Article pubs.acs.org/EF
Hydraulic Fracturing Additives and the Delayed Onset of Hydrogen Sulfide in Shale Gas Payman Pirzadeh, Kevin L. Lesage, and Robert A. Marriott* Department of Chemistry, University of Calgary, Calgary, Alberta, Canada S Supporting Information *
ABSTRACT: As natural gas production has shifted further from deep prolific gas reservoirs to shale gas, several questions are being addressed regarding fracturing technologies and the fate of chemical additives. A less investigated issue is the unexpected increase in produced hydrogen sulfide (H2S) from hot shale gas reservoirs. Understanding the source of H2S in shale reservoirs and managing low-levels of recovered elemental sulfur affects plans for future treatment, corrosion mitigation, and fracture fluid formulations. In this work we demonstrate that some typical ingredients of hydraulic fracturing fluids are not as kinetically stable as one might expect. Surfactants and biocides such as sodium dodecyl sulfate and glutaraldehyde are shown to undergo hydrolysis and thermochemical sulfate reduction reactions under moderate reservoir conditions, with H2S as the final product accompanied with long chain alcohols and hydrogen sulfate as long-lived intermediate species. This finding suggests that fracture fluid additives can be responsible for the delayed production of natural reservoir H2S.
1. INTRODUCTION Recent increases in unconventional gas production has greatly influenced the ability of oil and gas companies to meet the global energy demand, particularly in North America.1−4 The combination of lower-cost horizontal drilling and hydraulic fracturing technologies has been a major driving force for extensive exploration and production of gas from various lowpermeability reservoirs.5−7 A feature of many shale gases produced from hot reservoirs (T > 100 °C) is the presence of hydrogen sulfide (H2S), whose concentration may range from tens to several hundred parts per million (ppm), accompanied by various amounts of carbon dioxide (CO2), e.g., productions from Haynesville (T = 157 °C), Barnette shale (T = 82 °C), and Horn River (T = 80−160 °C) in western Canadian sedimentary basins.8−15 An anomaly observed in many H2S-producing shale gases, currently not well-documented in the open literature, is the absence of H2S in initial assessments of the (hot) reservoirs. In these cases, early fluid testing does not reveal significant amounts of H2S; therefore, facilities may be designed for sweet gas production.11,12 After production is brought on-stream, often H2S then begins to appear after a few months, where the concentration can increase up to a constant low-level (< 2%) within a period of six months (the fluid is souring). This presents a challenge for shale gas producers, as facilities must be modified to mitigate for extraneous corrosion mechanisms and treat raw gas for H2S (remove and dispose or recover in the form of elemental sulfur). While a variety of technologies can safely treat for H2S and mitigate corrosion, the corresponding modification of facilities and operation can impose an unexpected cost to producers, especially where sulfur dioxide emissions are tightly regulated.16 Thus, identifying the sources of H2S is an important aspect for future shale gas production. Four possible sources of souring in a shale gas reservoir after hydraulic fracturing are (i) H2S seeps into the producing shale from a nontargeted zone; (ii) H 2 S is generated by © 2014 American Chemical Society
aquathermolysis reaction with Kerogen, similar to H2S generation in steam assisted gravity drainage for bituminous reservoirs;17−20 (iii) H2S is native to the reservoir and is selectively adsorbed onto mineral surfaces, hence desorbing at lower pressures or later in the production life; (iv) H2S is produced by thermophilic sulfate reducing bacteria.21,22 The bacterial production of H2S is well-known; therefore, fracturing fluids often contain biocide additives before injection and/or after the fracture fluid is recovered.7,23,24 Water injected during fracturing of hot reservoirs could explain bacterial souring, as the water can contain the necessary bacteria and sulfate species; however, the presence of biocides in the fracturing fluid lowers the probability of bacteria involvement and bacterial souring does not explain why lower-temperature reservoirs do not apparently experience similar H2S increases. Aquathermolysis is a high-temperature process and is normally considered for much higher temperatures than the shale gas cases.17−20 For conventional carbonate reservoirs, in