually called replacement. A third reason may be that, at high temperatures, metastable greigite transforms to its stable counterparts, pyrrhotite and pyrite. However, the solid-state transformation of greigite to pyrite at low temperatures is kinetically hindered. The problem is that it requires reduction of the greigite-FeIII and oxidation of the greigite-S
II apart from a pervasive structural rearrangement. Structurally, the structures of mackinawite, greigite and pyrite are based on anion cubic close packing (ccp) but whereas mackinawite and greigite show arrays of simple atoms, the pyrite anion ccp is conceptual. It results from a rationalization of the arrangement of the average center points of the S2
2 dumbbells. Thus, whilst mackinawite can readily transform to greigite through a rearrangement of Fe atoms in a ccp sulfur lattice (Fig. 5.9), pyrite formation requires significant rearrangement of the sulfur lattice. Chemically, the formation of FeS2p from FeSm requires that the FeSm–S is oxidized but the FeS–Fe is not. In the FeSm / Fe3S4g transition, the FeSm–FeII is oxidized to FeIII whereas the FeSm–S
II remains unchanged. Pyrite is a low-spin semiconductor with a particularly low unit-cell volume (Persson et al., 2006) whereas the Fe in greigite and mackinawite is high-spin. These observations are consistent with the mackinawite / greigite transition being a solid-state transformation (highspin Fe / high-spin Fe), whereas solid-state transformations in the mackinawite / pyrite and greigite / pyrite transformations (high-spin Fe / low-spin Fe) are not possible at low temperatures under normal conditions. In fact, thermally-induced spin transitions of the type exhibited by iron sulfide mineral transformations at elevated temperatures are the most common types of spin-crossovers and help explain the temperature dependence of the processes. In low-temperature experimental syntheses of pyrite where the initial reactant is a dissolved Fe(II) salt, the solubility product for FeSm (and Fe3S4g) is almost always exceeded. The reasons for this are technical. In order to synthesize sufficient quantities of pyrite for analysis, the concentration of S(
II) and Fe(II) salts is normally >0.001 M. If the yield of the synthesis is 100% such concentrations would produce around 12 mg FeS2/100 mL solution. At these Fe(II) and S(
II) concentrations, the FeSm solubility product is exceeded at all pH values between ca 4 and 10 at 25  C. Since the rate of FeSm precipitation is rapid, FeSm cannot be avoided in such experiments. One way around this is to use a continuous flow system, which permits the gradual accumulation of sufficient quantities of FeS2p for analysis whilst avoiding FeSm precipitation. However, there are still some practical problems with this approach. For example, the S(
II) solutions at <0.001 M must be handled very carefully since they are prone to oxidationdand the lower the S(
II)

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concentration, the more the cryptic oxidation will affect the totals. One benefit of the system is, of course, that if FeSm precipitation can be avoided, Fe3S4g will not form since this is only produced from pre-existing FeSm. The use of acidic solutions, where the solubility product of FeSm is so high that precipitation might be avoided, but where FeS2p is still stable, appears to be precluded because of the formation of marcasite (and possible sulfur) under these conditions. The result of these technical problems has been that nearly all low-temperature pyrite syntheses have, effectively, FeSm as a reactant.

6.2. Sulfide-controlled Reaction Kinetics The reaction of FeS, either as pyrrhotite or mackinawite, with H2S to form pyrite stems from an original observation by Berzelius (1845), which was revisited by Wikjord et al. (1976), Huber and Wa¨chtershauser (1997), Rickard (1997) and Rickard and Luther (1997). The overall reaction is FeS þ H2 S/FeS2p þ H2

(18)

Reaction (18) is thermodynamically favorable at low temperatures (DGor ¼
29 kJ mol
1, at 25  C and 100 kPa total pressure, where FeS ¼ FeSm). Thus any discussions about its significance in the natural environment must solely be dependent on the kinetics. The reaction has become popular in experimental studies since it is a “clean” reaction, with simple, welldefined reactants compared with the practical complexities of handling systems with rhombic sulfur, for example. This H2S mechanism is important since S(
II) species are involved in all reported aqueous syntheses of pyrite. The key conceptual feature of this reaction is the counterintuitive idea that H2S can act as an oxidizing agent. An oxidizing agent is simply a chemical compound that can acquire electrons in a redox reaction. It is a relative term: a substance is always oxidizing or reducing with respect to something else. In this case, H2S is oxidizing relative to FeS. The reason that the idea that H2S can act as an oxidizing agent is counterintuitive in sedimentary geochemistry may be related to its commonly mistaken description of sulfidic sediments as “reducing”dwhich is using reducing as an absolute, rather than a relative, term. H2S is less nucleophilic than HS
since the highest molecular orbital (HOMO) for H2S is relatively stable (ca
10 eV) (Drzaic et al., 1984 Rickard and Luther, 1997). By contrast, the lowest unoccupied molecular orbital (LUMO) for H2S is around
1.1 eV and therefore H2S is a good electron acceptor (i.e. oxidizing agent). This compares with O2 itself where the LUMO is
0.47 eV. The key point here is that, although H2S can act as an oxidizing reagent with respect to [FeS], HS
cannot. By contrast, HS
is more nucleophilic than H2S and reacts faster with Fe(II) to form FeSm (Rickard et al., 1995). H2S is a key reactant for pyrite formation in aqueous solutions whereas HS
can be the

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dominant reactant for FeSm formation. Since pK1 for H2S is close to 7, the relative concentrations of these sulfide species decline logarithmically on either side of neutral pH (e.g. Fig. 2.10). Thus at pH 5, HS
makes up only c. 1% of the total dissolved S(
II) whereas, in normal seawater at pH 8.1, H2S makes up c. 9% of the total dissolved sulfide. The mechanism of the reaction between Fe(II) and H2S to form FeS2 thus again involves [FeS] as a reaction intermediate but, in this case, H2S does not substitute for [FeS] –S2
in a nucleophilic attack, but is involved in a redox reaction resulting in the oxidation of [FeS] –S2
to S2
2 . In both reactions, FeS is a nucleophile, but in the polysulfide pathway it is also an electrophile because higher order and more nucleophilic polysulfides can bind to the Fe(II) in FeS. The kinetics of this reaction involve a rate-controlling step involving H2S (Rickard, 1997; Rickard and Luther, 1997). The mechanism involves the formation of an inner–sphere complex between [FeS] and H2S, followed by electron transfer between S(
II) and H(I) to produce S2(
II): ½FeS þ H2 S/½Fe
S/SH2 /FeS2 þ H2

(19)

where [Fe–S / SH2] is a reaction intermediate. In this reaction, H2S acts as an oxidizing agent with respect to [FeS]. Butler et al. (2004) confirmed this mechanism isotopically. They demonstrated that the d34S of the product FeS2p from the reaction equaled a 1:1 mixture of the d34S of the reactant FeSm and H2S. The rate equation for the H2S mechanism is   (20) v FeS2p vt ¼ k½FeSm ½H2 S where [FeSm] is the molar concentration of FeSm (applicable because the reactant FeSm had a comparable initial surface area in all experiments), [H2S] is the molar concentration of H2S and k is the rate constant. Butler et al. (2004) noted that the so-called “iron loss pathway” reported by Wilkin and Barnes (1996) was actually a modification of the H2S reaction mechanism. The reaction stoichiometry proposed by Wilkin and Barnes (but not actually analytically demonstrated) was the sum of two well-established reactions: FeSm þ 2Hþ ¼ Fe2þ þ H2 S

