Experimental investigation on thermochemical sulfate reduction by H2S initiation

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Geochimica et Cosmochimica Acta 72 (2008) 3518–3530 www.elsevier.com/locate/gca

Experimental investigation on thermochemical sulfate reduction by H2S initiation Tongwei Zhang a,1, Alon Amrani a, Geoffrey S. Ellis b, Qisheng Ma a, Yongchun Tang a,* a

Power, Environmental, and Energy Research Center, Division of Chemistry and Chemical Engineering, California Institute of Technology, Pasadena, CA 91125, USA b U.S. Geological Survey, Box 25046, MS 939, Denver Federal Center Denver, CO 80225, USA Received 11 September 2007; accepted in revised form 30 April 2008; available online 15 May 2008

Abstract Hydrogen sulfide (H2S) is known to catalyze thermochemical sulfate reduction (TSR) by hydrocarbons (HC), but the reaction mechanism remains unclear. To understand the mechanism of this catalytic reaction, a series of isothermal gold-tube hydrous pyrolysis experiments were conducted at 330 °C for 24 h under a constant confining pressure of 24.1 MPa. The reactants used were saturated HC (sulfur-free) and CaSO4 in the presence of variable H2S partial pressures at three different pH conditions. The experimental results showed that the in-situ pH of the aqueous solution (herein, in-situ pH refers to the calculated pH of aqueous solution under the experimental conditions) can significantly affect the rate of the TSR reaction. A substantial increase in the TSR reaction rate was recorded with a decrease in the in-situ pH value of the aqueous solution involved. A positive correlation between the rate of TSR and the initial partial pressure of H2S occurred under acidic conditions (at pH 3–3.5). However, sulfate reduction at pH 5.0 was undetectable even at high initial H2S concentrations. To investigate whether the reaction of H2S(aq) and HSO4  occurs at pH 3, an additional series of isothermal hydrous pyrolysis experiments was conducted with CaSO4 and variable H2S partial pressures in the absence of HC at the same experimental temperature and pressure conditions. CaSO4 reduction was not measurable in the absence of paraffin even with high H2S pressure and acidic conditions. These experimental observations indicate that the formation of organosulfur intermediates from H2S reacting with hydrocarbons may play a significant role in sulfate reduction under our experimental conditions rather than the formation of elemental sulfur from H2S reacting with sulfate as has been suggested previously (Toland W. G. (1960) Oxidation of organic compounds with aqueous sulphate. J. Am. Chem. Soc. 82, 1911–1916). Quantification of labile organosulfur compounds (LSC), such as thiols and sulfides, was performed on the products of the reaction of H2S and HC from a series of gold-tube non-isothermal hydrous pyrolysis experiments conducted at about pH 3 from 300 to 370 °C and a 0.1-°C/h heating rate. Incorporation of sulfur into HC resulted in an appreciable amount of thiol and sulfide formation. The rate of LSC formation positively correlated with the initial H2S pressure. Thus, we propose that the LSC produced from H2S reaction with HC are most likely the reactive intermediates for H2S initiation of sulfate reduction. We further propose a three-step reaction scheme of sulfate reduction by HC under reservoir conditions, and discuss the geological implications of our experimental findings with regard to the effect of formation water and oil chemistry, in particular LSC content. Ó 2008 Elsevier Ltd. All rights reserved.

*

Corresponding author. Fax: +1 626 858 9250. E-mail address: [email protected] (Y. Tang). 1 Present address: Bureau of Economic Geology, The University of Texas at Austin, University Station, Box X, Austin, TX 787138924, USA. 0016-7037/$ - see front matter Ó 2008 Elsevier Ltd. All rights reserved. doi:10.1016/j.gca.2008.04.036

1. INTRODUCTION The most prominent abiotic alteration process in hot carbonate reservoir rocks is the reduction of sulfate to sulfide coupled with the oxidation of hydrocarbons (HC) to

Thermochemical sulfate reduction by H2S initiation

carbon dioxide, which collectively is termed thermochemical sulfate reduction (TSR). TSR is well documented in numerous geological observations from around the world (Le Tran et al., 1974; Orr, 1974; Krouse et al., 1988; Sassen, 1988; Claypool and Mancini, 1989; Heydari and Moore, 1989; Machel et al., 1995; Rooney, 1995; Connan et al., 1996; Worden and Smalley, 1996; Carrigan et al., 1998; Worden et al., 2000; Cai et al., 2003; Li et al., 2005; Zhang et al., 2005). There are also many examples where TSR reactions are implicated in the deposition of metal sulfide ore bodies in both high temperature and sedimentary settings (Powell and Macqueen, 1984; Leventhal, 1990; Sun and Puttmann, 2000; Alonso-Azcarate´ et al., 2001; Bechtel et al., 2001; Garven et al., 2003). Efforts to simulate TSR experimentally have generally focused on a determination of the reaction kinetics and on attempts to understand the details of the reaction mechanisms involved (Toland, 1960; Dhannoun and Fyfe, 1972; Kiyosu, 1980; Trudinger et al., 1985; Kiyosu and Krouse, 1990; Hoffmann and Steinfatt, 1993; Kiyosu and Krouse, 1993; Goldhaber and Orr, 1995; Cross et al., 2004). More recently, the effect of TSR on the alteration of the gas chemical compositions and on carbon isotopic composition was investigated (Manzano et al., 1997; Pan et al., 2006; Zhang et al., 2008). The effect of oil composition on TSR reactions, in particular the lowering of the onset temperature of TSR by the presence of labile organosulfur compounds (LSC), was examined (Zhang et al., 2007; Amrani et al., 2008). Sulfate reduction rates were reported to increase with the addition of hydrogen sulfide (H2S) as an initiator in the oxidation of HC, and also to increase with increases in both temperature and H2S pressure (Goldstein and Aizenshtat, 1994; Goldhaber and Orr, 1995). However, reduction was not measurable, at the experimental time scale used, without having H2S initially present. Several authors have suggested that the catalytic effect of H2S on TSR is due to its reaction with sulfate to produce elemental sulfur, which in turn reacts with organic matter to produce H2S (Feely and Kulp, 1957; Orr, 1974; Powell and Macqueen, 1984; Goldhaber and Orr, 1995). The catalysis of TSR by lower valence sulfur species has been shown to be more effective at lower pH conditions (Toland, 1960), which suggests that this mechanism may not be important for natural TSR occurrences where pH values are likely buffered to higher levels by carbonate rocks. Further study is needed to determine the relative reactivity of sulfur with different valence states under geologic conditions, and to elucidate the role that these intermediaries play in TSR. A detailed understanding of the mechanism of H2S catalysis of sulfate reduction is necessary to constrain kinetic models derived from laboratory experiments, which can then be extrapolated to geologic conditions. The objectives of this study are to (1) identify the main controls on sulfate reduction by H2S initiation in order to constrain kinetic models, (2) investigate LSC formation from the reaction of H2S and HC, and (3) propose a general reaction scheme for H2S initiation of sulfate reduction. To address these objectives, we developed a new method to quantitatively load H2S gas into small gold-tube reactors.