addition to aquathermolysis and/or bacterial H2S generation, the thermochemical sulfate reduction (TSR) process is another pathway for generation of H2S under hydrothermal conditions.25−34 With TSR, sulfate species are reduced under high-temperature and pressure at the expense of hydrocarbon oxidation; this process results in the formation of H2S, CO2 and lighter hydrocarbons.25−34 While TSR is a commonly accepted mechanism for the souring of conventional hot sour gas reservoirs, it is also well accepted that TSR proceeds very slowly over a geological time scale.8−12 A simplified aqueous TSR mechanism for Cx+1H2x+4 in a conventional sour reservoir is given by 3/4·x H+ + 3/4· xCaSO4 (s) ⇌ 3/4 · xCa 2 + + 3/4· x HSO4 −
(1)
Received: May 9, 2014 Revised: July 14, 2014 Published: July 20, 2014 4993
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Table 1. Chemicals Added to Fracturing Fluidsa additive type
chemical
surfactant surfactant biocide gel breaker oxygen scavenger scale inhibitor friction reducer gel iron controller pH adjusting agent cross-linker propant
sodium dodecyl sulfate isopropanol, ethanol, methanol, 2-butosyethanol gluteraldehyde, ammonium chloride, hydroxymethyl phosphonium sulfate ammonium persulfate, magnesium peroxide, magnesium oxide ammonium bisulfate ethylene glycol polyacrylamide, mineral oils, SDS guar gum, hydroxyethyl cellulose citric acid sodium or potassium carbonate borate salts silica and quartz sand
diluted acid
hydrochloric acid
a
purpose friction reduction, emulsion inhibition formulation stabilizer, winterizer eliminate bacteria delays breakdown of gel polymer until after propant is deposited removes oxygen prevents deposits within wellbore and pipelines interfacial friction reduction (slick water) viscosity modification for carrying propant chelating of metal oxides maintains the effectiveness of some compounds, such as cross-linkers viscosity modification for carrying propant (high-temperature) props fractures open for gas migration after fracture fluid has been recovered mineral dissolution for some acid stimulation
Adapted from refs 7, 20, and 21.
Table 2. Oxidation of 1-Dodecanol through TSR final
initial oxidant
T/°C
p/MPaa
t/h
103·SO42−/molb
103·SO42−/mol
105·H2S/mol
105·CO2/mol
% recoveryc
MgSO4(aq) MgSO4(aq) NaHSO4(aq) H2SO4(aq) H2SO4(aq) (NH4)2S2O8
300 300 200 200 200 200
6.5 15.7 36.9 71.2 36.9 20.9
40 40 40 40 40 40
5.06 5.06 4.68 4.68 4.68 4.98
3.60 3.95 4.01
0.47 0.52 3.0 121 136 0.0006
0.30 0.75 0.18 0.090 0.30 1.4
71 78 86 26 65 95
0.37 4.73
a
The total pressure at the reaction temperature was calculated based on the measured pressure and temperature upon opening the vessel (see Materials and Methods). bThe initial concentration of sulfate anion corresponds to a solution of 1.65 M (see refs 27−30). cSmall quantities of S2−(aq), SO32−(aq), and S2O32−(aq) were found, but omitted from the table for clarity. “% recovery” is the fraction of sulfur species accounted for after all analyses, i.e., the % sulfur mass balance.
3/4·x HSO4 − + 9/4·x H 2S + 3/4· x H+ ⇌ 3xS° + 3x H 2O
(2)
3xS° + Cx + 1H 2x + 4 + 2x H 2O → 3x H 2S + xCO2 + CH4
(3)
3/4·xCO2 + 3/4· x H 2O → 3/4· x HCO3− + 3/4· x H+
(4)
than what has been previously reported. In the present work, we have demonstrated that aqueous additives such as sulfate based surfactants and biocides can be involved in generation of H2S on a much shorter time scale than one would expect from carbonate reservoir hydrocarbons. For example, sodium dodecyl sulfate (SDS), also known as sodium lauryl sulfate (SLS), a surfactant and friction reducer, quickly hydrolyses at T > 80 °C and then undergoes TSR in a very short time (days at T = 150 °C). In a sour reservoir, the sulfate intermediate would contribute to new H2S and would rapidly remove any native H2S which would be rereleased after some time, i.e. the native H2S would be temporarily scavenged by excess sulfate. As a friction reducer, the SDS associates with the rock face and would remain in the reservoir after the fracture water has been recovered (after flowback and cleanup). Feeding of a shale gas reservoir with SDS under hydrothermal conditions would then result in a delayed generation of H2S. Other components of the fracturing fluid such as glutaraldehyde, a biocide, and ethylene glycol may facilitate TSR or may participate in parallel TSR reactions with other sulfate-producing compounds, thus resulting in the observed levels of H2S. We note that in many of the presented experiments, H2S concentrations were higher than one would expect from microbial activities.