(21)

FeSm þ H2 S ¼ FeS2p þ H2

(22)

2FeSm þ 2Hþ ¼ FeS2p þ Fe2þ þ H2

(23)

Butler et al. noted that the reaction of FeSm (d34S ¼ 2.8&) and H2S (d S ¼ 3.3&) by these processes would produce pyrite with an isotopic composition close to that of the reactant FeSm (and H2S), as observed by Wilkin and Barnes. The problem with the H2S reaction is the electron mass balance. In Eqn (22) the reduced product of the redox reaction is H2. In fact, this mass balance has 34

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257

never been demonstrated experimentally. Although H2 has been detected in the products with varying yields, dependent on temperature (Wikjord et al., 1976; Huber and Wa¨chtershauser, 1997; Rickard, 1997), the quantities suggested by Eqn (22) have not been recovered at low temperatures. Several different explanations have been given. H2 is notoriously difficult to trap experimentally and could escape from the reaction vessels, or H2 could be sequestered in the solid product. H2 could re-react rapidly with other components of the system or even be taken up by bacteria in the unsterilized apparatus. A more radical explanation for the absence of H2 at lower temperatures (i.e. < 100 C) is that the reaction is actually the polysulfide reaction and that the reaction with H2S does not become dominant until higher temperatures (e.g. > 100 C) have been reached. This idea stems from the above discussion that all H2S solutions have an intrinsic Sn(-II) content and thus increasing the H2S concentration increases the Sn(-II) concentration. The observation that the rate is a function of the H2S concentration is then indistinguishable from the rate being dependent on the Sn(-II) concentration as in the polysulfide pathway. The difference is that the polysulfide pathway does not result in H2 as a product. The lack of recovery of stoichiometric amounts of H2 in the experimental system is not a major problem in interpreting the kinetics and mechanism of the reaction. However, it would be nice to determine the fate of the reduced product of the reaction. More importantly, H2 is an important metabolite for a number of microorganisms, including sulfate-reducing bacteria. The production of H2 is important in investigations of the deep biosphere as well as in those concerning the early evolution of life. And the pyrite-forming reaction provides a ready and widespread source for this nutrient.

6.3. Polysulfide-controlled Pyrite Reaction Kinetics The reaction of polysulfide with iron salts to form pyrite was first described by Bunsen (1847). By 1886, Van Bemmelen could write that this reaction was well-known, as noted in the citation heading of Chapter 2. In experimental syntheses, pyrite is produced quite efficiently through a mixture of aqueous sulfide, sulfur and FeSm at low temperatures, and this was a favored route of the early experimentalists. Rickard (1975) showed that the rate of pyrite formation in these systems increases with increased polysulfide concentration. That is, the rate-controlling step in pyrite formation involves a reaction between a Fe species and polysulfide. This is unsurprising in view of the equilibrium thermodynamic background to pyrite formation discussed above. However, it is important in demonstrating that sedimentary pyrite forms through a solution reaction and not via a series of solid-state transformations. Rickard found that the rate of pyrite formation is first-order with respect to polysulfide and secondorder with respect to the FeSm surface area. The rate equation is    v FeS2p vt ¼ k A2FeSm ASð0Þ fSð
IIÞgT Hþ

(24)

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Sedimentary Pyrite

where AFeSm and AS(0) are the surface areas of FeSm and S(0), respectively, in cm2, {S(
II)}T is the total dissolved-sulfide activity and {Hþ} is the hydrogen ion activity. The kinetics show that pyrite formation from FeSm involves a solution reaction. The Arrhenius energy for the reaction is consistent with this: it is relatively high, showing that the rate-controlling step is a chemical reaction and not transport-controlled (e.g. diffusion) as would be the case in a solid-state transformation. Rickard concluded that pyrite formation occurred through the reaction between an aqueous Fe(II) species and polysulfide. The reaction is a complex heterogenous system. It includes several solid and aqueous reactants and produces a number of products. The rate law Eqn (23) therefore takes account of this complex mixture and necessarily includes the rates of a series of partially interdependent reactions. These include the formation of polysulfides through the reaction between rhombic sulfur and aqueous sulfide, the dissolution of FeSm and the formation of pyrite. Kamyshny et al. (2003) examined the kinetics of equilibration of aqueous polysulfides and found that this process was fast. The kinetics of the reaction between S8 and S(
II) were originally reported by Hartler et al. (1967), Boulegue and Michard (1978) and Kleinjan et al. (2004). Hartler et al. (1967) found that the rate was first-order with respect to total sulfide in the presence of polysulfide and possibly second-order with respect to S(
II) in the absence of polysulfides. They found that the Arrhenius energies for both reactions are low and that the reactions are probably mainly diffusion-controlled: that is, the chemical reactions are so fast that the rate is limited by the transport of reactants. The rate is not affected significantly by OH
, which suggests that the rate of dissolution is mainly independent of pH. Boulegue and Michard (1978) found that the rate of dissolution of S8 is first-order with respect to HS
. The problem with these transport-controlled reactions is that they are sensitive to the small variations in the nature of the reactant sulfur, through the surfacearea dependence, and are dependent on the relative velocity solution with respect to the particle surface. This means that even the reaction-vessel shape can affect kinetics of transport-controlled reactions (Sjo¨berg and Rickard, 1983). Rickard (1975) investigated the reaction under conditions where transport was not rate-controlling; he found that the rate of sulfur dissolution in the pyrite-forming system he investigated to be relatively fast and equilibration was approached. The second-order dependence of FeSm suggests that FeSm is involved in more than one rate-determining reaction, and the first-order dependence on {Hþ} at constant S(
II) suggests that the FeSm dissolution is involved in a rate-determining step. The rate of FeSm dissolution has been investigated by Pankow and Morgan (1979) as discussed in Chapter 5.3.5. Their data show that FeSm reaches equilibrium rapidly with dissolved Fe and S. The rate of pyrite formation is slow relative to the rate of dissolution of FeSm and thus FeSm acts as a continuous source of dissolved Fe and S components for pyrite formation.

6. Kinetics and Mechanisms of Sedimentary Pyrite Formation

259

Rickard (1975) suggested that the mechanism involved an aqueous FeS reaction intermediate which was attacked by nucleophilic polysulfides with formation of FeS2. The basic mechanism may be illustrated with respect to S
2 5 by the sequence Fe2þ þ Sð
IIÞ/½FeS ½FeS þ S2
5

/ FeS2 þ

S2
4

rapid rate
controlling

(25) (26)

where [FeS] is an aqueous iron(II)-sulfide reaction intermediate. Generally, for Sn(
II), where n 2, the most stable species are protonated and where n  3, the more stable species are unprotonated. Luther (1990) reviewed Sn(
II) nucleophilicity. The relative nucleophilicity of the species follow the energies of the HOMOs of the Sn(
II) species. Longer-chain polysulfides are more nucleophilic than shorter-chain polysulfides. The

