H.L. Barnes

The kinetics of silica-water reactions

Abstract A differential rate equation for silica-water reactions from 0–300°C has been derived based on stoichiometry and activities of the reactants in the reaction SiO2(s) + 2H2O(l) = H4SiO4(aq) ( ∂a H 4 SiO 4 ∂t ) P.T.M. = ( A M )(γ H 4 SiO 4 )(k+a SiO 2 a 2 H 2 O − k_a H 4 SiO 4 ) where ( A M ) = (the relative interfacial area between the solid and aqueous phases/the relative mass of water in the system), and k+ and k− are the rate constants for, respectively, dissolution and precipitation. The rate constant for precipitation of all silica phases is log k − = − 0.707 − 2598 T (T, K) and Eact for this reaction is 49.8 kJ mol−1. Corresponding equilibrium constants for this reaction with quartz, cristobalite, or amorphous silica were expressed as log K = a + bT + c T . Using K = k + k − , k was expressed as log k + = a + bT + c T and a corresponding activation energy calculated: a b c Eact(kJ mol -1) Quarts 1.174 -2.028 x 103 -4158 67.4–76.6 α-Cristobalite -0.739 0 -3586 68.7 β-Cristobalite -0.936 0 -3392 65.0 Amorphous silica -0.369 -7.890 x 10-4 3438 60.9–64.9 Upon cooling a silica-saturated solution below the equilibrium temperature, the decreasing solubility of silica causes increasing super saturation, which tends to raise the precipitation rate, but the rate constants rapidly decrease, which tends to lower the precipitation rate. These competing effects cause a maximum rate of precipitation 25–50°C below the saturation temperature. At temperatures below that of the maximum rate, silica is often quenched into solution by very slow reaction rates. Consequently, the quartz geothermometer will give the most accurate results if samples are taken from the hottest, highest flow rate, thermal springs which occur above highly fractured areas.

The size distribution of framboidal pyrite in modern sediments: An indicator of redox conditions

Abstract Pyrite framboids are densely packed, generally spherical aggregates of submicron-sized pyrite crystals. In this study, a survey was made of framboid size distributions in recently deposited sediments from euxinic (Black Sea; Framvaren Fjord, Norway; Pettaquamscutt River Estuary, Rhode Island, USA), dysoxic (Peru Margin), and oxic (Wallops Island, Virginia, USA; Great Salt Marsh, Delaware, USA) environments. Pyrite framboids in sediments of modern euxinic basins are on average smaller and less variable in size than those of sediments underlying dysoxic or oxic water columns. Down-core trends indicate framboid size distribution is a sediment property fixed very early during anoxic diagenesis, generally within the top few centimeters of burial. Size distributions in modern sediments are comparable with those in ancient sedimentary rocks, evidence that framboid size is preserved through advanced stages of diagenesis and lithification. It is proposed that where secondary pyrite growth is limited, as to preserve primary pyrite textures, framboid size distribution may be used to indicate whether fine-grained sedimentary rocks were deposited under oxic or anoxic conditions. The Crystal Size Distribution Theory relates framboid size to growth time and rate. On the basis of this theory, the characteristic smaller sizes of framboids in sediments of modern euxinic basins reflect shorter average growth times relative to oxic or dysoxic environments. In euxinic environments, framboid nucleation and growth occurs within anoxic water columns, and growth times are, on average, shorter because of hydrodynamic effects than when framboid nucleation and growth occurs within anoxic sediment porewaters underlying oxic water columns. A maximum framboid growth time of 0.4 years is indicated for framboids forming in the water columns of euxinic basins.