3/4·xCa 2 + + 3/4·x HCO3− → 3/4 ·xCaCO3(s) + 3/4 ·x H+
(5)
where the net reaction for the oxidation of C2+ species is 3/4· xCaSO4 (s) + Cx + 1H 2x + 4 → 3/4· x H 2S + 1/4· xCO2 + 1/4· x H 2O + 3/4· xCaCO3(s) + CH4·
(6)
Because the hydrocarbon reduction (reaction 3) is considered rate limiting and is very slow for light hydrocarbons, heavier hydrocarbons are required for more significant TSR rates. Simple saturated hydrocarbons, such as those found in low molecular weight gas reservoirs, do not react at an appreciable rate to see changes over the life of a production. Furthermore, shale gas reservoirs rarely contain sulfate species unlike the conventional carbonate reservoirs, where anhydrite is often responsible for the reservoir seal (caprock).8−10 In this study, we considered two alternative sources for the appearance of H2S during shale gas production which both have roots in the degradation chemistry of the fracturing fluid additives (see Table 1): (1) the degradation of additives and (2) the temporary scavenging of native H2S by additive products.7,23,24,35 It is suggested that fracture fluid additives can provide both the reductants (hydrocarbons) and oxidants (sulfate) for the TSR mechanism. A required condition for these mechanisms then is that the rates of TSR be much faster
2. MATERIALS AND METHODS 2.1. Materials. Sodium dodecyl sulfate (Catalog No. S-329) and sodium bisulfate (Catalog No. S-240) were obtained from Fisher Scientific Company. Magnesium sulfate was purchased from BDH Inc. 1-Dodecanol (Catalog No. 126799), glutaraldehyde (Catalog No. G6257), ethylene glycol (Catalog No. 293237), propargyl alcohol 4994
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Table 3. Generation of H2S from Hydrolysis and TSR Reactions Involving Aqueous SLS final
initial T/°C
p/MPaa
t/h
103·SLS/molb
103·SO42−/mol
105·H2S/mol
105·CO2/mol
103·k/h−1
% recoveryc
300 250 200 200 200 150 150 150 200 200 200 200 200 200
15.1 20.9 26.2 17.0 27.6 28.7 37.2 16.7 18.5 19.0 19.6 18.4 16.2 16.9
40 40 40 40 40 40 40 40 132 132 132 132 132 132
6.32 6.38 6.34 6.35 6.35 6.35 6.35 6.33 4.92 3.84 2.76 1.68 0.634 0.634
2.95 4.31 5.77 5.32 5.51 4.92 4.76 4.69 4.28 3.48 2.53 1.61 0.51 0.53
69 16 0.38 0.20 1.1 0.084 1.7 0.13 0.71 0.45 0.18 0.12 0.009 0.024
90 4.9 3.8 3.4 5.2 4.0 3.9 3.6 3.7 2.4 1.8 1.1 0.61 0.50
5.08 1.92 0.17 0.05 0.33 0.02 0.26 0.02
60 71 91 84 87 77 75 74 87 91 92 96 81 86
a
The total pressure at the reaction temperature was calculated based on the measured pressure and temperature upon opening the vessel (see Materials and Methods). bThe initial concentration of sulfate anion corresponds to the initial concentration of SDS ranging from 0.15−1.5 M. cSmall quantities of S2−(aq), SO32−(aq), and S2O32−(aq) were found but omitted from the table for clarity. ‘% recovery’ is the fraction of sulfur species accounted for after all analyses, i.e., the % sulfur mass balance. targeted isobaric condition of p = 24.1 MPa (3500 psi).26−32 Total pressure was increased at room temperature using UHP N2, where initial pressures were targeted using the compressibility calculations from NIST’s RefProp software.47 A schematic diagram of the experimental setup is provided in the Supporting Information. The vessels were wrapped in a preheated resistive heating jacket, and the jacket temperature was adjusted accordingly to maintain a target temperature for the vessel. After the required reaction time, the vessel was removed from the jacket and air-cooled to the room temperature for sampling. Note that for the high SDS concentrations, significant general acid attack of the vessels occurred; therefore, vessels were replaced for the lower concentrations, where attack was less noticeable. Inductively coupled plasma (ICP) analyses of the aqueous phase confirmed some titanium dissolution. 2.6. Analytical Methods. 2.6.1. Gas Chromatographic Analysis of Reaction Mixture Headspace. After cooling to room temperature, each reaction vessel was connected to a sampling manifold, ca. 1.7 mL, to measure the final pressure of the product gas mixture. These data were combined with mole fractions obtained from GC analysis in order to estimate the density of the gas mixture using the software RefProp and calculate the pressure prior to opening.47 Aliquots of the pressurized subsample were released to a Varian CP-3800 gas chromatograph equipped with a thermal conductivity detector (TCD) and a pulse flame photodetector (PFPD) for sensitive detection of sulfur species (Restek Rt-U-Bond 30 m × 0.53 mm I.D. (part no. 19750) and GS-GasPro 30 m × 0.32 mm I.D. (part no. 1134332), respectively). The columns were kept at T = 60 °C for 5 min, and then the temperature was raised to T = 190 °C at a rate of 20 °C min−1. 