2

2 nucleophilicity varies in the sequence S
2 5 > S4 > HS > HS2 > S3 > H2S.
2 Note that longer-chain Sn species are actually more nucleophilic than HS


and HS
2 is less nucleophilic than HS . This means that, although HS2 may be the most abundant polysulfide in many environmentally significant pH regimes, it is less nucleophilic than HS
, and likely to react more slowly. These frontier molecular orbital (FMO) considerations are interesting because they help explain why the addition of excess rhombic sulfur in a sulfide solution is a preferred route for the rapid synthesis of pyrite at low temperatures. As shown in Figure 2.14, the dominant Sn(
II) species in the S8 stability field are species with n  5. These longer-chain polysulfides are the most nucleophilic species and thus their relative abundance is expected to enhance the rate of the rate-controlling reaction in pyrite formation: the nucleophilic attack on the [FeS] reaction intermediate. Luther (1991) reacted pure S
2 n solutions with Fe(II) and independently reproduced the kinetics found by Rickard (1975). Fig. 6 shows the proposed mechanism, which indicates that the FeS2 formed should have all S atoms from S
2 n species. Butler et al. (2004) confirmed this

FIGURE 6 Pyrite formation via polysulfides. The reaction with FeS (a) is a special case of the generic reaction (b), where higher-chain nucleophilic polysulfides attack Fe centers of complexes with ligand L, to produce FeS2. The S from FeS (black) is exchanged and the resultant pyrite contains the S atoms from the polysulfide (gray).

260

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isotopically. They showed that the sulfur-isotopic composition of the pyrite product was similar to that of the polysulfide and was distinct from the sulfur in reactant FeS. The scission of the longer-chain polysulfides at the second bond is consistent with the relatively longer-bond length compared to the first S–S bond in the sulfur chain (Hordvik, 1966), which is reflected in the computed net charge on the individual sulfur atoms in polysulfides (Fig. 2.15). The net result is that nucleophilic attacks of longer-chain polysulfides on Fe(II) result in the preferential formation of the S2
2 ion. It can be seen that the process involving FeSm is a special case of a generic reaction where nucleophilic Sn(
II) attack Fe atoms of complexes (Fig. 6b). It has been suggested (Luther, 1991) that, on a molecular scale, the initial reaction is always the formation of an [FeS] transition complex even in systems where phases like FeSm or FeSaq are not formed. Luther pointed out that the key process is the change from relatively reactive, highspin Fe in Fe reactants to kinetically unreactive, low-spin Fe in pyrite. This could occur by clustering of Fe(II), S(
II) and Sn(
II) as occurs in the FeS centers of ferredoxins. This suggestion is consistent with the observation that pyrite does not form experimentally in aqueous solutions in the absence of S(
II) and provides an alternative explanation to the empirical interpretation based on the polysulfide equilibrium measurements described above. Furthermore, Guilbaud et al. (2011) reported that this spin-crossover was responsible for the Fe isotopic fractionation observed in this reaction (see Chapter 12.2.1). The H2S reaction is important in this discussion since it shows a mechanism whereby Fe(II) is initially associated with S(
II) before further reaction to form pyrite. In other words, both reactions involve the initial formation of an [FeS] intermediate. It appears that the molecular mechanisms of the polysulfide and sulfide pathways are essentially the same. The reaction between FeSm and H2S is clearly shown to actually be a reaction between a dissolved FeSaq species (in the form of a cluster) and H2S. In this system, FeSaq is reduced relative to H2S. It is possible that all dissolved or surface [FeS] species are reduced with respect to H2S, including the transition complexes formed during the polysulfide reaction.

6.4. Relative Rates of Pyrite Formation Although understanding the detailed molecular mechanism or pyrite formation is interesting theoretically, in practical terms the rate-controlling reactions of the two processes are competitive and complicated. The kinetics of pyrite formation in sediments are described by combining the rate Eqns (20) and (24) to produce a master equation Eqn (27) for the rate of pyrite formation which includes the polysulfide and H2S mechanisms (Rickard and Morse, 2005):     v FeS2p vt ¼ kH2 S ½FeSm ½H2 S þ Ksn ð
IIÞ ½FeSm 2 ½Sð0Þ½Sð
IIÞT Hþ (27)

6. Kinetics and Mechanisms of Sedimentary Pyrite Formation

261

In this formulation, KH2 S and Ksn ð
IIÞ are the experimentally measured rate constants for the H2S and polysulfide reaction, respectively, and the surface area of FeSm and S(0) in the polysulfide reaction has been assumed to be directly proportional to their concentrations since they can be approximated as standard materials. It is also assumed that [S(
II)]T z {S(
II)} and [Hþ] z {Hþ} compared with the original formulation as the errors propagated by these assumptions are well within the uncertainties of the calculations. Equation 27 is interesting because it appears to combine many of the most probable components involved in pyrite formation in sedimentary systems. As discussed above, the [S(0)][S(
II)]T product in Eqn (27) is mainly a measure of the kinetics of polysulfide formation. Eqn (27) reveals that, as this product tends to zero, the second term becomes very small and the H2S reaction dominates. The equation therefore takes into account sedimentary environments where significant concentrations of S(0) occur. Similarly, the [FeSm] term in Eqn (27) is an indirect measure of the formation of the FeSaq. However, the equation can also be applied to sedimentary systemsdmainly fresh waterdwhere mackinawite is formed. The one term that is missing is a term describing environments where greigite may be a reactant. Such systems are relatively rare and again mainly occur in fresh water sediments. The greigite term will basically involve the rate of dissolution of greigite since the process involves the dissolution of this phase. Of course, since greigite itself is formed by the solid-state transformation of mackinawite, the greigite term in the general equation will be a function of the FeSm concentration, and the rate of greigite formation will be competitive with the rate of FeSm dissolution. Overall, the missing greigite term will not significantly affect the general application of Eqn (27) to pyrite formation in sediments except in very specific environments. Although the experimental reaction between aqueous Fe(II) and S(
II) has not been demonstrated to produce pyrite directly at low temperatures, this may be because of the limitations of current experimentation in systems where the concentrations of Fe(II) and S(
II) are lower than the FeSm solubility product or FeSaq stability constant. In such systems, the quantity of pyrite formed is extremely low and difficult to detect, and the reaction is inhibited by the nucleation kinetics of pyrite, which require extreme limiting supersaturations (see below). The fact that pyrite crystals grow readily in sedimentary systems suggests that the process involves a simple reaction between dissolved Fe(II) and S2(
II) species. The growth of pyrite crystals in a sediment requires the transport of iron and sulfur to the site of crystal growth to maintain supersaturation. This would appear to be most easily achieved by the transport of relatively simple and widely available molecular species. Using Eqn (27), Rickard and Morse (2005) calculated the abiologic rates for a typical total dissolved sulfide concentration in anoxic aquatic and sedimentary systems of between 1 mM and 1 mM which, at pH ¼ 8, is equivalent to between 0.1 mM and 0.1 mM H2S. Typical sedimentary concentrations of acid volatile sulfide lie between 1 and 100 mmol gdw
1 sediment. Assuming all acid