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We conducted a series of isothermal gold-tube hydrous pyrolysis experiments in the presence of H2S and CaSO4 with and without HC under constant confining pressure. Our experimental results provide important insights into the reaction mechanisms of TSR and the effect of pH, temperature, and H2S partial pressure on the rate of sulfate reduction. 2. SAMPLES AND EXPERIMENTAL METHODS 2.1. Sample description The model HC used in this study included pure n-hexadecane (n-C16) and a paraffin mixture (Sigma–Aldrich, St. Louis, MO) that is a research grade paraffin wax composed of C21 to C35 normal alkanes, with a melting point of 52– 58 °C. A gas chromatogram for the HC composition of the mixture is shown in Zhang et al. (2007). The model compounds used contain no sulfur. Inorganic reagents used in our experiments included calcium sulfate (CaSO41/ 2H2O), magnesium chloride (MgCl26H2O), sodium chloride (NaCl). Because all three inorganic reagents were prepared as a mixed aqueous solution prior to loading to gold-tube reactors, there is no difference by using bassanite, anhydrite or gypsum as the source of sulfate under our experimental condition. Magnesium silicate (talc: 3MgO4SiO2H2O) and silica gel (SiO2) were all purchased from Alfa Aesar and Aldrich. Dolomite (CaMg(CO3)2), which was donated by the British Museum, is from Silurian Interlake Group, Manitoba, Canada. 2.2. Experimental methods 2.2.1. Quantitatively loading H2S gas into gold-tube reactor 2.2.1.1. Loading H2S gas into gold-tube reactor. We developed a new method of quantitatively loading H2S gas into gold-tube reactors. Gold was the material of choice because of its chemical inertness and flexibility that allows volume expansion and contraction for external control of the confining pressure. This method makes use of our existing vacuum line system that is used for gas transfer and total gas volume quantification, which was readily adaptable for the quantification of the gas yields from our TSR experiments. Each gold tube was between 110- and 120-mm-long, with an internal diameter of 4.3 mm and a wall thickness of 0.45 mm, giving a total reactor volume of approximately 1.4 mL. Prior to loading the samples, the open-ended tubes were heated to 600 °C to remove any residual organic material. One end of each tube was then crimped and sealed using an argon arc-welder. The solid reactants were accurately weighed and transferred to the tubes by a small funnel with an outside diameter slightly smaller than the inner diameter of the gold tubes. Liquid organic reactants and aqueous solutions, with an ionic strength of 3.91 M, were loaded into the tubes by means of a 100-lL pipette. A schematic diagram of the apparatus for loading H2S gas into gold-tube reactors is shown in Fig. 1. The tubes, filled with the solid and liquid reactants, were then connected to a stainless steel manifold with 3/16-in Swagelok nuts, ferrules, and a custom-made titanium insert.

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T. Zhang et al. / Geochimica et Cosmochimica Acta 72 (2008) 3518–3530 Digital pressure Helium for flushing air from gold-tube reactor

2

1

Vent connecting to vacuum line

3

H2S sampling cell with sulfur-inert coated tube and valve

Pressure device

Gold-tube reactor

Air as driving force of pressure device

Fig. 1. Schematic diagram showing the loading of H2S gas into gold-tube reactors. Numbered valves correspond to descriptions in text of Section 2.2.1.

ing, valve 4 was closed and the gold tube was cut just beneath the 3/16-in. nut of the manifold using a knife. The gas-filled gold tube was isolated from the manifold by the pressure device, and the open end of the gold tube was crimped. The crimped end was welded using an argon arc-welder while the other end was submerged in a liquid nitrogen trap (196 °C) to trap any volatiles created during the welding process. Verification of the amount of H2S loaded into the goldtube reactors was made by gas chromatography thermal conductivity detection (GC-TCD) and compared with our calculated values based on the initial gas pressure in the gold tube. These results are listed in Table 1. Following the above-described H2S loading procedure, we prepared four gold tubes with equivalent amounts of solids and liquids and variable amounts of H2S. Without heating the gold tubes, the gas was then recovered and analyzed following the procedure described in Section 2.2.3; the amount of H2S measured is listed in Table 1. The calculated amount of H2S is based on the gas pressure during loading and the effective volume of the gold tube, assuming ideal gas behavior (PV = nRT). The comparative results show that the recovered amount of H2S is within 3% of the amount loaded, over amounts ranging from 15 to 200 lmol (Table 1), confirming that the method of loading H2S gas into gold-tube reactors is accurate and workable.

The stainless steel manifold, with a total volume of about 6.1 mL, consisted of 1/8 in. OD tubing, two-way Swagelok valves with 1/8-in. fittings and connections, a 4.3-mL sampling cell with a two-way Swagelok valve on each end, a digital pressure gauge, and a vent connected to a vacuum system (Fig. 1). All the valves, tubing, fittings, and connections were made sulfur-inert by a special treatment by Restek. Prior to loading H2S gas into the gold tube, a leak check was performed by introducing 120 psi of He gas into the gold tube and monitoring the pressure. After ensuring that there was no gas leakage (leak testing is essential because H2S gas is highly toxic), the gold tube was flushed with He for 5 min to remove air. Valves 1 and 2 were closed, then the background pressure was recorded with a digital pressure gauge. Opening valve 3 allowed H2S gas collected in the sampling cell to be introduced into the gold-tube reactor due to the pressure difference. After about 5 min for gas equilibration, the gold tube was crimped at a point approximately 70 mm above the bottom and held by a compression device, which supplied about 6000 psi of hydraulic pressure driven by nitrogen gas. The pressure of the mixed gas (He + H2S) in the gold tube was recorded by means of the digital pressure gauge. Prior to isolating the gold-tube reactor from the manifold, the remaining gas in the manifold was vented by opening valve 4, which was connected to a vacuum line. After about 3 min of vent-

Table 1 The amount of H2S in gold tube reactors (before heating) as measured by gas chromatography compared with the calculated amounts based on the gas pressure and the volume of the gold-tube reactors (assuming ideal gas behavior) Gold-tube length (cm)

9.4 9.4 9.4 9.4

Filled gas pressure (psi)

Filled H2S amount (lmol)

Helium

H2S mixture

Calculated

Measured

18.7 17.7 18.7 19.7

144.7 75.7 43.7 30.7

200.4 93.6 40.5 13.9

205.2 nm 40.5 15.0

Note: H2S mixture gas is 60% H2S:40% Ar. nm, not measured.