2.6.2. Ion Chromatographic Analysis. After sampling the headspace gas phase, the vessel was repressurized with nitrogen, and the liquid content was then extracted from the vessel. The aqueous phase normally was expunged from the vessel first. A 100 μL sample of the aqueous phase was added to 10 mL of a mixture solution of 10 mM mannitol and 50 mM sodium hydroxide, to prevent oxidation of sulfide anions. A Dionex DX320 with a IonPac AS17 hydroxideselective anion-exchange column was used for anion analysis. Samples were run in 6 mL vials by a Dionex automated sampler equipped with a CD25 conductivity detector and a parallel AD25 absorbance detector. Deionized-deoxygenated water was used for preparation of potassium hydroxide solutions used in Dionex EG40 eluent generator. A concentration gradient of potassium hydroxide was applied ranging from 30 to 70 mM. 2.6.3. Gas Chromatography−Mass Spectrometry (GC−MS). Upon extracting the materials from the vessel, the reaction vessel
(Catalog No. P50803), and ammonium persulfate (Catalog No. 248614) were purchased from Sigma-Aldrich. All chemicals were used without further purification. 2.2. Preparation 1-Dodecanol for TSR. To investigate the TSR reaction involving 1-dodecanol in the aqueous phase, the solubility limit of magnesium sulfate at room temperature was considered (25.5 g/100 mL); the mole ratio of water to magnesium was kept at ca. 33.5, compared to previous studies of ∼17 in the solid phase.26−32 The mole ratio of water to hydrocarbon was kept near that of previous literature reports, ca. 35.29−32 All water was polished to 18.2 MOhms and degassed under a vacuum for 12 h. The reaction mixture was prepared by evacuating a 5 mL high-pressure titanium reaction vessel and allowing 1 mL of hydrocarbon to be drawn in through a single valve. This was followed by 3 mL of 1.65 M magnesium sulfate, sodium bisulfate, or sulfuric acid solutions into the same titanium reaction vessel. After liquid injection, ultrahigh-purity nitrogen (99.998% purity from Praxair) was used to generate total pressures, which would reproduce the similar high-pressure conditions at reaction temperature.26−32 The reaction was carried out for 40 h at T = 200.0 and 300.0 °C, as listed in Table 2. Reactions were quenched by allowing the temperature to return to room temperature and analyzed within 24 h. 2.3. SDS Solution Preparation for TSR. For comparable experiments with SDS, mole ratios similar to those in 1-dodecanol experiments were used; SDS solutions of 1.58 M were prepared and 4 mL of the solution was drawn into the reaction vessel. The reactions were carried out at T = 150.0, 200.0, 250.0, and 300.0 °C for 40 h, as listed in Table 3. SDS solutions of 0.15−1.23 M were prepared to explore the dependence of TSR reactions on SDS concentration. The latter experiments were carried out at T = 200.0 °C for 132 h. 2.4. Preparation of Aqueous Additives for TSR. The molar ratios of water to hydrocarbon and water to sulfate anion were kept as above. The solutions of glutaraldehyde, ethylene glycol, and propargyl alcohol with ammonium persulfate were freshly prepared and kept in refrigerator to minimize any potential reactions prior loading into the reaction vessels. 2.5. Experimental Setup. The reaction vessels consisted of hollowed-out 1″ female grade II titanium HiP plugs (HP pressure gland fitting) mated to a 1″ to 1/16″ adapter fitting. A titanium 1/16″ line then connected to low volume SSI UHPLC needle valve for sampling and reactant charging. The reaction vessel was connected to a sampling manifold containing a transducer during pressurizing the vessel and sampling the headspace gas mixture (depressurizing) before and after the reaction, respectively. All reactions were carried out at a 4995
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was opened, and the interior of the vessel was rinsed with chloroform. These rinsing aliquots were added to the aqueous phase and shaken in order to partition more organic material from of the aqueous phase. The final solution was left for at least half a day to settle. After settling, a 10 times-diluted sample of the chloroform phase was prepared and injected into a Bruker Scion-SQ GC-MS with a 15 m × 0.25 mm I.D. BR-5 ms column. A thermal gradient program was applied; initial temperature of GC oven was held at 50 °C for 1 min and then increased up to T = 320 °C at a rate of 10 °C min−1. Data collection started after 3 min of injection and continued for 34 min. 2.7. Kinetic Modeling. In our simplified model, we assumed (i) an instantaneous rate of hydrolysis (reaction 7), (ii) a fast reaction between sulfate and native H2S (rapid reaction 8), and (iii) a rate limiting oxidation of alcohols (reactions 9 and 10). The rate constants for reactions 9 and 10 were given the same value, which was 5−8 orders of magnitude less than the rate constant for reaction 7. We have considered a very simple bed breakthrough reaction to simulate a simplified plug flow reservoir with stationary sulfate. For this calculation, the bed begins with excess SDS and receives a constant feed of low level H2S gas, followed by a similar volumetric production (bed exit).