262

Chapter | 6

Sedimentary Pyrite

volatile sulfide is in the form of FeSm (a limiting assumption for which there is little actual evidence) and with a porosity of >90%, this is equivalent to 1 and 100 mmol FeSm L
1 wet sediment. In this system, where [S(0)] ¼ 0 at 25  C, and with k ¼ 10
4 mol L
1 s
1, the rate of pyrite formation is then between 10
11 and 10
6 mmol L
1 sediment s
1 or 10
2 to 103 mmol FeS2p L
1 sediment a
1. Converting back to dry weight terms this is equivalent to 3  10
4 to 3  101 mmol FeS2p gdw
1 a
1. At pH ¼ 7, for the same conditions, [H2S] is between 0.5 mM and 0.5 mM, and the rate of pyrite formation 1.5  10
3 to 1.5  102 mmol FeS2p gdw
1 a
1. As [H2S] / 0 in alkaline sulfidic environments (e.g. pH > 9), the first term becomes very small and the rate is dependent on [S(0)] and [S(
II)] and pH. As discussed above, the [S(0)][S(
II)] term in Eqn (27) actually describes the rate of formation of polysulfides. The application of Eqn (27) to sedimentary conditions is not straightforward since it depends on knowledge of both the concentration of S(0) and FeSm in the sediment and does not take into account the potential catalytic effects of microorganisms on the system. To illustrate the relative abiologic rates, Rickard and Luther (2007) considered the case where the concentrations of S(0) and FeSm are of the same magnitude. The rate of pyrite formation for the polysulfide mechanism under the same environmental conditions as for the H2S mechanism, except that S(0) is present in equal molar quantities to FeS, then ranges between 10
2 and 10
11 mmol FeS2p gdw
1 a
1. This is between 10
4 and 10
8 times slower than the H2S mechanism in the absence of S(0). Even accepting considerable errors in the estimate of [FeSm] and [S(0)], the differences are so great that the polysulfide mechanism is relatively slow except under conditions where S(0) is present in significant concentrations. Thus, although the rate of pyrite formation may be described mainly by the kinetics of the polysulfide reaction initially near the redoxcline, pyrite is actually formed relatively slowly by this process, at least in abiologic systems. However, the microbial ecology of the anoxic/oxic boundary in sulfidic environments includes sulfur-disproportionating microorganisms which are closely involved in sedimentary-pyrite formation in this part of the system. Isotopic evidence (Canfield et al., 1998) shows that pyrite forms at similar rates through both the polysulfide and H2S processes during bacterial sulfur disproportionation, with total rates up to 105 times faster than the purely inorganic process would suggest. The nature of the catalytic processes enabled by these organisms is unknown. However, it is associated with extremely large sulfur isotope fractionations and the bacterial processes leave the pyrite with this signature. The effect of organisms, especially microorganisms, in the natural lowtemperature sulfide system is to bring the systems closer to equilibrium. In other words they increase the reaction rates. As discussed in Chapter 2, the inorganic sulfate–sulfide system is not reversible in low-temperature sulfidic environments because of the recalcitrant kinetics of inorganic sulfatereduction. Reversibility is achieved through the activities, for example, of

6. Kinetics and Mechanisms of Sedimentary Pyrite Formation

263

sulfate-reducing bacteria. The effect of microorganisms is therefore to increase the rates of reactions through catalytic effects of their biochemical pathways. This means that the experimentally measured rates of low-temperature inorganic systems may not describe the rates in any particular natural system. The inorganic studies provide a sort of base-line rate which may be locally enhanced depending, particularly, on the make-up the local microbial ecology. As noted above, some workers have limited the H2S reaction to highertemperature systems and have assumed that this reaction is far slower than the polysulfide reaction. In fact there is very little reason for this and those reports do not include calculations of the relative rates under the conditions addressed. In most experimental systems involving polysulfides, and in all normal natural environments, the pyrite-forming environment includes both polysulfides and sulfides, as shown thermodynamically above. So both reactions occur in natural sulfidic environments (Canfield et al., 1998). In the presence of FeSm, the rate is dependent specifically on the concentration of H2S in the one hand and polysulfides, especially the longer-chain species with n ¼ 4 or 5, on the other. Therefore, the concentration of polysulfides or H2S will be rate-limiting. In systems undersaturated with respect to FeSm, the concentration of Fe(II) may become rate-limiting depending on the sulfide or polysulfide concentrations. The concentration of both polysulfide and H2S are pH dependent at similar total sulfide loadings. Thus the H2S concentration increases as the pH decreases and the polysulfide concentration, as a fraction of the total sulfide, increases as the pH becomes more alkaline. Butler et al. (2004) considered the relative rates of these reactions under experimental conditions (FeSm present and P(H2S) ¼ 0.03 atm). They showed that, under these conditions, pyrite would form at a rate of w2  10
8 mol L
1 s
1 at pH ¼ 7 and 25  C through the H2S reaction. In the presence of excess S(0) under the same conditions, the polysulfide mechanism would produce pyrite at a rate of w8  10
11 mol L
1 s
1, or about two magnitudes slower than the H2S reaction. The polysulfide reaction also becomes more significant near the S(
II)/ SO2
4 redox boundary, where polysulfides may become important dissolved sulfur species. Thus in a natural marine sulfidic system with a pronounced redoxcline and an alkaline pH (w8.1), the polysulfide concentration may limit the rate of pyrite formation initially. As the system becomes more reduced and acidic (pH / 7), the H2S reaction may become rate-limiting.

6.5. Kinetics of Pyrite Nucleation and Crystal Growth Pyrite “formation” or “precipitation” consists of two processes: (1) nucleation, which requires relatively high supersaturations and/or catalytic effects of active surfaces or added trace components, and (2) crystal growth, which is relatively rapid and occurs through the reaction between Fe2þ and S2(
II). Kinetic studies mainly reveal the rate of the slowest reaction in a process. In the case of pyrite formation, supersaturations necessary to initiate pyrite

264

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Sedimentary Pyrite

nucleation are relatively high but, once achieved, nucleation is rapid. Harmandas et al. (1998) showed that, like many compounds with extremely low solubilities, pyrite requires extreme supersaturations, Up, in order to nucleate. They investigated the limiting supersaturation for pyrite nucleation, Up . This is the supersaturation limit up to which a phase cannot be expected to precipitate spontaneously. They found that for pyrite formation according to Eqn (14), Up needs to be in excess of 1014 at 25  C and pH ¼ 6.5 in the presence of pyrite seeds. In this sense, the observation of Schoonen and Barnes (1991b) is understandable: pyrite does not nucleate readily at Up < 1014. Grimes et al. (2001) synthesized pyrite from undersaturation with respect to FeSm or FeS0 within plant cells. In these experiments, initially precipitated FeSm in the open, outer xylem cells was shown to dissolve, and FeS2p formed in the closed, interior parenchyma cells. Calculations by Rickard et al. (2007) based on the parameters in the report by Grimes et al. (2001) showed that the biological surface provided a similar catalytic effect as a pyrite surface, with Up around 1011. It therefore appears that pyrite can nucleate from the reaction between hexaaqua Fe2þ and S2(
II) in the presence of an active surface. It is possible that other iron sulfides such as FeSm and Fe3S4g provide a similar role. Rickard and Luther (2007) calculated that Up in the presence of nanoparticulate FeSm is of the same order to that for pyrite formation in the presence of biological surfaces. As expected, nanoparticle surfaces show similar reactivities with respect to pyrite nucleation as biological surfaces. Harmandas et al. (1998) showed that the pyrite surface itself could act as a catalyst to allow rapid nucleation of FeS2p such that the rate-controlling reaction was crystal growth. This is in contrast to the suggestion by Schoonen and Barnes (1991b) that pyrite nucleation was the rate-limiting process. Rickard (1997) also showed that the formation of large numbers of equidimensional pyrite crystals during pyrite syntheses demonstrated that pyrite nucleation was not the rate-limiting process. This conclusion is important for understanding the formation of the framboidal texture, where up to one million pyrite crystals form at the same time within a limited volume (Butler and Rickard, 2000). The investigation of pyrite formation rates with polysulfide (Rickard, 1975) and with H2S (Rickard, 1997) were therefore mainly concerned with the rate and mechanism of pyrite crystal growth. The actual experimental method, in which the mass of pyrite formed with time is measured, demonstrates this. The mass of pyrite formed during nucleation is relatively small compared with the mass produced during crystal growth. In these investigations the concentrations of sulfur species present little or no thermodynamic problems. The equilibrium data discussed above show that the disulfide species HS
2 is present in sufficient concentrations in any experimental or natural S(
II) solution to account for pyrite formation. The experimental systems were saturated with respect to FeSm and the resultant FeS2p was produced by reaction between FeSaq and sulfur species. This is demonstrated by the observations that (1) the mechanism involves a solution

