Gas Hydrates Phase Equilibria and Formation from High

Article pubs.acs.org/jced

Gas Hydrates Phase Equilibria and Formation from High Concentration NaCl Brines up to 200 MPa Yue Hu,† Taras Y. Makogon,‡ Prasad Karanjkar,§ Kun-Hong Lee,∥ Bo Ram Lee,*,†,∥ and Amadeu K. Sum*,† †

Hydrates Energy Innovation Laboratory, Chemical & Biological Engineering Department, Colorado School of Mines, Golden, Colorado 80401, United States ‡ Flow Assurance Department, Wood Group, Houston, Texas 77084, United States § ConocoPhillips, Production Assurance, Bartlesville, Oklahoma 74004, United States ∥ Department of Chemical Engineering, Pohang University of Science & Technology, Cheongam-Ro, Nam-Gu, Pohang, Gyeongbuk 37673, Korea S Supporting Information *

ABSTRACT: Gas hydrate phase equilibrium and kinetics at high NaCl concentrations (near and at saturation in solution) and very high pressures (up to ∼200 MPa) are investigated to study the interplay of hydrate formation and salt precipitation. Limited experimental data above 80 MPa exist for hydrate phase equilibrium in high salinity systems. This study reveals the unusual formation of gas hydrates under these extreme conditions of high salinity and very high pressure. In particular, the results demonstrate that hydrates can form from saturated salt solutions, and the formation of hydrates and salt precipitation are competing effects. It is determined that hydrates will remain in equilibrium with a saturated salt solution, with the amount of salt precipitation determined by the amount of hydrates formed. These data are essential fundamental data for gas hydrates applications in the oil and gas production flow assurance and seawater desalination.

1. INTRODUCTION Gas hydrates are formed by the inclusion of small guest molecules inside cavities formed by hydrogen-bonded water molecules. The guest molecules generally consist of light hydrocarbons, such as methane (CH4), ethane (C2H6), and propane (C3H8), the size of which determines the crystal structure of the hydrate formed.1 Fundamentally, a hydrate can be considered a fourth phase of water, as it possesses a similar hydrogen-bonding network of water molecules as in ice. Moreover, hydrates represent a delicate balance of guest−host interactions that compensate the entropic losses with the energetic gain in enclosing hydrophobic molecules within a water network. It is well-known that salts, such as sodium chloride, calcium chloride, etc., are hydrate inhibitors. Because of strong electrostatic forces, water becomes tightly bound in the ion’s hydration shell and becomes unavailable to hydrogen bond to other water molecules and form the hydrate.2 As such, salts suppress the hydrate formation temperature at a given pressure.1 The degree of the suppression temperature depends on the salt type and its concentration in solution. While it is well established that hydrates can form from salt solution at moderate concentration, it is not well-known how hydrates may © XXXX American Chemical Society

form from near to saturated salt solutions. For salt solutions near the saturation concentration, any formation of hydrates would have to be coupled with salt precipitation or create supersaturation conditions. As such, the interplay of hydrates and salt precipitation is fundamentally interesting and important, as it involves the precipitation of two solids. More practically, hydrate formation from high salinity solutions is increasingly important during development and production of oil fields at ultradeep water depths, especially presalt production developments in the Gulf of Mexico. Field developments in ultradeep water encounter harsher conditions in terms of pressure and water salinity, posing significant flow assurance challenges, especially due to gas hydrate formation and scale precipitation in pipelines. These conditions can lead to the accumulation of solids in pipelines, subsea-transfer lines, and chokes, and eventually block them, causing production loss and placing flow lines at a risk for pressure buildup.3 The presalt reservoirs, a layer of oil bearing rock of carbonate composition, are positioned below sea level with depths Received: March 24, 2017 Accepted: May 17, 2017

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DOI: 10.1021/acs.jced.7b00292 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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between 5000 m and 6000 m, in ultradeep water (1900 m to 2400 m) beneath a thick salt layer (in some areas, up to 2000 m), resulting in reservoir pressures higher than 100 MPa and salt concentrations more than 20 wt %.4 Moreover, heavy brines are used during well completion to offset the high reservoir pressures with the brine hydrostatic pressure in order to avoid a well blowout. At the same time completion brines must remain hydrate-free at mudline (seabed) conditions in case light hydrocarbon gases migrate from the bottom of the well up the wellbore and come in contact with cold (∼277 K) brine, if a well is drilled in the offshore environment.5 If a chemical, such as pure methanol or a low dosage hydrate inhibitor formulated in methanol, is used to control hydrate formation, this may cause a salting out effect as scale forms due to colligative properties of water and inability to dissolve both salt and methanol simultaneously beyond the solubility limit.6 Even though it is imperative to know and understand the hydrate phase equilibrium conditions in near-saturated salt solutions and at very high pressures, there is little literature data available to fulfill these conditions due to challenges associated with experimental measurements under such severe conditions. Consequently, as shown in Figure 1, the majority of existing data are for methane hydrate phase equilibria with undersaturated NaCl solutions at relatively low pressures, below 80 MPa.7−12