be long-lived and/or be reacting with native H2S to delay the production of H2S. We do note that reactions 9 and 10 would be much more complicated than presented here, and one would expect several intermediate species beyond two simple alcohols. Initially, we investigated the TSR reaction of 1-dodecanol as the product of SDS hydrolysis in the presence of various sulfate species (see Table 2). Earlier works have demonstrated that TSR initiation and progress can occur without any initial elemental sulfur, sulfide species, or H2S,26−34 and many previously investigated TSR experiments have been carried out in the presence of excess solid magnesium sulfate as the source of sulfate species. In order to investigate TSR of hydrocarbons in aqueous media, while keeping molar ratios of species close to formerly reported experiments, a relatively concentrated aqueous solution of magnesium sulfate was prepared (see Materials and Methods). Previous literature studies were mainly carried out at T ≥ 330 °C.26−33 This temperature is not realistic for an actual shale reservoir and is close to the critical temperature of water, Tc = 373 °C; however, as discussed, most TSR reactions require excessive temperatures to increase the reaction rates in laboratory studies.26−33 Therefore, our experiments also began at hightemperature (T = 300 °C) in order to probe TSR rate in aqueous phase with 1-dodecanol, and then the reaction temperature was lowered to T = 200 °C (not ideal but closer to a hot reservoir) in order to verify our experimental techniques. Results within Table 2 suggest that 1-dodecanol is capable of undergoing TSR reactions in the presence of different initial sulfate species. As expected, sulfuric acid showed the highest yield of H2S compared to sodium bisulfate and magnesium sulfate; i.e., acidic conditions favor reaction kinetics. This is in agreement with previous TSR studies where dextrose was the reductant.40 The major components found in the gas phase were H2S and CO2. Rarely, smaller hydrocarbons, such as propane or derivatives of hexane, were observed, but not of analytical significance (concentrations less than 50 ppm for non-CO2 species detected on TCD and less than 1 ppm for sulfur compounds on PFPD are considered as “not analytically significant” in this study). Figure 1 shows representative GCMS chromatograms of the 1-dodecanol mixture after the TSR reaction. In the presence of magnesium ions and only sulfate species (higher pH), the population and diversity of different byproducts were larger, indicating the formation of (and longevity of) a variety of intermediate sulfur species with different oxidation states as well as the formation of smaller hydrocarbon species through oxidation of 1-dodecanol. These intermediate species are apparently short-lived with low-pH conditions, indicating that the oxidation of the hydrocarbons is acid catalyzed. Table 3 summarizes the results of TSR for SDS solutions with no additional source of sulfate, where only aqueous SDS was exposed to higher temperatures and pressures (T = 150− 300 °C). In order to keep the mole ratios of hydrocarbon and sulfate ions similar to those reactions in Table 2, a concentrated solution (45 wt%) of SDS was initially investigated. Our results show that the TSR reactions involving only aqueous SDS are much faster than the 1-dodecanol systems at T > 250 °C. As opposed to the initial reaction mixture, which was a clear homogeneous single-phase liquid solution, the extracted liquids were found to be two-phase mixtures with a dark oily layer floating on the top of a clear aqueous layer. In contrast, materials extracted from reactions performed at T ≤ 200 °C
3. RESULTS AND DISCUSSION 3.1. SDS Hydrolysis and TSR. SDS is a compound with vast applications in different fields of science and technology including biology, pharmaceuticals, and cosmetics as well as in chemistry such as oil and gas. It is well-known that aqueous SDS decomposes through hydrolysis to 1-dodecanol and sodium bisulfate;36 see reaction 7. It has been argued that this hydrolysis reaction is nearly completed at T = 80 °C after ca. 12 h.36 The hydrolysis of SDS is an autocatalytic process, where the presence of 1-dodecanol, low pH, and high-initialconcentration of SDS (above critical micelle concentration) accelerate the hydrolysis process.35−39 In addition, earlier studies have shown that alcohols can undergo TSR reactions and produce comparable amounts of H2S with respect to equivalent unsaturated alkenes (at very low reaction rates).26 Finally, it has been suggested that bisulfate anion (HSO4−) rather than sulfate anion (SO42−) is the more reactive species involved in TSR that is eventually reduced to form a sulfide species; i.e., reaction kinetics are favored by acidic conditions.26−34 Combining all these physicochemical features, it appears that SDS itself has all the required reactants to facilitate the TSR reaction under the right hydrothermal conditions. Following reactions 1−5, a simplified reaction mechanism for the degradation of aqueous SDS in a sour shale reservoir is as follows: H 2O + C12H 25SO4−(ads) + Na + → C12H 25OH + HSO4− + Na + (7)