265

7. Sedimentary Pyrite Textures

reaction and (2) the suppression of FeSaq leads to the inhibition of pyrite formation. Homogeneous nucleation in solution is unexpected at the best of times and homogenous nucleation of relatively insoluble salts like FeS2p is unlikely. Harmandas et al. (1998) measured the rate of pyrite crystal growth, Rp, and showed that it fitted the semi-empirical equation Rp ¼ k s Unp

(28)

where n ¼ 3.5  0.5, k is the growth rate constant and s is a function of the active growth sites on the crystal seeds. The rate of pyrite crystal growth is given in units of mol m
2 s
1 which varied from 2.3  10
8 to 2.1  10
7 at Up between 1.1  107 and 2.4  107. The rate law and high value for n suggest a surface diffusion-controlled growth mechanism which is expected for sparingly soluble salts. Harmandas et al. reported that a similar order was found for the spontaneous precipitation of FeS2p at 80  C. It is therefore quite clear that pyrite crystals grow quite rapidly from a reaction involving Fe2þ and S2(
II) and that there is no requirement for any “precursor” mineral phases, such as FeSm, in the reaction.

7. SEDIMENTARY PYRITE TEXTURES Sedimentary pyrite occurs in two main textural forms: single crystals and framboids. The occurrence of a third formdpyrite spherulesdhas been widely documented in sedimentary rocks but, as is discussed below, this may not always be produced by normal marine or fresh water sedimentary processes. In each of these, there are many variants. Thus single crystals vary from nanoparticles to macrocrystals with various habits, while framboids include a variety of crystal habits and arrangements, and aggregates grade from simple groups of pyrite crystals to radiating, spherical forms. Because sedimentary pyrite is a key component in studies of modern and ancient Earth surface environments, it is important to understand what these various textures represent and how they are formed. This subject is in its infancy and much of the information is qualitative. There is a further problem of widespread misunderstandings about the origin of some textural forms. However, progress in understanding pyrite formation has enabled us to provide more robust interpretations of sedimentary-pyrite textures.

7.1. Single Crystals Single crystals mostly result from limited nucleation and continued supply of nutrients at low supersaturations to the site of crystal growth (Fig. 7). As discussed above, pyrite nucleation requires extremely high supersaturations in order to occur spontaneously. These supersaturations, at least

266

Chapter | 6

Sedimentary Pyrite

FIGURE 7 Elements of nucleation and crystal growth in terms of the relative saturation, U, and time. When the relative saturation exceeds the limiting supersaturation, U*, pyrite nucleates. In a normal, fine-grained sedimentary system, where the rate of supply of dissolved Fe(II) and S(
II) is lower than the rate of nucleation, the relative supersaturation decreases, nucleation is no longer possible spontaneously and the pyrite crystals grow.

11 magnitudes greater than the equilibrium values, mean that nucleation can be limiting to pyrite formation in sedimentary systems. One consequence of the extreme limiting supersaturations for pyrite in sedimentary environments is that the total concentrations of Fe(II) and S(
II) in solution become high before nucleation occurs. In other words, there is a relatively large amount of Fe(II) and S2(
II) available for nucleation with the result that there is a tendency for multiple pyrite nuclei to be formed. The consequence is the quite common formation of large numbers of nanoparticulate pyrite crystals in fine-grained sediments, which we colloquially term “pyrite dust” (Fig. 8). Nanoparticulate pyrite has also been found to constitute up to 10% of the black smoke of deep-ocean hydrothermal vents (Yucel et al., 2011). Nanoparticulate pyrite is not commonly found in sedimentary rocks and we must assume that it dissolves during diagenesis,

FIGURE 8 Nanoparticulate pyrite (A) synthesized at 25 C (B) from recent sediment. The pyrite dust particles are 100 nm in size and display some ordering developing in the natural sample.

7. Sedimentary Pyrite Textures

267

possibly with the formation of single large crystals by a process akin to Ostwald ripening. As shown in Fig. 8B, the pyrite dust is sometimes partially ordered and grades into the common framboidal texture, which is discussed further below. It is worth noting that biological surfaces also act to reduce the limiting supersaturation by a factor of at least 103 compared with, for example, pyrite seeds (Rickard et al., 2007). This helps explain the close association of pyrite with organic matter in sediments, including the widespread formation of pyritized fossils (see Fig.16.1). Of course, organic matter can also provide the metabolites required for sulfide production, particularly through the activities of sulfate-reducing bacteria. By contrast, pyrite crystal growth is achieved at low supersaturations and, since the solubility of pyrite is so low, these are readily achieved in sedimentary systems. Thus, as long as pyrite nucleates, the crystals will tend to grow as long as fresh nutrients are supplied to the site of nucleation. A key point here is that, in fine-grained sediments especially, the rate of supply of new nutrients, Fe(II) and S(
II), to the nucleation site is less than the nucleation rate. Thus nucleation itself depletes the nutrient concentration in the solution, reducing the relative supersaturation and preventing further nuclei to form. With pyrite, however, the solubility product is so low that the sedimentary system is usually supersaturated with respect to pyrite and the mineral continues to crystallize. For the pyrite crystal to continue growing, the rate of diffusion from the bulk solution to the mineral surface must not be less than the rate of crystal growth. When the rate of diffusion is less than the rate of crystal growth, the crystal will stop growing. The rate of crystal growth is then dependent on the rate of diffusion in the sediment. This effect is often tracked by the variation of S-isotopes across pyrite crystals. The initial nucleation and early growth stage are functions of the Sisotopic composition of the solution, but outer parts of the crystal may show isotopic compositions determined by Rayleigh distillation, partly closed system, effects (see Chapter 11.4). The rate of pyrite formation at 25 C has been reported to be w3  10
6 mol L
1 s
1 in two independent studies (Rickard and Luther, 1997; Guilbaud et al., 2011). The major problemdas originally addressed by Rickard (1973) and Raiswell (1982)dis the rate of diffusion of Fe and S to the site of pyrite crystal growth and the rate of diffusion of counterions from that site. A back-of-the-envelope calculation assuming steady state, a standard diffusion constant (Rickard, 1973) and a diffusion boundary layer (DBL) of 1–10 mm in thickness shows that the flux required could be attained by Fe and S concentrations in the bulk solution of 10
5–10
7 mol L
1. This is a significant result since this is less than the limiting concentration of aqueous Fe and S in equilibrium with FeSm (Rickard et al., 2006). Of course, a later pulse of nutrients, as is seen in the distribution of sulfatereduction in the deep biosphere (Chapter 10), will cause an increase in crystalgrowth rates. Usually, most of the new pyrite will nucleate and grow on earlier