Thermochemical sulfate reduction by H2S initiation

2.2.1.2. Examination of potential impact of the dissolution of gold in aqueous H2S to our experimental results. It has been shown that aqueous sulfide solutions at elevated temperatures (>150 °C) are capable of dissolving gold through the formation of hydrosulfide complexes that also result in the formation of hydrogen gas (Benning and Seward, 1996; Tagirov et al., 2005 and references therein). In order to examine the potential impact of the dissolution of gold in the aqueous hydrogen sulfide solutions in our experiments, we conducted two blank experiments at pH about 3.5 and 5 in the presence of 190 lmol H2S at 330 °C for 24 h. The results are listed in Table 2. No detectable hydrogen from reaction of the gold tube with hydrogen sulfide was observed at pH 3.5, whereas at pH 5 the amount of hydrogen produced (1.2 lmol) was too small to have any affect on the TSR reaction. These observations are in agreement with a reported maximum of AuHS2  complex formation at near neutral pH (Benning and Seward, 1996; Tagirov et al., 2005). Under acidic conditions the solubility of gold is 1–2 orders of magnitude lower than at near neutral pH, and in these environments aqueous gold predominately forms the neutral complex Au(HS)° (Benning and Seward, 1996; Tagirov et al., 2005). As shown in Section 3.1.2, the rate of TSR observed in our experiments is significantly faster at pH 3.5 than at pH 5, which is in opposition to the expected effect of the increased concentration of hydrogen gas. Therefore, the fact that little to no hydrogen gas is generated in our blank experiments and that the concentration of hydrogen gas increases as the rate of TSR decreases, strongly indicates that the reaction of the gold tube with aqueous sulfide has a negligible effect on the observed rate of TSR under our experimental conditions. 2.2.2. Hydrous pyrolysis After individually sealed gold tubes were loaded with reactants, they were placed in separate stainless steel autoclaves and inserted into a pyrolysis oven. Pyrolysis experiments were conducted under isothermal conditions at temperatures of 300 and 330 °C for different reaction times, or under non-isothermal conditions from 300 to 370 °C at 0.1 °C/h heating rate. Temperature was controlled to within 1 °C of the set value, and was monitored using a thermocouple secured to the outer wall of each autoclave. Our experiments were intentionally conducted below the critical temperature of water (374 °C) to keep the reaction in the liquid state. A water pump was used to maintain a constant confining pressure at approximately 24.1 MPa (3500 psi), which is about 2 MPa greater than the saturated vapor pressure at the critical temperature of water, to prevent rupturing of the gold tubes and keep H2O in the liquid phase at these elevated temperatures. When the desired reaction tempera-

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ture or time was reached, the stainless steel autoclaves were withdrawn from the oven, and rapidly cooled to room temperature by quenching in water. Once the autoclaves were depressurized, the gold tubes were recovered for detailed analysis of their contents. 2.2.3. Gas analysis Each gold tube was loaded separately into a piercing unit and connected to a custom-made glass vacuum line with a residual pressure of 0.1 Pa. After isolation from the vacuum pump, each gold tube was pierced with a stainless steel device to allow the product gases (C1–C5, CO2, H2S, H2, and N2) to be volatilized into the glass vacuum line. The heavier gaseous compounds were cold-trapped with liquid nitrogen (196 °C), while the permanent gases were concentrated into a pre-calibrated volume using a mercury Toepler pump. Replacing the liquid nitrogen with a mixture of dry ice and acetone (80 °C) caused release of the other volatile species, excluding water and organic compounds heavier than C5. These gases were drawn into the same calibrated volume in order to determine total gas yields, and the total number of moles of gas were calculated assuming ideal gas behavior (PV = nRT). Identification and quantification of individual HC and non-HC gas components were carried out using a twochannel Hewlett–Packard 6890 Series Gas Chromatograph (GC) that was custom-configured by Wasson ECE Instrumentation. The details of the GC operation conditions are described by Zhang et al. (2007). 2.2.4. Quantification of remaining CaSO4 after reaction After pyrolysis and cool down, the gold tube was withdrawn from the autoclave and gently scored around the middle, taking care not to puncture the tube. The tube was then frozen in liquid nitrogen (196 °C) for 5 min to completely condense all volatiles except hydrogen and helium. Next, the gold tube was removed from the liquid nitrogen, manually folded to break it along the score, and placed into a 250-mL plastic bottle filled with 50 mL of a 10% NaCl solution that was immediately capped. The bottle was agitated overnight to allow all of the residual CaSO4 to dissolve from the gold tube. Analysis of aqueous SO4 concentration was performed by MWH Laboratories using an ion chromatographic method (EPA 300.OA). This analysis was conducted on a Dionex Model DX120 ion chromatograph consisting of a gradient pump, a guard column (Dionex ionpac AG4A-SC), an analytical column (Dionex ionpac AS4A-SC) coupled with a Dionex anion self-regenerating suppression system, detectors, an autosampler, and an autoion controller.

Table 2 Experimental observation of effectiveness of mineral buffer and gold-tube inertness in the presence of high H2S at 330 °C for 24 h Two testing conditions

H2S + 400 lL 18% NaCl + dolomite H2S + 400 lL (0.25% MgCl2 + 17.8% NaCl) + Talc–silica

Amount of mineral buffer (mg)