HSO4− + 3H 2S + H+ ⇌ 4S° + 4H 2O
(8)
3S° + C12H 25OH + 2H 2O → 3H 2S + CO2 + C11H 23OH
(9)
S° + 1/3C11H 23OH + 2/3H 2O → H 2S + 1/3CO2 + 1/3C10H 21OH
Na + + CO2 + H 2O ⇌ NaHCO3 + H+
(10) (11)
where the net reaction is NaC12H 25SO4 (aq) + 2/3 H 2O → NaHCO3 + H 2S + 1/3 CO2 + 2/3 C11H 23OH + 1/3 C10H 21OH
(12)
If reactions 9 and 10 are rate limiting reactions (hydrolysis reaction 7 is fast), then the intermediate bisulfate anion could 4996
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Figure 2. GC-MS chromatograms for the liquid-phase products of selected sodium dodecyl sulfate TSR experiments. The peaks corresponding to 1-dodecanol can be seen at approximately the same retention time as in Figure 1. Chromatograms represent the first eight reactions listed in Table 3.
Figure 1. GC-MS chromatograms for liquid-phase products of 1dodecanol TSR experiments. Representative chromatograms of the nonaqueous liquid phase of TSR reaction products of 1-dodecanol with MgSO4 at T = 300 °C, NaHSO4 at T = 200 °C, and H2SO4 at T = 200 °C after t = 40 h. The peaks corresponding to 1-dodecanol have been highlighted with a star (t ≈ 10.5 min). Chromatograms represent the reactions listed in Table 2.
and the higher temperatures for the Horn River. The rate of this reaction was much faster than we expected. The concentration of SDS in fracturing fluids is not nearly as large as what has been explored in experiments reported in Table 3. Another set of experiments were carried out with more dilute SDS solutions, 4.5−35 wt%, at T = 200 °C for ca. 5.5 days. Again, H2S and CO2 were detected in headspace gas mixture samples. The result of 4.5 wt % is comparable with those generated from 10 times more concentrated SDS solutions at T = 150 °C (see Table 3). The results of this set of experiments suggest that SDS can hydrothermally undergo TSR reaction even at lower concentrations; however, actual injected SDS additive is normally a proprietary value and not known in this case. Furthermore, local concentration of SDS in a near-wellbore region and after water flow-back may be quite large. Figure 3a reveals that production of H2S appears to have a second-order dependence on SDS initial concentration. Yet, we suggest more accurate pressure control is required in order to confidently demonstrate the dependence of TSR reaction on SDS concentration. Therefore, for simplicity, it was assumed that sulfate reduction (or formation of H2S) is a first-order reaction with respect to SDS; accordingly, a temperature-
were primarily aqueous with foamy residues on the top of the aqueous phase (1-dodecanol and water). Sampling the headspace gas of these reactions revealed the presence of H2S and minor concentrations of various sulfur species such as methanethiol, ethanethiol, isomers of propanethiol, and sometimes carbonyl sulfide. In all cases, the initially measured pH of the system was within the range of pH = 7−8, while it decreased to a value of pH = 1 after the reaction was quenched, indicating the rapid hydrolysis of SDS according to reaction 7.35 The corresponding GC-MS data in Figure 2 also confirms formation of 1-dodecanol (due to hydrolysis of SDS) as well as its derivatives, e.g., ethers and esters. As the temperature decreased, the rate of H2S and CO2 production decreased, and the resulting hydrocarbon species underwent less oxidation. At lower temperatures, there was a drastic decline in the amounts of various nonsulfate sulfur species detected in the aqueous phase (see Table 3). A very important observation is that we have been able to detect levels of H2S from TSR reactions carried out at hot reservoir conditions (∼150 °C), which is the first observation of its kind to the best of our knowledge. In fact, this temperature is typical in the Haynesville shale gas play 4997
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phase partitioning; however, reactions are expected to proceed in the aqueous phase only. Assuming SDS hydrolysis is faster under current experimental conditions and that the bisulfate anion (pKa of 1.78) completely dissociates at the beginning of the reaction, one can simply approximate SDS number of moles with the produced number of moles of sulfate anions in the system.36 Figure 3b presents the results of lnk versus the inverse of temperature (1/T). Fitting the Arrhenius equation, the estimated activation energy for TSR reaction of SDS was found to be ca. 83 ± 7 kJ mol−1 (a very similar number, 80 ± 11 kJ mol−1, was estimated with a second-order dependence on SDS concentration). This range of activation energy has previously been reported for TSR reactions in the presence of NaHSO4, H2S, and NaHS within the pH range of 2−9, having a reduced sulfur species as a catalyst,41 but appears to be much smaller than the previously estimated activation energy numbers for TSR in the