268

Chapter | 6

Sedimentary Pyrite

pyrite surfaces, because of the effect of the pyrite surface on reducing the limiting supersaturation, as discussed above. The net result is the formation of pyrite overgrowths with differing, usually heavier, S-isotope compositions (e.g. Nishizawa et al., 2010). In the more extreme cases, of course, for example where there is a late hydrothermal pulse, the limiting supersaturation may be reached and new pyrite crystals nucleated. Some progress has been made in understanding the factors determining pyrite crystal habit. More than one hundred varieties of habit and more than four hundred sixty types of faces have been reported in euhedral pyrite crystals. However, above all, cubes a{100}, octahedra o{111}, pyritohedra e{210} and their intermediates, such as truncated octahedra and cubo-octahedra, are the most common habits observed in natural pyrite (Fig. 10). The variation of the three basic pyrite habits are frequently observed in sedimentary rocks and appear to be dependent on locality and coexisting minerals (e.g. Amstutz and Ligasacchi, 1958; Amstutz, 1963), sulfide concentration (Bush et al., 1960) and grain size, which is assumed to affect the appearance of e{210} as a{100} develops (Sunagawa, 1957). The surface energies of pyrite faces vary (see Chapter 7.3) and, near equilibrium, small pyrite particles will approach the shape of minimal surface energy. In the case of pyrite, the (100) face has the lowest surface energy and therefore low degrees of supersaturation produce cubic crystals. As the degree of saturation increases and the system becomes far from equilibrium, other faces may be exposed. So, for example, octahedral (111) planes may predominate. Looking at it the other way from a precipitation reaction with an initial high limiting supersaturation, we would expect the first pyrite crystals to be octahedra. As the concentration on nutrients in the solution decrease due to precipitation, cubic planes appear (producing cubo-ctahedra: Fig. 9c) and finally cubes.

7.2. Framboids Pyrite framboids are a common texture of sedimentary pyrite. Vallentyne (1963), for example, isolated 100,000 framboids per gram dry sediment from a North American lake. This is equivalent to a pyrite concentration of the order of 0.2 wt% dry sediment. The oldest framboids reported to date may be from the late Archean ( 2.9 Ga) sediments of the Witwatersrand Formation (Hallbauer, 1986; Guy et al., 2010). They are common in Phanerozoic black shales with modal sizes between 20 and 30 mm in diameter (Love and Amstutz, 1966; Rickard, 1970; Sassano and Schrijver, 1989; Wignall and Newton, 1998) and sizes ranging between 5 and around 100 mm. They occur commonly as infillings of fossils (e.g. Schallreuter, 1984; Roberts and Turner, 1993) including plant fossils (e.g. Grimes et al., 2002). They are found suspended in the water columns of euxinic systems (e.g. Ross and Degens, 1974; Skei, 1988) and meromictic lakes (e.g. Perry and Pedersen, 1993). These suspended framboids are generally 10 mm (Wilkin et al., 1996).

7. Sedimentary Pyrite Textures

269

FIGURE 9 Pyrite cubes and octahedra. (a) Pyrite cubes synthesized at 25  C (Photo H. Ohfuji). (b) Pyrite octahedron modified by cube synthesized at 25  C (Photo H. Ohfuji). (c) Layer of pyrite cubes from 2.32 Ga shales in the Pretoria Group, South Africa (Reflected light; photo E. Williams). (d) Large truncated octahedron with hopper faces from Early Jurassic shales, UK (Carstens, 1986).

It is important to note that framboidal pyrite is not restricted to sediments. It is a common texture in hydrothermal ore deposits (e.g. Kanehira and Bachinski, 1967), volcanogenic massive sulfide deposits (e.g. Rickard and Zweifel, 1975), sedimentary exhalative deposits (e.g. Love and Amstutz, 1966; Skei, 1988) and deep-ocean hydrothermal vents (e.g. Duckworth et al., 1994). Pyrite framboids have been observed in hydrothermal veinlets in volcanic rocks (Love and Amstutz, 1969) and meteorites (Jedwab, 1965). It is also noteworthy that the framboidal texture is not restricted to pyrite. Framboids of magnetite (Jedwab, 1971), hematite (Lougheed and Mancuso, 1973), and greigite (Nuhfer and Pavlovic, 1979) have also been reported. I have also observed framboidal marcasite and covellite. However, the most common mineral displaying the framboidal texture is pyrite and some of the reported occurrences of other minerals displaying this texture appear to have been derived from original pyrite textures. The history of framboidal pyrite is discussed in Chapter 1. Since the early suggestions by Schneiderho¨hn and Love that framboids were fossilized colonies of sulfur bacteria, the abiologic formation of framboids has been established. Pyrite framboids have been widely synthesized in inorganic chemical systems at low temperatures (Ohfuji and Rickard, 2005). There have been

270

Chapter | 6

Sedimentary Pyrite

FIGURE 10 Variations on a theme: the interrelationships between the various pyrite habits. The basic faces (a, cube; o, octahedron; e, pentagonal dodecahedron or pyritohedron) reflect various growth conditions. The change in growth conditions during pyrite crystal growth leads to the various combinations (After Sunagawa, 1957).

several problems in discussions of framboids and their origins, not least regarding a definition of what constitutes a framboid. There are an infinite variety of subspheroidal aggregates of pyrite microcrystals not only produced synthetically but also found in sediments and sedimentary rocks. Many of these are the products of simple intergrowths of pyrite microcrystals and are formed by similar processes to those discussed above. The word “framboid” derives from framboise, the French for raspberry, and was originally coined on account of their raspberry-like appearance (Rust, 1931). The key attributes of framboids are the subspheroidal outer form and the internal crystalline structure. A particular astonishing feature is the self-organization displayed by some pyrite framboids (Fig. 11b, d). However, the degree of self-organization is highly variable and most framboids display only limited or even no obvious internal organization (Fig. 11a, c). Ohfuji and Rickard (2005) defined framboids as: “Microscopic spheroidal to sub-spheroidal clusters of equant and equidimensional microcrysts.” This definition proscribes three fundamental characteristics of framboids: (1) They are spheroidal to subspheroidal in form (Fig. 11a–d). Framboids had classically been assumed to be spheroidal and the origin of the spheroidicity was a feature of earlier discussions regarding their origin. Morrissey (1972) originally described hexagonal sections of framboidal pyrite, anddwith

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271

FIGURE 11 SEM images of 4 types of sedimentary framboid (From Ohfuji and Rickard, 2005). (a) Disorganized (From the Uonuma Formation, Early Pleistocene, Japan). (b) Backscattered electron image of organized framboid (From the Chattanooga Shale, Devonian, USA). (c) Partially organized, indicated by arrows (From the Nishiyama Formation, Pleistocene, Japan). (d) Backscattered electron image of a sub-spheroidal (icosahedral) form with external faces. (From the Miocene Teradomari Formation, Japan).