Calculated in-situ pH

Measured gas yield (umol) H2S

CO2

H2

41.6 65.3

5 3.5

190.6 191.5

19.8 0

1.22 0

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To verify that the amount of CaSO4 used in our experiments (13 mg) is completely soluble in the 50-mL of 10% NaCl solution that was used to recover residual sulfate, we added variable amounts of CaSO4 then had the solution analyzed for dissolved SO4 concentration by MWH Laboratories. The difference between the amount of CaSO4 measured in solution and the original weight was less than 0.5 mg, which is within the measurement error of our balance. 2.2.5. Quantification of LSC in the gas phase The quantity of individual volatile and semi-volatile organic sulfur species in pyrolysates from the experiments was determined using a gas chromatograph equipped with a flame photometric detector (GC-FPD). The instrument configuration consisted of an SRI 8610 GC (SRI Instruments, Las Vegas, NV), a 1-mL gas sample loop mounted to a splitless injection port, a Gas Pro fused silica capillary column (60 m  0.32 mm ID and 1-lm film thickness), and a flame photometric detector. The capillary column had an initial flow of ultra-high purity helium of 4.98 mL min1, and the GC oven was programmed to be held at 60 °C for 2.5 min and then heated at 10 °C min1to 250 °C where it was held for 3 min. The high-sensitivity FPD was operated at room temperature. Compound identification and quantification was made against a series of standard gas analyses. 3. RESULTS AND DISCUSSION 3.1. The pH of the aqueous solution significantly affects the rate of TSR 3.1.1. pH value of minerals buffering aqueous solution at high temperature condition To investigate the effect of the pH of aqueous solutions on the rate of TSR, four aqueous solutions containing the same concentrations of Ca2+ and SO4 2 ions and different concentrations of Mg2+ were prepared. Because the solubility of CaSO4 is proportional to the salinity of the aqueous solution (Blount and Dickson, 1969), we held the ionic strength of the aqueous solutions constant (3.91 M) by adjusting the NaCl concentration. A detailed list of the concentrations of the solutes in the four aqueous solutions used in these experiments is provided in Table 3. Since H2S is soluble in water, the equilibrium between aqueous (H2S(aq)) and gaseous (H2S(g)) H2S is dependent

on the temperature, the partial pressure of H2S, and the salinity of the aqueous solution (Suleimenov and Krupp, 1994). The dissociation of aqueous H2S produces HS and H+ and drives the pH of the solution down to 2.5– 3.2 at temperatures above 300 °C (Miyasaka et al., 1989). Therefore, it is critical to select an effective mineral buffer for keeping in-situ pH within the desired range. To buffer the pH of high-temperature solutions, a talc– silica mineral buffer was added to solutions numbers 1, 2, and 4 and dolomite was added as a different mineral buffer to solution 3 (Table 3). The approach used to regulate insitu chemical conditions in this study relied on chemical reactions that are known to proceed rapidly at the temperature and pressure conditions of the experiments. For solutions 1 and 2, the precipitation of talc was used as an in-situ control for pH changes (Seewald et al., 2000) according to Eq. (1): 3Mg2þ þ 4SiO2ðsilicaÞ þ 4H2 OðlÞ () Mg3 Si4 O10 ðOHÞ2ðtalcÞ þ 6Hþ ð1Þ

Thermodynamic equilibrium within solutions containing talc and silica leads to an inverse correlation between the pH and the concentration of Mg2+ (Saccocia and Seyfried, 1990; Seewald et al., 1998) with an increase in the concentration of Mg2+ resulting in a decrease of the pH of the aqueous solution at temperatures above 200 °C (Seewald et al., 2000). To allow differences in TSR reaction rates resulting from pH variations to be distinguished from those due to variations in Mg2+ activity, dolomite was added to solution 3 (Mg2+-free solution) instead of the talc–silica mineral buffer. Fluid–dolomite equilibrium controls pH according to the reaction:
Mg2þ Ca2þ þ 2H2 CO3 () Ca; MgðCO3 Þ2ðdolomiteÞ þ 4Hþ ð2Þ It is difficult to measure the in-situ pH of aqueous solutions in pyrolysis systems, and therefore we used aqueous thermodynamic modeling (Stream Analyzer, OLI systems, Inc.) to predict the water chemistry at our experimental conditions. It should be noted that the uncertainties associated with the calculated pH values from our thermodynamic model at temperatures above 300 °C are not known due to a lack of experimental data. For this study, we have assumed that the observed trends of pH and dominant aqueous and solid species as a function of temperature are

Table 3 Ion concentrations and pH values of the aqueous solutions used in our experiments Solution No.

1 2 3 4

Ion concentration (mol/l)

Mineral buffer

Ca

SO4

Mg

Na

Cl

0.04 0.04 0.04 0

0.04 0.04 0.04 0

0.62 0.026 0 0.62

1.89 3.69 3.75 1.89

3.13 3.74 3.75 3.13

Talc–silica Talc–silica Dolomite Talc–silica

Ionic strength

pH

(M)

25 °C

>300 °C

3.91 3.93 3.91 3.75

6.2 6.2 7.8 6.3

3 3.5 5.0 3

Note: The predicted pH of solution 3 above 300 °C is in the presence of 0.1 mol hydrogen sulfide. The pH of all solutions at 25 °C was measured by an Oakton pH meter.

Thermochemical sulfate reduction by H2S initiation

unchanged at higher temperatures (>300 °C). Fig. 2 shows the predicted variation of pH with temperature in solutions 1, 2, and 3 from 125 to 350 °C. At temperatures above 300 °C the different concentrations of Mg2+ are predicted to produce pH values of 3, 3.5, and 5.0, respectively. 3.1.2. CaSO4 reduction by paraffinic mixture under different pH conditions at 330 °C A series of isothermal hydrous-pyrolysis experiments at 330 °C for 24 h was conducted using the paraffinic mixture with CaSO4 in the presence of variable amounts of H2S (10–200 lmol) at different pH conditions. Table 4 presents the results of these experiments. At pH 3 and 3.5, the rate of CaSO4 reduction is positively correlated with the amount of H2S initially loaded. Our results at these lower pH conditions are consistent with Goldhaber and Orr (1995). In contrast, CaSO4 reduction was not observed with increasing H2S loading at pH about 5.0 (Fig. 3), indicating that the reduction of sulfate might not occur, or that at least the rate is extremely slow even when large amounts of the catalyst (H2S) are present. These results might indicate that low pH conditions in aqueous solutions may play the most important role in sulfate reduction by HC. As shown in Table 4 and Fig. 3, the rate of reaction is positively correlated with the initial H2S partial pressure at low pH conditions. The following mechanism has been suggested to explain the reactivity of H2S in TSR (Feely and Kulp, 1957; Powell and Macqueen, 1984): SO4 2 þ3H2 S () 4S þ2H2 Oþ2OH 4S þ1:33ðCH2 Þþ2:66H2 O ) 4H2 S þ1:33CO2

ð3Þ ð4Þ

SO4 2 þ1:33ðCH2 Þþ0:66H2 O ) H2 Sþ1:33CO2 þ2OH ð5Þ Here S° represents sulfur in some intermediate oxidation state with low valance (Orr, 1974; Goldhaber and Orr, 1995; Seewald, 2003). The reduction of sulfate (+6 valence) to sulfite (+4 valence) is a rate-limiting step for TSR reaction, therefore reaction (3) is the rate-limiting step for the overall H2S initiation reaction because the rate of reaction (4) is rapid (Amrani et al., 2008; Ma et al., in press). Once reaction (3) starts to produce elemental sulfur or