presence of magnesium cation.26−34 Therefore, the relatively low activation energy and observation of thiol species suggest the formation of elemental sulfur as an intermediate compound, which has been previously identified as the intermediate species of a TSR reaction.34 It should be noted that the acidity of the environment can also greatly reduce the TSR activation energy (also see Table 1), as this is the case in SDS solutions due to dissociation of the bisulfate anion. Comparing the results of TSR reactions of SDS with those of hydrocarbons in Table 2 reveals interesting features. The amount of H2S produced from SDS at temperatures above 250 °C is comparable with those of 1-dodecanol and sulfuric acid at T = 200 °C. This suggests that the hydrolysis reaction is not the rate-limiting reaction. Again, at lower temperatures, elemental sulfur reaction with hydrocarbon has been suggested as the rate-limiting reaction.34,42 The generated elemental sulfur can undergo further reactions to produce various intermediate sulfur compounds, such as alkylthiols,43 and traces of these species were observed in headspace sampling. 3.2. TSR Reaction with Other Selected Fracturing Additives. As summarized in Table 1, fracturing fluids contain various aqueous additives such as glutaraldehyde, ethylene glycol, and ammonium persulfate utilized as biocide, scale inhibitor/winterizing agent, and gel breaker, respectively. One can speculate that those compounds not containing sulfate can participate in TSR reactions as reducing agents and those containing sulfate can participate as both reducing and oxidizing agents.7,23,24,44 A third set of experiments were carried out in order to investigate the possibility of TSR reactions of glutaraldehyde (GLA), ethylene glycol (EG), and
Figure 3. H2S production from aqueous SDS at high pressure. (Top) Produced number of moles of H2S shows a second-order dependence on the initial concentration of SDS. (Bottom) Using Arrhenius equation and assuming eq 13 holds, an activation energy of 83 ± 7 kJ mol−1 has been estimated for the overall thermochemical sulfate reduction reaction of SDS. Data points are from Table 3; two outliers are not plotted, and accordingly are eliminated for slope calculation.
dependent reaction rate constant (k) has been estimated as follows: Δn H2S Δt
= knSDS, Δt ⇒ k =
Δn H2S nSDS, Δt Δt
(13)
where ΔnH2S, nSDS,Δt, and Δt are the number moles of produced H2S, number moles of SDS consumed, and reaction duration (which was 40 h), respectively. These concentrations ignore the
Table 4. Oxidation of Other Fracture Fluid Additives Reduction TSRa final
initial reactants GLA + MgSO4 EG + MgSO4 GLA + (NH4)2S2O8 EG + (NH4)2S2O8 PA + (NH4)2S2O8
T/°C
p /MPab
t/h
103·SO42−/molc
103·SO42−/mol
105·H2S/mol
105·CO2/mol
% recoveryd
200 300 200 300 200 200 200
15.6 16.3 17.4 17.4 30.9 32.7 18.9
40 40 40 40 40 40 40
6.59 6.59 6.76 6.76 6.65 6.67 6.64
6.25 5.75 5.39 5.53 1.22 4.26 2.33
n.d. 0.14 0.0005 0.0002 52 28 1.4
17 51 0.019 52 61 46 67
95 87 80 82 26 68 35
a
GLA, EG, and PA are gluteraldehyde, ethylene glycol, and propargyl alcohol, respectively. bThe total pressure at the reaction temperature was calculated based on the measured pressure and temperature upon opening the vessel (see Materials and Methods). cThe initial concentration of sulfate anion corresponds to the initial concentration of 1.65 M. dSmall quantities of S2−(aq), SO32−(aq), and S2O32−(aq) were found but omitted from the table for clarity. ‘% recovery’ is the fraction of sulfur species accounted for after all analyses, i.e., the % sulfur mass balance. 4998
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propargyl alcohol (PA) in the presence of magnesium sulfate and ammonium persulfate; see Table 4 for results. Both GLA and EG were found to undergo TSR reactions with magnesium sulfate at T = 300 °C with H2S as the product. However, H2S production significantly declined at T = 200 °C, particularly in the case of glutaraldehyde. Regardless, these experiments provided evidence that both molecules are capable of producing H2S under relatively mild (more neutral pH) but hightemperature aqueous conditions (T = 200 °C). Like SDS, ammonium persulfate is hydrolyzed in an aqueous solution.45,46 Reaction of 1-dodecanol with ammonium persulfate (see Table 2) generated a very small amount of H2S, likely due to higher pH values of the solution (pH = 2.30 at 25 wt% in our experiments) and absence of a possible catalyst such as magnesium ion.45,46 In contrast, all GLA, EG, and PA with ammonium sulfate produced a significant amount of H2S at 200 °C. The high H2S yields could be due to the presence of combined multiple functional groups in each of the reacting hydrocarbons. Having multiple hydroxyl, aldehyde, and/or unsaturated carbons may facilitate activation of the hydrocarbons with the sulfate anion in the acidic environment.25 It is noted that the concentrations of additives in the non-SDS experiments are comparable to our earlier SDS experiments (particularly the 45 wt% SDS solutions). The results point to the fact that components of fracturing fluid are capable of undergoing TSR reactions and generate H2S independent of the native sulfate content in the reservoir. An important step in fracturing process is the recovery of the water used in fracturing the well (flowback). Although it has been reported that a great portion of the initially injected water can be recovered in a short time after production begins, the portion of the fracturing chemicals that is recovered has not been well investigated/reported.5,24 As friction modifiers such as SDS are designed to adsorb to the interface (mineral), after water flow-back, the in-reservoir concentrations during production will be greater than the concentrations injected.23 Thus, it is expected that SDS and other aqueous additives can be a finite source of produced H2S; however, without production information specific to individual wells the overall H2S originating with the SDS is unkown. These additives can still play a large role in delaying the onset of H2S which is native to the reservoir before fracturing. 3.3. SDS and the Delay of H2S Breakthrough. Though fractured reservoir production is very complicated, we did investigate a simplified kinetic model based on reactions 7−10. Figure 4 shows the result of excess SDS in this breakthrough model at three different temperatures. While further work is required to obtain more detailed kinetic parameters and to simulate real reservoir conditions, these exploratory calculations show the type of profiles one might expect from this delayed H2S production mechanism. Figure 4 shows the delay (temporary scavenging) and peak in H2S production which can indicate the overall rates of TSR in the actual reservoir. For the hottest condition, we predict short H2S breakthrough time followed by a peak in H2S production before returning to the native H2S level. A hot condition is analogous to the Haynesville Shale Gas. At lower temperatures (e.g., Barnett or Horn River), one might expect longer delays before H2S production and a smaller (broad) peak in H2S. Realtime production information from shale producers would greatly benefit the investigations in this area. Future high-pressure breakthrough studies, possibly by following isotopic ratios, will help verify this type of production profile. We also note that
Figure 4. A simplified H2S production model based on the hydrolysis and thermal sulfate reduction of SDS (reactions 7−10). This is a fixed bed model, where the bed contains excess SDS and receives a fixed concentration of H2S.
reservoir flow is much more complicated, where one might expect multiple peaks in H2S production. Other potential issues as a result of TSR with SDS could include temporary increases in general acid corrosion, increased levels of produced thiols and/or excess amine foaming with the amine treatment (H2S separation process) if any long chain alcohols are introduced to the treatment facilities.
4. CONCLUSIONS In conclusion, these experiments support the concept that, in addition to the release of H2S native to a reservoir and that generated through potential sulfate reducing bacteria, the chemical additives within fracturing fluids are capable of generating H2S under hydrothermal conditions or delaying the production of native H2S; i.e., many fracture fluid additives can be responsible for the souring of sweet reservoir fluids without bacterial activity. In this regard, either a different set of ingredients, particularly surfactants and biocides, can be considered or the appropriate H2S treating and recovery processes can be added to design criteria based on offset data (nearby well production information). Both these options affect the value of the gas in-place for a producer and the steps required by regulatory agencies. To further this understanding, our current studies are focusing on (i) the selective highpressure adsorption of H2S onto shale minerals, (ii) further kinetic studies leading to the modeling of souring kinetics (possibly catalyzed by mineral surfaces and with better pressure control), (iii) controlled pH experiments, (iv) variable ionic strength, (v) identification of the TSR products for SDS and other additives in the fracturing fluid, and (vi) alternative sulfur recovery processes for low-level H2S fluids. Finally, these results are important for addressing the fate of fracture fluid additives.
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ASSOCIATED CONTENT
S Supporting Information *
SI-Figure 1: A schematic presentation of the experimental setup used for hydrothermal thermochemical sulfate reduction (TSR) reactions. This material is available free of charge via the Internet at http://pubs.acs.org. 4999
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The authors are grateful for Discovery Grant support from the Natural Science and Engineering Research Council of Canada (NSERC).
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