increasing use of scanning electron microscope (SEM) techniquesdit has become apparent that the spheroidal form of many framboids derives from the development of outer, crystal-like faces (Fig. 11d). These forms may involve curved faces and polyhedron-like, icosahedral forms (Ohfuji and Akai, 2002), which are classically forbidden crystallographic symmetries. (2) They are composed of discrete microcrysts (Fig. 11a–d). A key feature of the framboidal texture is the microcrystals that constitute the material. That is, the interior of the framboid is not solid, homogeneous pyrite. The number of pyrite microcrystals in any single framboid is large, up to more than 106 (Rickard, 1970) but usually of the order of 104. Various attempts have been made to relate the mean crystal size, d, to the mean diameter of the framboids, D (Skipchenko and Berber’yan, 1976; Wilkin et al., 1996). The D/d ratio has been suggested as a paleoenvironmental indicator (see Chapter 13.7.6). In particular, the D/d ratio has been suggested to be related to the formation of framboids in a water column rather than within sediments (Wilkin et al., 1996). However, it is unclear what the controlling

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factor in the relation between crystal size and framboid diameter is. Rickard (1970) showed that the density of small framboids can be less than 1 g cm
3, and that they would float. However, the D/d ratio is an apparent correlation at present and, like all correlations, is intrinsically uncertain unless a robust mechanism is available to explain it. In general, the number of ordered framboids in a sediment or sedimentary rock appears to be a small fraction of the total framboid population. For example, Rickard (1970) estimated 10% ordered or partly-ordered framboids, and Love and Amstutz (1966) estimated that 28% of the framboids from the Chattanooga Shale are ordered or partly-ordered, and this is the classical locality for ordered framboids. However, some form of cryptic or localized, partial-order may occur in a large number of framboids. Partially-ordered microcrystalline or even nanoparticulate pyrite in irregular, nonspheroidal, aggregates occurs widely (Fig. 8b), especially in recent sediments. But these lack the obvious spheroidal characteristics of framboids. However, the occurrence of this type of material in nature demonstrates that microcryst ordering is unrelated to the outer form of the aggregate. (3) The microcrysts are equant. A key feature of the framboid texture is that the microcrysts are all approximately equant (Fig. 11a–d). A number of other spheroidal pyrite forms are found in nature and synthesized experimentally in which the material is radiating, acicular pyrite. Other spheroidal pyrite forms appear massive and result from interlocking anhedral pyrite crystals. Obviously, these forms have been produced by different processes to framboids. Framboidal microcrystal sizes vary between approximately 0.1 and 2 mm. The form of the microcrysts varies between simple cubes and octahedra to more complex cubo-octahedral forms. The occurrence of higher-order habits, such as pentagonal dodecahedra, is uncertain. There appears to be a relationship between the crystal habit and the crystal packing geometries. (4) The microcrysts are equidimensional. Within any single framboid the pyrite microcrystals all tend to have the same size (Fig. 11a–d). This is obviously a consequence of the pyrite-formation process. It also may be related indirectly to the overall spheroidicity of the framboids, in that dramatic variations in crystal size in any individual pyrite aggregate would not readily form an equidimensional spheroidal form. Pyrite framboids found in sediments range from <1 to 250 mm (normally, tens of micrometers) in diameter and consist of 102 w 107 microcrystals (<0.1–20 mm in diameter). Natural framboids are divided into two types by their internal structures; ordered and disordered types. In the ordered framboids, the microcrystals are arranged regularly, probably in a ccp packing, with a common crystal orientation (Love and Amstutz, 1966; Kalliokoski and Cathles, 1969) (Fig. 11b), whereas in the disordered framboids, the microcrystals are not packed in any

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273

regular arrays, but assemble in random orientations (Love and Amstutz, 1966; Rickard, 1970) (Fig. 11a). One key characteristic of framboidsdexactly what they aredwas not determined until the early years of the current millennium. Ordered framboids commonly show domain structures where the microcrystals are arranged in threefold and fivefold symmetries in sections (Morrissey, 1972; Skipchenko and Berber’yan, 1976; Large et al., 2001; Ohfuji and Akai, 2002). Such framboids are icosahedral and consist of twenty tetrahedral units of regularly arranged microcrystals (Ohfuji and Akai, 2002) (Figs. 11d and 12a, b). This is

FIGURE 12 False-color SEM images of ordered framboids. Both are sections through an icosahedral symmetry (From Ohfuji and Akai, 2002). (a) Trigonal domains. (b) Pentagonal domains. For a color version of this figure the reader is referred to the online version of this book.

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significant since icosahedral symmetry is a crystallographically forbidden symmetry and cannot arise through simple crystallization processes. Using single-crystal X-ray diffraction techniques, Ohfuji et al. (2006) demonstrated that even perfectly ordered framboids are not single pyrite crystals, although they appeared to show multiple domains where the microcrystals have similar orientations. That is, the individual microcrysts are not in simple crystallographic continuity. This was further refined by Ohfuji et al. (2005), who used electron backscatter diffraction techniques to show that the crystallographic orientation of individual microcrystals is random in disordered framboids but becomes more regular in ordered framboids. The crystallographic misorientation in ordered framboids is caused by slight physical misalignments of individual microcrystals and the gross symmetry characteristics of the pyrite structure which permits ordered packing to be maintained even after 90 relative microcrystal rotation. The results are important since they demonstrate that framboids are not produced by sequential crystallization of pyrite microcrysts but result from aggregation of existing microcrysts. The exceptional ordering is acquired by rotation of microcrysts during aggregation driven by a minimalization of surface free-energy between neighboring microcrystals. The aggregation process explains why there is a tendency for the ordering geometries to reflect the dominant microcrystal habit. Pyrite framboids have been synthesized by several of the recipes for pyrite synthesis listed in Table 1. However, the results do not appear to be reproducible (Ohfuji and Rickard, 2005). One of the simplest methodsdwith the cleanest productsdis through the reaction between H2S and FeSm (Butler and Rickard, 2000). We used this, together with a continuous flow system, to investigate the reason for the lack of reproducibility in experimental syntheses. We found that a keydand hitherto missingdcomponent of the syntheses was control of the hydrodynamics of the system. Thus framboids tend to form in zones with low advective flow. This is consistent with the abundance of framboids in mud and mudstones and within the tests of organisms. The process responsible for framboid formation is now clear. The evenness of pyrite microcrystal size within framboids is a key to aggregation and selforganization. This process is caused by the extremely high limiting supersaturations required for pyrite nucleation. When these values are exceeded, a nucleation burst occurs with large numbers of pyrite nuclei forming at the same time. Pyrite crystal growth is rapid but limited by the rapid reduction in concentration of dissolved Fe and S species in environments where the supply of further nutrients is restricted by poor advection, and is diffusion-limited. The result is the formation of large numbers of small equant pyrite crystals which aggregate under the force of a reduction in surface free-energy. The orderliness of the aggregation is a function of the homogeneity of the microcrystal sizes and shapes. This may be further perturbed by the inclusion of foreign particles in the system. The chances of perfectly ordered framboids are therefore relatively low, which is consistent with the situation observed in sediments.

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275

One further feature of natural framboids, which has not as yet been reproduced in the laboratory, is the inclusion of organic matter between the pyrite microcrysts. This has now been shown to be biofilm (Large et al., 2001; MaClean et al., 2008) produced mainly by extracellular polysaccharides by the assemblage of bacteria related to microbial sulfate-reduction. It is equivalent to the coacervates of the older workers (Papunen, 1966; Rickard, 1970). Again, formation of pyrite within biofilms is consistent with low-energy hydrodynamic regimes where the transport of Fe and S is mainly through diffusion. The organic geochemistry of ancient biofilms within pyrite framboids has not been investigated as yet but would make an interesting study.