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low-valence sulfur-intermediates, it will quickly react with HC at our temperature conditions and form additional H2S, which further drives reaction (3) to reduce sulfate. As a result, the sulfate concentration decreases as the reaction proceeds. However, our experimental results clearly show that, at pH 5.0, there is no observable sulfate reduction occurring even though a high partial pressure of H2S is present. These reaction conditions are closer to actual pH conditions within carbonate reservoirs than those where we observed TSR to occur. Therefore different mechanism must be involved for the catalysis of TSR by H2S. 3.1.3. Experiments in the absence of HC The notion that hydrogen sulfide can be oxidized by aqueous sulfate can be traced back to the experimental work of Feely and Kulp (1957) who were studying the origins of elemental sulfur deposits associated with Gulf Coast salt domes. A subsequent study by Davis et al. (1970) showed that thermodynamic considerations indicate that this reaction is favorable especially at lower pH conditions. However, even using very sensitive radioisotopic methods they were not able to observe any oxidation of hydrogen sulfide unless they intentionally allowed atmospheric oxygen to leak into their experimental system. These authors concluded that the results that Feely and Kulp (1957) reported were actually due to oxygen leakage into their experimental system (Davis et al., 1970). Some more recent reports have questioned the applicability of the experimental results of Davis et al. (1970) in natural systems citing the relatively low H2S pressure (1.5 psi) and low temperature (70 °C) that were used (Orr, 1974; Trudinger et al., 1985). Consequently, the oxidation of hydrogen sulfide by dissolved sulfate still remains a controversial subject. To further examine whether reaction (3) (i.e., H2S directly reacting with sulfate to form elemental sulfur or low sulfur intermediates) occurs at our experimental conditions, an additional series of isothermal hydrous-pyrolysis experiments were conducted. These experiments were run at 330 °C for 24 h with H2S and CaSO4 in the absence of the paraffin mixture (HC-free system), and aqueous solution 1 was

Fig. 2. Model prediction of pH values for three different solutions over the temperature range used in our experiments.

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added to buffer the pH to 3. As shown in Fig. 4, no increase in CaSO4 reduction was observed with the increase in the amount of H2S initially loaded without the presence of paraffin mixture at pH 3. The reduction of CaSO4 obviously occurs in the presence of the paraffin mixture, and the rate of reduction increases proportionally to the amount of H2S initially loaded under the same pH conditions. There are two possible explanations for the above experimental observations. One is that the initiation of sulfate reduction by the presence of H2S is not via the formation of elemental sulfur or low sulfur intermediates as proposed in reaction (3). The actual mechanism might involve the presence of both H2S and HC forming some active species that catalyzes the TSR reaction. It is possible that H2S reacts with HC to form reactive intermediates, in particular LSC, that later react with active sulfate species thereby enhancing the rate of TSR (Zhang et al., 2007; Amrani et al., 2008). A second alternative is that the rate of H2S reaction with SO4 2 or HSO4  to form elemental sulfur (3) is slow, whereas the reaction of HC with elemental sulfur (4) is fast. Any small amount of elemental sulfur produced from reaction (3) will rapidly be consumed by reaction (4). Given the relatively large measurement error for dissolved CaSO4 (±3.7 lmol) in our experiments, small changes in the amount of residual sulfate in solution may have gone undetected. More sensitive methods, such as sulfur isotopic tracers should be employed in future investigations. Our experimental observations can not conclusively rule out the possibility that elemental sulfur forms from H2S reacted with HSO4 2 HSO 4 , although there is an indication that this reaction is less significant than the HC–H2S mechanism.

Reduced CaSO4 amount (umol)

35

pH=3

25

15

pH=3.5

5 pH=5.0

-5 0

50

100

150

200

250

Loaded H2S amount (umol)

Fig. 3. Comparison of CaSO4 reduction by the paraffin mixture in the presence of H2S at three different pH conditions at 330 °C for 24 h with a confining pressure of 3500 psi.

3.2. Formation of labile organosulfur compounds (LSC) from H2S reacting with hydrocarbons Field observations have shown that TSR-altered oils are enriched in thiophenic and sulfidic compounds, which are formed in reservoirs via reactions between H2S and hydrocarbons (Cai et al., 2003; Ho et al., 1974; Manzano et al., 1997). Kelemen et al. (2008) reported that sulfur-rich organic solids from TSR-associated reservoir rocks from the Mississippian Madison Limestone and Devonian Nisku Formation are highly enriched in aromatic carbon structures, have little or no nitrogen, and contain organic sulfur almost exclusively in aromatic form. The sulfur-rich organic solids formed in the laboratory TSR simulation experiments

Table 4 Amount of CaSO4 reduction due to thermochemical sulfate reduction in the presence of H2S at 330 °C for 24 h with a confining pressure of 3500 psi H2S loading amount (lmol)

Solution type

Mineral buffer

Estimated pH

Series 1: three different pH conditions with the presence of hydrocarbon 211.6 3 Dolomite 5 95.7 3 Dolomite 37.3 3 Dolomite 11.8 3 Dolomite

Hydrocarbon source Paraffin mixture

Amount of CaSO4 reduced (lmol) 0.0 0.0 0.0 1.6

186 117 40 15

2 2 2 2

Talc–silica Talc–silica Talc–silica Talc–silica

3.5

14.3 10.0 0.0 0.0

211.6 95.7 37.3 11.8

1 1 1 1

Talc–silica Talc–silica Talc–silica Talc–silica

3

26.0 13.0 8.2 0.0

Series 2: reaction of H2 S þ HSO4  in the absence of hydrocarbon 211.6 1 Talc–silica 95.7 1 Talc–silica 37.3 1 Talc–silica 11.8 1 Talc–silica

3

No paraffin mixture

The maximum measurement error of CaSO4 is ±3.7 lmol. Note: CaSO4 = 96 lmol; paraffin mix = 5 mg; talc–silica = 30 mg each; dolomite = 30 mg; solutions = 400 lL.