Formation processes of framboidal pyrite

Abstract Pyrite framboid formation may be the result of four consecutive processes: (1) nucleation and growth of initial iron monosulfide microcrystals; (2) reaction of the microcrystals to greigite (Fe 3 S 4 ; (3) aggregation of uniformly sized greigite microcrystals, i.e., framboid growth; and (4) replacement of greigite framboids by pyrite. The uniform morphology, uniform size range, and ordering of the microcrystals in individual framboids, as well as the range of observed framboid structures from irregular aggregates to densely packed spherical aggregates and polyframboids, are consequences of these processes. Using DLVO theory (Derjaguin, Landau, Verwey, and Overbeek), we have evaluated the stability of colloidal, iron monosulfide suspensions with ionic strengths typical of marine and lacustrine waters. In addition to van der Waals attractive and double-layer repulsive forces, a term is included to account for the ferrimagnetic properties of greigite. Numerical models predict that magnetically saturated greigite particles >0.1 μm in diameter will rapidly aggregate in either marine or fresh water. The aggregation model is in agreement with the sequence of greigite formation followed by pyrite framboid formation established in a previous experimental study (Sweeney and Kaplan, 1973) and is consistent with the occurrence of framboids composed of other magnetic minerals, e.g., greigite, magnetite, and magnesioferrite. Based on the temperature-dependent magnetic properties of greigite and aging experiments in hydrothermal solutions, this mechanism for framboid formation via precursor greigite could operate to temperatures of ∼200°C, consistent with the occasional occurrence of pyrite framboids in the paragenesis of metalliferous ore deposits.

Oxidation of pyrite in low temperature acidic solutions: Rate laws and surface textures

Rate laws have been determined for the aqueous oxidation of pyrite by ferric ion, dissolved oxygen and hydrogen peroxide at 30°C in dilute, acidic chloride solutions. Fresh, smooth pyrite grain surfaces were prepared by cleaning prior to experiments. Initial specific surface areas were measured by the multipoint BET technique. Surface textures before and after oxidation were examined by SEM. The initial rate method was used to derive rate laws. The specific initial rates of oxidation (moles pyrite cm−2 min−1) are given by the following rate laws (concentrations in molar units): rsp,Fe3+ = −10−9.74M0.5Fe3+M−0.5H+ (pH 1–2)rsp,o2 = −10−6.77M0.5O2 (pH 2–4)rsp,h2o2 = −10−1.43MH2O2 (pH 2−4) An activation energy of 56.9 ± 7.5 kJ mole−1 was determined for the oxidation of pyrite by dissolved oxygen from 20–40°C. HPLC analyses indicated that only minor amounts of polythionates are detectable as products of oxidation by oxygen below pH 4; the major sulfur product is sulfate. Ferric ion and sulfate are the only detectable products of pyrite oxidation by hydrogen peroxide. Hydrogen peroxide is consumed by catalytic decomposition nearly as fast as it is by pyrite oxidation. SEM photomicrographs of cleaned pyrite surfaces indicate that prior to oxidation, substantial intergranular variations in surface texture exist. Reactive surface area is substantially different than total surface area. Oxidation is centered on reactive sites of high excess surface energy such as grain edges and corners, defects, solid and fluid inclusion pits, cleavages and fractures. These reactive sites are both inherited from mineral growth history and applied by grain preparation techniques. The geometry and variation of reactive sites suggests that the common assumption of a first-order, reproducible dependence of oxidation rates on surface area needs to be tested.