7.3. Spheroidal and Nodular Pyrite Various spheroidal forms of pyrite have been reported in sedimentary rocks (e.g. Fig. 13). Forms such as these are far more abundant in hydrothermal systems and are important constituents of colloform textures. Many of these spheroidal and nodular forms appear macroscopically and in simple polished sections to be made of massive pyrite (e.g. Fig. 14). In fact, they are normally constituted by radiating, acicular crystals and these can be seen on etching (e.g. Fig. 13). Pyrite should always be etched, with nitric acid or ammonium ceric nitrate for example, while being examined with a reflected-light microscope. The identification of the source of the pyrite can be important. The spheroidal form in Fig. 13, for example, is a constituent of the pyrite analyzed by Bekker et al. (2004) and used to identify the onset of the oxygenation of the Earth’s atmosphere.

FIGURE 13 A 1 mm pyrite spherule, etched in ammonium ceric nitrate, from 2.3 Ga shales of the Pretoria Formation, South Africa. (Reflected light photo, Emily Williams). For a color version of this figure, the reader is referred to the online version of this book.

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FIGURE 14 Reflected-light micrograph of unetched colloform pyrite. (A) General view. (B) Detail of spheres of radiating pyrite showing organic-like forms which have been described as similar to budding or dividing organisms.

Some of these forms have previously been assumed to be inversions of the metastable marcasite, orthorhombic FeS2, because it is assumed that cubic pyrite does not form elongated crystals. However, this is not the case. Pyrite can develop extremely elongated habits producing acicular, wire-like or columnar crystals (e.g. Bonev et al., 1985). The radiating pyrite spheroids are a common feature of “colloform” pyrite aggregates. We have synthesized these radiating pyrite spherules in the Cardiff laboratory and have shown how the individual acicular crystals grow rapidly in conditions where the crystal protrudes into higher supersaturation, advective region beyond a thin diffusion boundary layer (Butler, 1994; Ohfuji, 2004). All stages in the process have been tracked, from individual acicular pyrite growing out from a heterogeneous nucleus to the completion and infilling of the spheroid. If the hydrodynamics of the environment changedso that the rate of pyrite crystal growth is not constantdthe result is the formation of pyrite bodies which are nonspheroidal and sometimes display eerily organic-like forms. These forms have been likened to budding organisms. The textures actually reflect growth processes in changing hydrodynamic conditions and the interference between coevally developing spheroids. The key aspect about the conditions giving rise to these spheroids is the maintenance of high supersaturation with respect to pyrite in the bulk solution. The supersaturation must be less than the limiting supersaturation, where spontaneous nucleation of pyrite may occur, which has been found to be as low as 1011 Ksp (pyrite). These textures develop in systems where new nutrients are rapidly supplied to the growth site over extended time periods. Large pyrite nodules and elongated forms (>1 cm in diameter) occur frequently in sediments and sedimentary rocks. There is no doubt that these may be formed diagenetically (e.g. Raiswell, 1982). I have studied pyrite nodules from modern Baltic Sea sediments which are up to 10 cm in diameter. These are hefty bodies with a volume of around 2500 cm3 and a weight of 12.5 kg. Each nodule contains w100 mol of FeS2. To form this, 100 mol of Fe and 200 mol of S must have reacted in a 2.5 L volume. The growth of the nodule must have made 2.5 L space in the sediment. All this is possible in the upper

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unconsolidated layers (with >90% water) of the organic-rich mud (or gyttja) in which the nodules were found. The rate of pyrite formation at 25  C has been reported to be w3  10
6 mol L
1 s
1 in two independent studies (Guilbaud et al., 2011; Rickard and Luther, 1997). This would suggest that the nodule could have taken just 10 years to grow. Obviously this is a minimum time period and the experimentally determined rate was achieved under ideal conditions which may not be characteristic of natural sedimentary systems. However, even if this is too fast by an order of magnitude, it demonstrates that the formation of large pyrite nodules is possible during sedimentary diagenesis. It also suggests that, as a first approximation, nodule formation can be treated as a steady-state system. The conditions responsible for the formation of sedimentary-pyrite textures are compared in Table 2 in terms of supersaturation with respect to pyrite and the dominant hydrodynamic regime. Euhedral crystals develop at low supersaturations. Extreme supersaturations, exceeding the limiting supersaturation, result in spontaneous nucleation. There is a trend in pyrite crystal growth from cubic at high supersaturations to octahedral as the supersaturation approaches the pyrite solubility product. Of course, when the saturation is less than the solubility product, pyrite dissolves. Examples of non-oxidative dissolution of pyrite are common in sediments and result in the redistribution of pyrite in the system, usually evidenced by overgrowths and late pyrite formationdoften with heavy sulfur isotopic signatures (see Chapter 11). In order to maintain crystal growth, new material must be supplied continuously to the site. The rate of diffusion must not be less than the rate of crystal growth. This can be achieved in a sedimentary system where diffusive conditions dominate but where,

TABLE 2 Summary Table of the Conditions Responsible for the Formation of Pyrite Textures in Sediments in Terms of Supersaturation, U, Solubility Product, Ksp (pyrite) and Hydrodynamic Regime Pyrite texture

Supersaturation / Low U / Ksp (pyrite)

Euhedral crystals

Cubic Y Octahedral

Spheroids

Acicular aggregates High Y Spheroids

Microcrystals

Nanocrystals Y Framboids

/ Ulimit

Dominant hydrodynamic regime Various

Stagnant (gently advective)

Diffusive

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in contrast to framboid formation, the Fe and S concentrations are maintained at levels above the solubility product for pyrite. It can also occur in rapidly advecting hydrothermal systems where the DBL is thin, and acicular crystals do not develop. In the original Berner steady-state model of diagenesis, a particle moves steadily through chemical gradients as the sediment accumulates. In the case of a euhedral pyrite crystal, this means that it will continue to grow whilst in a regime which allows Fe and S to diffuse to the site of crystal growth. As burial continues, the concentration of available Fe and S decreases and the resupply of dissolved Fe and S from producing levels via diffusion becomes slower than the rate of crystal growth, and the pyrite crystal stops growing. Again, simple estimations based on experimental measurements suggest that small euhedral pyrite crystals (<1 cm) develop in periods of a few years. Spheroids and nodules develop through the growth of multitudes of acicular crystals from the same point into the medium. In this case the DBL is relatively thick, and the crystals grow into the bulk solution. In order to maintain a spherical form, the supply of material and the advective regime must be fairly constant and the growth relatively rapid. This requires a gently advective or stagnant regime with a good supply of Fe and S. Spheroids and nodules are therefore usually found in systems where the supply of Fe and S is enhanced over normal marine conditions, by hydrothermal inputs or in fresh or brackish waters, for example. Microcrystals and nanocrystals of pyrite reflect conditions where a nucleation burst has occurred and where subsequent crystal growth has been limited by a restricted nutrient supply. This is achieved in diffusion-limited systems with little advection. The nucleation burst itself depletes the system of Fe and S and, in the absence of further nutrient supply, crystal growth is limited. Framboidal textures develop in those systems where aggregation of the microcrystals occurs as a result of minimization of surface free-energies. The contrast between the nucleation burst for framboids and the years it takes to grow a euhedral pyrite crystal is consistent with the widespread observation of diagenetically early framboids and later euhedra (e.g. Raiswell, 1982).

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