1.7 0.0 0.5 0.0

Thermochemical sulfate reduction by H2S initiation

Reduced CaSO4 amount (umol)

35

25

with paraffin

15

5

without paraffin

-5 0

50

100 150 200 Loading H2S amount (umol)

250

Fig. 4. Comparison of CaSO4 reduction by the reaction of H2S and HSO4  , with and without the presence of paraffin, at pH 3 and 330 °C for 24 h with a confining pressure of 3500 psi.

are consistent with naturally occurring TSR-solid bitumens (Zhang et al., 2008). Abnormally high thiol contents have been observed in TSR-altered oils and condensates (Orr, 1974) and natural gas (Cai et al., 2003). Therefore, experimental evidence of the formation of thiols and sulfides in reactions involving H2S and hydrocarbons may provide some useful information for understanding TSR in nature. The presence in oil of aliphatic sulfur compounds such as thiols and sulfides, termed labile sulfur compounds (LSC) because of their instability at high temperatures, can significantly lower the onset temperature of TSR (Zhang et al., 2007; Amrani et al., 2008). The fact that TSR reaction rates are significantly enhanced by the presence of LSC might indicate a potential explanation for the catalysis of TSR by H2S at reservoir condition with relatively high pH. The proposed mechanism is related to the formation of labile organosulfur intermediates that are highly reactive during TSR. To investigate the formation of LSC from the reaction of H2S with HC, two sets of hydrous pyrolysis experiments were conducted over a range of temperatures from 300 to 370 °C at 0.1 °C/h heating rate, with saturated HC (paraffin mixture or hexadecane)

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reacted with H2S in the presence of aqueous solution 4 (pH buffered to 3 without CaSO4). Gas yields from these experiments, including volatile organosulfur compounds, are listed in Table 5. Measurable amounts of methanethiol (CH3SH), ethanethiol (C2H5SH), propanethiols (i-C3H7SH and nC3H7SH), and dimethylsulfide (CH3–S–CH3) were generated in these experiments. Methanethiol and dimethylsulfide are the most dominant volatile organosulfur species formed by the reaction of H2S with the paraffin mixture at 300 °C and 192 h (Table 5). Our results show that, at the same thermal stress, the yield of all quantified sulfur compounds increases as the amount of H2S that is initially loaded increases. This indicates that the production of organosulfur species from H2S reacting with HC is correlated with the H2S partial pressure or the amount loaded. As shown in Fig. 5, the absolute amount of LSC produced, as well as the relative contribution of LSC to the total amount of gas produced (on a mole basis), is directly proportional to the initial H2S partial pressure. The generation of LSC from the reaction of H2S with saturated HC is highly temperature-dependent. The gas yields from the reaction of H2S and hexadecane from 300 to 370 °C at 0.1 °C/h heating rate at pH 3 are listed in Table 5. The amounts of H2S and hexadecane used were fixed at about 37 and 221 lmol, respectively. As shown in Fig. 6, the yield of LSC from the reaction of H2S with hexadecane slightly increases as the temperature increases up to 350 °C, substantially rises from 350 to 360 °C, and then slightly increases as the temperature further increases. This relation shows that the reaction of H2S with hydrocarbons is a competitive process between the generation and thermal decomposition of LSC. The rapid increase in the yield of LSC, in particular ethanethiol and propanethiol, is closely related to the generation of ethylene and propylene from hexadecane thermal cracking at high temperatures (Fig. 6c). The higher yield of hydrogen is an indication of the enhancement of hydrocarbon cracking, which might indicate, in turn, that the following reactions take place in the reaction of H2S with saturated HC:

Table 5 Quantification of gas yields, including labile organosulfurm compounds, from the reaction of H2S with hydrocarbons at an in-situ pH 3 and a confining pressure of 3500 psi H2S (lmol)

Amount of hydrocarbon (mg)

Temperature (°C)

Gas yield (lmol/g paraffin or C16) C1

C2–C5

C2ene

C3ene

H2

CH3–SH

Series 1: H2S + paraffin + 400 lL (5.6% MgCl2 + 10% NaCl) + talc–silica mineral buffer at 300 °C for 192 h 212 5.8 300 6.30 11.36 0.08 0.06 89.07 14.08 95 5.7 300 7.98 13.13 0.08 0.05 41.45 10.60 41 6.3 300 8.59 8.70 0.05 0.04 18.70 2.48 13 6 300 10.78 9.68 0.07 0.03 3.72 0.70 0 5.9 300 2.76 6.92 0.19 0.16 7.76 0.00 Series 2: H2S + C16 + 400 lL (5.6% MgCl2 + 10% NaCl) + talc–silica 36.0 54.9 336 1.62 7.17 36.9 54.4 346 3.21 19.50 37.0 51.7 360 7.32 40.45 37.9 54.2 365 10.02 86.28 38.3 51 370 11.70 97.03

mineral 0.03 0.06 0.15 0.27 0.37

C2H5–SH

CH3–S–CH3

i-C3H7–SH

n-C3H7–SH

Sum

2.39 0.99 0.22 0.14 0.00

1.08 4.66 2.36 0.81 0.00

1.65 0.48 0.16 0.13 0.00

2.18 0.43 0.12 0.08 0.00

21.37 17.16 5.35 1.87 0.00

0.14 0.42 0.49 3.59 2.61

0.06 0.35 0.56 5.51 3.55

1.80 2.96 3.62 16.70 17.41

buffer from 330 to 370 °C at 0.1 °C/h heating rate 0.58 0.57 1.50 0.02 0.08 1.56 2.59 2.01 0.08 0.09 3.57 9.98 2.26 0.21 0.10 6.94 16.72 5.88 1.25 0.47 8.56 36.61 10.03 1.16 0.05

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Cn H2nþ2 ¼ Cn H2n þ H2 Cn H2n þ H2 S ¼ Cn H2nþ1 SH

100

ð6Þ ð7Þ

80 0.004 60

40 0.002

LS/(LS+C1-5) (%)

Labile sulfur amount (umol)

0.006

20

0.000 0

20

40

60

0 100

80

loading H2S amount (umol)

Fig. 5. Generation of labile sulfur compounds from H2S reacting with the paraffin mixture at pH 3 and 300 °C for 192 h with a confining pressure of 3500 psi. The solid line with filled circles is the sum of the yields of the five quantified labile sulfur species (CH3– SH, C2H5–SH, CH3–S–CH3, i-C3H7–SH and n-C3H7–SH); the dashed line with solid squares is the mole percentage of labile organosulfur compounds in the total produced gas.

LS yield (umol/g C16)

20.0

15.0

ð8Þ

These radicals may start a radical chain-reaction to form more radicals and olefins that will increase the thermal degradation rate of the HC. A sulfur radical mechanism has been proposed to explain the generation of oil by high-sulfur kerogens (Type II-S) at relatively low thermal maturity (Lewan, 1998), and has been postulated to be a potentially important factor in controlling the thermal stability of petroleum in sulfur-rich environments (Seewald, 2001).