PYRITE FORMATION BY REACTIONS OF IRON MONOSULFIDES WITH DISSOLVED INORGANIC AND ORGANIC SULFUR SPECIES

Abstract Pyrite formation has been investigated at 70°C and pH 6–8 by aging precipitated, disordered mackinawite, Fe9S8, and greigite, Fe3S4, in solutions containing aqueous H2S, HS−, Sx2−, S2O32−, SO32−, colloidal elemental sulfur, and the organic sulfur species thiol, disulfide, and sulfonate. Pyrite formed in all experiments where unoxidized iron monosulfides were aged with species containing zero-valent sulfur, i.e., polysulfides and colloidal elemental sulfur, but not with hydrogen sulfide (or bisulfide), the sulfoxy anions, or the organic sulfur species. Pyrite formation also occurred in experiments where the starting monosulfides were air-exposed prior to aging in sulfide solutions, or when air was bubbled through a reaction vessel containing iron monosulfides suspended in a sulfid sulfide solution. The experiments indicate the rate of conversion from iron monosulfides to pyrite is not only a function of solution chemistry (i.e., pH and aqueous speciation), but also depends on the surface oxidation state of the precursor iron monosulfides. Measurements of δ34S of reactants and products from pyrite-forming experiments suggest that the conversion from iron monosulfides to pyrite may proceed via loss of ferrous iron from, rather than via addition of zero-valent sulfur to, the precursor monosulfides. The sulfur isotopic composition of pyrite in sedimentary environments should reflect the sulfur isotopic composition of the precursor iron monosulfide plus sulfur sources incorporated during surface-controlled growth processes. Pyrite forms produced in this study ranged from poorly developed octahedral grains, in experiments where initial pyritization rates were the slowest, to framboidal aggregates in experiments where initial pyritization rates were the fastest. Although greigite formation occurred in experiments that produced framboids, not all experiments that produced greigite led to framboid formation. The formation of pyrite with framboidal texture is apparently favored when iron monosulfides rapidly convert to pyrite.

Reaction pathways in the Fe–S system below 100°C

Abstract The formation pathways of pyrite are controversial. Time resolved experiments show that in reduced sulphur solutions at low temperature, the iron monosulphide mackinawite is stable for up to 4 months. Below 100°C, the rate of pyrite formation from a precursor mackinawite is insignificant in solutions equilibrated solely with H2S(aq). Mackinawite serves as a precursor to pyrite formation only in more oxidised solutions. Controlled, intentional oxidation experiments below 100°C and over a wide range of pH (3.3–12) confirm that the mackinawite to pyrite transformation occurs in slightly oxidising environments. The conversion to pyrite is a multi-step reaction process involving changes in aqueous sulphur species causing solid state transformation of mackinawite to pyrite via the intermediate monosulphide greigite. Oxidised surfaces of precursors or of pyrite seeds speed up the transformation reaction. Solution compositions from the ageing experiments were used to derive stability constants for mackinawite from 25°C to 95°C for the reaction: FeS(s) +2 H + ⇔ Fe 2+ + H 2 S The values of the equilibrium constant, logKFeS, varied from 3.1 at 25°C to 1.2 at 95°C and fit a linear, temperature-dependent equation: logKFeS=2848.779/T−6.347, with T in Kelvin. From these constants, the thermodynamic functions were derived. These are the first high temperature data for the solubility of mackinawite, where Fe2+ is the dominant aqueous ferrous species in reduced, weakly acidic to acidic solutions.

Reactions forming pyrite and marcasite from solution: II. Via FeS precursors below 100°C

The formation of pyrite and marcasite from a FeS precursor has been examined experimentally. In aging experiments at 65°C, the conversion of precursor amorphous FeS depends on these geologically relevant variables: concentration of sulfur-contributing species, acidity, redox state, time, Fe(II)/S(−II) ratio in solution, and addition of an organic ligand (citrate). The results indicate that pyrite and marcasite formation proceed at a significant rate only if intermediate sulfur species (i.e., polysulfides, polythionates, or thiosulfate) are present in solution. In the absence of any sulfur contributor or with only hydrogen sulfide or bisulfide present, no FeS2 formed within 16 days. Sulfidation of the precursor proceeds through progressively more sulfur-rich, Fe-S phases: am FeS (Fe1.11S-Fe1.09S) → mackinawite (FeS0.93-FeS0.96) → greigite (Fe3S4) → pyrite/marcasite FeS2. Greigite is absent under very reduced environments. The conversion sequence found in this study is in good agreement with iron-sulfide distribution patterns found in modern marine sediments.