10.0

5.0

335

340

345

350

355

360

365

370

375

120.0

Gas yields ( umol/g C16)

RSH ¼ R þ HS

3.3. Proposed mechanism for H2S initiation of sulfate reduction

Sum of (CH3-SH+CH3-S-CH3 +C2H5-SH+C3H7-SH)

0.0 330

100.0 80.0 C2-5

60.0 40.0

C1

20.0 0.0 330

335

340

345

350

355

360

365

370

375

10.0

Gas yield (umol/g C16)

It is important to note that the type of hydrocarbon involved can play a significant role in the reaction of H2S and HC under reservoir conditions (Machel, 1998, 2001). Nitrogen-, sulfur-, and oxygen-containing compounds (NSO) in oils are less thermally stable than saturated hydrocarbons, and aromatic hydrocarbons more thermally stable. As a result, H2S could be more reactive with NSO compounds in oils. The thermal cracking of thiols and organic sulfides has been shown to generate H2S and sulfur radical species according to the following equation (Martin, 1993; Katritzky et al., 2001):

C2ene+C3ene

8.0 6.0 4.0

H2

2.0 0.0 330

335

340

345

350

355

360

365

370

375

Tem p erature (ºC)

Fig. 6. Gas generation from hexadecane cracking in the presence of H2S at pH 3 and a 0.1-°C/h heating rate from 330 to 370 °C with a confining pressure of 3500 psi. (A) Labile organosulfur compounds generation; (B) C1 and C2–C5 gas yields; (C) H2 and (ethene + propylene) yields.

As discussed above, the H2S initiation reaction during sulfate reduction is highly pH-dependent. The HSO4  ion is the dominant sulfate species in acidic aqueous solutions. Molecular modeling calculations have shown that HSO4  ions are much more reactive than SO4 2 ions in aqueous solution; the calculated activation energy of the HSO4  ion reacting with ethane is about 55 kcal/mol, which is about 23 kcal/mol lower than that of SO4 2 ion with ethane (Ma et al., in press). Kiyosu (1980) observed that sulfuric acid and sodium bisulfate were reduced to H2S by dextrose at temperatures above 300 °C, with an initial pH of the aqueous solution in the range of 0.9–1.35 (at 25 °C). However, the reduction of sodium sulfate to H2S did not occur when the initial pH of the aqueous solution was 7. Both experimental and theoretical results suggest that the reduction of the HSO4  ion by HC is more energetically favorable than the SO4 2 ion. Because the concentration of bisulfate ½HSO4   increases with increased acidity, the rate of TSR is highly dependent on the pH of the aqueous solution. In addition to the pH of the aqueous solution, the presence of H2S is an important factor in controlling the rate of sulfate reduction. Our experimental results clearly show that the rate of sulfate reduction by HC is positively correlated with the amount of H2S initially loaded or with the partial pressure, under acidic conditions. However, at pH 5.0, the rate of sulfate reduction by HC in the presence of an excess of H2S is extremely slow and cannot be measured at 330 °C. Furthermore, the results also indicate that the reaction of the HSO4  ion with H2S to form elemental sulfur as an intermediate under acidic conditions is questionable.

Thermochemical sulfate reduction by H2S initiation

Therefore, we suggest that the oxidation of H2S to elemental sulfur by sulfate ions is unlikely to occur in the subsurface, particularly in carbonate reservoirs. Alternatively, we propose that the formation of LSC produced from the reaction of H2S with HC is likely the main mechanism for H2S catalysis of TSR. Our experimental results show that the production of LSC, such as thiols and sulfides, is positively correlated with the partial pressure of H2S. The presence of LSC in oils can significantly lower the onset temperature required for sulfate reduction (Zhang et al., 2007; Amrani et al., 2008). Labile S compounds are more reactive than inorganic S species (H2S and S8) in TSR, with the reactivity of pentanethiol being about 20 times greater than H2S at the early stages of TSR (Amrani et al., 2008). A strong correlation is observed between the formation of LSC and the enhancement of TSR reaction rate under acidic conditions. Our experimental results show, however, that even under high partial pressures of H2S the reduction of sulfate does not occur at pH 5.0, which is closer to natural conditions. Typical petroleum reservoir formation waters (pH 6.5–8.5) have low concentrations of HSO4  ions (Collins, 1975). Therefore the reaction of HC with the HSO4  ion is too slow to generate substantial amounts of H2S and elemental sulfur at reservoir temperatures. Alternative reactive sulfate species, other than the HSO4  ion, probably play a role at reservoir temperature conditions. Theoretical calculations show that the energy required to reach a transition state is the same for MgSO4 contact ion pair (CIP) and HSO4  reacting with HC (Ma et al., in press). The calculated range of concentrations of MgSO4 CIP in formation waters is much higher than HSO4  , and is sufficient to sustain TSR reaction rates that are geologically reasonable (Ellis et al., 2007). Because HSO4  concentrations in formation waters under geological conditions are

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so low, MgSO4 CIP has been proposed to be the reactive species in geologic settings (Tang et al., 2005; Ma et al., in press). We further propose a new reaction scheme of sulfate reduction by HC as shown in Fig. 7. The overall simplified reaction of hydrocarbon oxidation by sulfate may be described by three steps: Step 1. Sulfate reduction with hydrocarbons prior to H2S presence. The first condition for TSR to be initiated is the activation of the sulfate ion to form HSO4  ion or MgSO4 contact ion pair (Ma et al., in press). Once this is achieved, the reduction of sulfate by HC without catalysis by H2S may start at a slow rate. This reaction produces lower valance sulfur species including SO3, S2O3, S8 and H2S. Step 2. Sulfur incorporation into hydrocarbons by reacting with H2S. The reduced S species from step 1 react rapidly with HC and form LSC including thiols, sulfides, and disulfides and H2S. If a sufficient amount of H2S originating from (a) thermal cracking of oil (Zhang et al., 2007), (b) bacterial sulfate reduction (Machel, 2001), or (c) migration from other reservoirs (Cai et al., 2005; Zhang et al., 2005), is present, the reaction may jump over the induction period (step 1), and proceed directly to step 3. Step 3. Sulfate reduction by reactive, labile sulfur compounds oxidation. The LSC that formed in step 2 can be readily oxidized by the HSO4  or MgSO4 contact ion pair (Rudolph et al., 2003), thereby enhancing the overall rate of TSR. The lower valance sulfur species that are formed in this process will react with HC to form more H2S and LSC and thus will further sustain the auto-catalyzed reaction.