Solubility of gold in aqueous sulfide solutions from 150 to 350°C☆

Abstract The solubility of gold was measured in aqueous sulfide solutions at pH from 3 to 8, 150° to 350°C, and at pressures determined by the liquid-vapor pressure of the solution, with oxidation state fixed or buffered by either sulfate-sulfide equilbria or H 2 (g). High solubilities were measured in solutions with near neutral pH with a maximum measured gold concentration of 0.036 m (7224 mg/kg) at 350°C in a solution containing 0.66 m H 2 S and 0.28 m NaHS. The results are consistent with the aqueous complex Au(HS) 2 − . Log equilibrium constants for the reaction Au + H 2 S ( aq ) + HS − = Au ( HS ) 2 − + 1 2 H 2 ( g ) at 150°, 200°, 250°, 300°, and 350°C were determined as −2.39 ± 0.2, −1.89 ± 0.2, −1.56 ± 0.3, −1.35 ± 0.3, and −1.22 ± 0.2, respectively. These values are in reasonable agreement with published data at both 25°C and elevated temperatures. The high stability of Au(HS) 2 − indicates that geologically significant quantities of gold can be transported in typical hydrothermal solutions. Calculated gold solubility for the Ohaaki geothermal system in New Zealand shows that Au(HS) 2 − can easily account for the measured hydrothermal gold concentration. Gold may be precipitated from solution by both pH and redox changes. In addition, decreasing the activity of sulfide in solution is an effective mechanism for gold deposition. Analysis of the effect of temperature on the solubility of gold shows that a decrease in temperature may increase or decrease solubility. Deposition by cooling depends upon the pH-oxidation state path of the solutions.

Marcasite precipitation from hydrothermal solutions

Pyrite and marcasite were precipitated by both slow addition of aqueous Fe2+ and SiO32− to an H2S solution and by mixing aqueous Fe2+ and Na2S4 solutions at 75°C. H2S2 or HS2− and H2S4 or HS4− were formed in the S2O32− and Na2S4 experiments, respectively. Marcasite formed at pH < pK1 of the polysulfide species present (for H2S2, pK1 = 5.0; for H2S4, pK1 = 3.8 at 25°C). Marcasite forms when the neutral sulfane is the dominant polysulfide, whereas pyrite forms when mono-or divalent polysulfides are dominant. In natural solutions where H2S2 and HS2 are likely to be the dominant polysulfides, marcasite will form only below pH 5 at all temperatures. The pH-dependent precipitation of pyrite and marcasite may be caused by electrostatic interactions between polysulfide species and pyrite or marcasite growth surfaces: the protonated ends of H2S2 and HS2 are repelled from pyrite growth sites but not from marcasite growth sites. The negative ions HS2 and S22− are strongly attracted to the positive pyrite growth sites. Masking of 1πg* electrons in the S2 group by the protons makes HS2 and H2S2 isoelectronic with AsS2− and As22−, respectively (Tossellet al., 1981). Thus, the loellingitederivative structure (marcasite) results when both ends of the polysulfide are protonated. Marcasite occurs abundantly only for conditions below pH 5 and where H2S2 was formed near the site of deposition by either partial oxidation of aqueous H2S by O2 or by the reaction of higher oxidation state sulfur species that are reactive with H2S at the conditions of formation e.g., S2O32− but not SO42−. The temperature of formation of natural marcasite may be as high as 240°C (Hannington and Scott, 1985), but preservation on a multimillion-year scale seems to require post-depositional temperatures of below about 160°C (Rising, 1973; McKibben and Elders, 1985).

Reactions forming pyrite and marcasite from solution: I. Nucleation of FeS2 below 100°C

Reaction paths for nucleation and growth of pyrite and marcasite from solution have been investigated experimentally. Conditions were chosen to avoid the precipitation of metastable Fe-S phases which can act as precursors for FeS2 formation. The experiments indicate that FeS2 nucleation is extremely slow below 100°C. Instead of FeS2 nuclei, the reaction of ferrous ions and polysulfide ions produces initially amorphous FeS. Although the nucleation of FeS2 is inhibited below 100°C, pyrite and marcasite can grow from solutions devoid of polysulfides and undersaturated with respect to possible Fe-S precursor phases. The inability of pyrite to rapidly nucleate explains high supersaturation with respect to pyrite and marcasite in anoxic environments. Although pyrite is the stable Fe-S phase in these environments, it will not control the Fe2+ and H2S (or HS−) concentrations until its growth rate exceeds the dissolution rate of far more soluble, metastable FeS precursor phases.