Fig. 7. Proposed mechanism of sulfate reduction initiated by H2S.

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3.4. Implications of the proposed mechanism for TSR in nature Previous reports have suggested that the formation water chemistry may play a significant role in controlling the rate of TSR (Tang et al., 2005; Ellis et al., 2006, 2007). Our experimental observations clearly demonstrate that the rate of TSR is proportional to pH value or the concentration of ½HSO4   under the same temperature condition. The lower pH value of formation water is, the faster is TSR reaction. The reactive sulfate species is HSO4  under high temperature experimental conditions rather than SO4 ¼ which was assumed as reactive sulfate species in TSR by most of previous other researchers (Machel, 2001). However, HSO4  anion concentration in the typical formation water of carbonate reservoirs in which are commonly associated with TSR may be too low to be an effectively reactive species due to the relatively high pH conditions that are maintained by carbonate mineral buffering. Alternatively, MgSO4 contact-ion-pair is proposed as a reactive sulfate species because of its lower activation energy than that of SO4 ¼ (Ma et al., in press) and higher concentration than that of HSO4  in the typical oilfield brines (Ellis et al., 2007). Consequently, the difference in water chemistry from place to place may cause a significant variation in hydrogen sulfide generation from sulfate reduction. Manzano et al. (1997) reported that sour gas condensate pools located only a few kilometers apart display a wide variation in H2S concentrations in the Nisku Formation in the Brazeau river area. A possible interpretation is those gas condensates pools may have isolated hydrological system which determines aqueous chemical compositions and leads to different reaction rate of TSR under approximately same temperature condition. The geochemical changes that Orr (1974) observed in the H2S contents and stable sulfur isotopic compositions of some of the oils and gases in the Big Horn basin, particularly the unusually high concentrations of thiols, led him to propose that a dynamic sulfurization/desulfurization process occurs in association with TSR. He also suggested that TSR may be an autocatalytic process with the generation of H2S catalyzing the further reduction of sulfate, and this concept has been substantiated by numerous experimental studies (Toland, 1960; Trudinger et al., 1985; Goldhaber and Orr, 1995 and references therein). Based on the previous experimental work of Toland (1960) and Feely and Kulp (1957), Orr (1974) indicated that the mechanism of TSR in natural environments involved the reaction of hydrogen sulfide and dissolved sulfate to form reduced sulfur species (primarily elemental sulfur and/or polysulfides). The results of the present study confirm the notions that TSR is an autocatalytic process and that organic sulfur compounds are both produced and eliminated throughout the reaction; however, a novel mechanism has been proposed for explaining how the process occurs. Our data indicate that the presence of H2S can significantly catalyze the sulfate reduction and enhance the rate of TSR by means of formation of labile sulfur compounds through H2S reacting with oils. Higher H2S partial

pressures produce greater amounts of LSC. Our previous experimental investigation showed that the onset temperature of TSR for three oils containing variable amounts of LSC decreased significantly from 415 to 355 °C with the increase of LSC from 1.1 to 17.3 mg/g oil (Zhang et al., 2007). High-sulfur oil, in particular high LSC content, may play an important role in controlling the onset temperature and rate of TSR in nature. Those observations may help geochemists to understand the origin of the Zhaolanzhuang sour gas accumulation in the Tertiary of the Jixian Sag of Bohai Bay Basin, China. Gas accumulations in this field contain H2S concentrations of up to 92%, and the associated oils contain very high amounts of sulfur up to 14.7% (Cai et al., 2005). The origin of H2S in the Zhaolanzhuang sour gas accumulation is controversial. Zhang et al. (2005, 2006) have proposed that H2S was generated by TSR occurring at a greater depth and then accumulating in the shallower, cooler reservoir by migration along faults. On the other hand, Worden and Cai (Cai et al., 2005; Worden and Cai, 2006) have provided evidence that supports the notion that the H2S was generated in-situ by bacterial sulfate reduction (BSR). Their main argument is that the temperature of the gas reservoirs range from 75 to 100 °C, which is too low for significant amounts of thermochemical sulfate reduction to have taken place (Worden et al., 1995). However, these authors have neglected to consider the effect that the high-S content in these oils might have on the onset temperature of TSR. To quantitatively predict TSR reaction rates under geologic conditions, the following factors should be considered: formation water chemistry (HSO4  or MgSO4 CIP concentration), hydrocarbon type (LSC content), H2S partial pressure (related to various sources and sinks) and temperature. 4. CONCLUSIONS To quantify the reaction of H2S with HC, we developed a new method for quantitatively loading H2S into gold-tube reactors. This method has proven to be a simple and effective means to investigate the effect of water chemistry, such as pH and H2S partial pressure, on sulfate reduction rate. These results provide new insights into the role of H2S in TSR catalysis. The most important factor in the TSR reactions in our experiments is the pH or HSO4  concentration, which activates the sulfate ion for reaction with HC. Once the sulfate ion is activated, the presence of H2S can significantly increase the rate of the TSR reaction. The initial H2S partial pressure in the system positively correlates with TSR rate. However, when HC is not present in the system, no sulfate reduction could be detected even under high H2S pressures, which indicates that the direct reduction of sulfate by H2S is not a significant pathway under our experimental conditions. These results also point to the importance of the HC as the reducing agent in TSR. Sulfur incorporation into HC during the reaction of H2S with HC (300–370 °C) results in the formation of appreciable amounts of thiols and sulfides (LSC). The formation of LSC is most likely the main path for the initiation of TSR by H2S. Based on

Thermochemical sulfate reduction by H2S initiation

our experimental observations, we proposed a new mechanism for TSR reaction, which may provide some potential implication to TSR process occurred in nature. ACKNOWLEDGMENTS This research was supported by the Joint Industrial Program on Thermochemical Sulfate Reduction at the Power, Environmental, and Energy Research Center at the California Institute of Technology. Industrial sponsors include BP, Chevron, ENI, ExxonMobil, SaudiAramco, Shell Oil and Total. We thank Dr. Hans Machel, Dr. Martin Fowler, Dr. Richard Worden and an anonymous reviewer for their critical and constructive reviews of our paper. This work benefited from collaborations and discussions with many colleagues, and we especially thank M. Haught, for his help in setting up the H2S gas-filling system. Helpful reviews of an early draft of this manuscript were provided by Dr. Paul Lillis and Dr. Zeev Aizenshtat.

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