Figure 2. (a) Schematic phase equilibrium plot for methane hydrate with increasing salt concentration up to saturated conditions. (b) Illustrative phase diagram for NaCl + water + methane showing the temperature limits for methane hydrates. The symbols shown in the plots correspond to the conditions for hydrate phase equilibrium of brine−hydrate−vapor. The quadruple points (brine−hydrate−vapor− salt) end at the salt solubility line.

that is based on the binary (H2O−NaCl) phase diagram and with the following assumptions: (i) methane is immiscible in the aqueous solution at all pressure conditions. This assumption implies that dissolved methane does not affect the solubility of NaCl in water, and conversely, NaCl has a negligible effect on the solubility of methane in water; (ii) salt solubility line is independent of pressure, which assumes that pressure has no effect on salt solubility; and (iii) there is excess liquid or gas is the limiting phase. For the case with excess gas, the salt solubility curve does not exist below the quadruple point and all solution would be converted to hydrate and solid salt. If we superimpose the hydrate equilibrium conditions on the phase diagram, for under-saturated solutions (0, 12, 23, and 26 wt % NaCl at 293.15 K), the temperatures of hydrate phase equilibrium are represented by open squares and circles, which correspond to three-phase (L−H−V) equilibrium at 65 and 185 MPa, respectively. According to Schreinemaker’s Rule, N univariant loci can emanate from any invariant point, where N is the number of phases which coexist in the invariant point.14 The four univariant loci (CH4(g)−NaCl(s)−solution(l), CH 4 (g)−solution(l)−hydrate(s), CH 4 (g)−NaCl(s)−hydrate(s), and solution(l)−NaCl(s)−hydrate(s)) emanate from the quadruple point indicated by the solid squares (65 MPa) and circles (185 MPa), which are obtained from oversaturated

Figure 1. Existing literature data7−12 of methane hydrate phase equilibria in the presence of NaCl. There are no data in the open literature for the region above 80 MPa.

This study introduces the development of a new system for hydrate formation measurements for pressures up to 200 MPa and suitable for high salinity solutions. This system is used to obtain hydrate phase equilibria and hydrate formation rates at these extreme conditions of very high pressures and near/at saturation salt solutions. To understand the effect of high NaCl salinity on the stability of methane hydrates, experimental measurements of hydrate phase equilibria were performed for undersaturated and oversaturated NaCl solutions; NaCl saturation in water is ∼26.4 wt % at 293.15 K.13 Depending on the NaCl concentration, the measured data either correspond to three (CH 4 (g)−solution(l)−hydrate(s)) or four (CH 4 (g)− NaCl(s)−solution(l)−hydrate(s)) phase equilibrium. Figure 2 shows the phase diagram13 for the CH4−H2O−NaCl system B



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Figure 3. Experimental setup for hydrate phase equilibrium studies with high salinity solutions and very high pressure conditions. Because of corrosion, the protective cylinder (item L), baffle (item M), and RTD (resistance temperature detector) sensor (item I) are made of Inconel 600.

solutions (27 and 30 wt % NaCl, total salt amount).14 In the measurements with the oversaturated NaCl solutions, only one temperature is actually measured for the hydrates dissociation at constant pressure because the solution must be in equilibrium with the hydrate at the saturation concentration; any excess NaCl is precipitated. On the basis of the phase diagram shown in Figure 2, we report the effect of total NaCl content on the hydrate phase equilibrium and growth rate of hydrates.

connected with a refrigerated/heating circulator. The cell contents were stirred at 500 rpm using a magnetic bar and strong magnetic plate. The baffle in the cell allows for a good mixing at the gas−liquid interface. After flushing the system three times with the experimental gas to remove air and impurities, the system is pressurized with methane (99.97 mol % purity, General Air) to the desired pressure with a gas booster pump. Measurements for the hydrate phase equilibria used an isochoric method, which is based on the variation of the pressure with temperature for a fixed volume system. As shown in Figure 4a, temperature and pressure data are monitored based on the following procedure: (i) fast cooling (linear line of pressure−temperature data with no hydrate); (ii) hydrate formation (identified by a sudden pressure drop and exothermic reaction); the amount of hydrate formation is controlled by increasing the temperature after about 3 MPa in pressure drop; (iii) fast heating at a rate of 5 K/hour; and (iv) slow stepwise heating so the hydrates dissociate and the system returns to its initial state. During step iv, as shown in the inset in Figure 4b, the temperature is increased stepwise by 0.2 K and held for 2 h so that the system pressure stabilizes at each temperature. It is particularly important to use this stepwise temperature increase because hydrates are dissociating and the system re-equilibrating, not only in terms of the amount of the phases present, but also in terms of the salt concentration in solution, which is gradually changing upon hydrate dissociation (hydrates consume the water). The inflection point, intersection of the hydrate dissociation line and the heating/ cooling line, corresponds to the hydrate equilibrium temperature and pressure for a given salt concentration. After reaching the inflection point (hydrate phase equilibrium point), the pressure increment at each step sharply decreases, since there is no more gas released due to hydrate dissociation, and the pressure is increased only due to the gas expansion with temperature, as shown in Figure 4b. For the measurement of hydrate formation kinetics, the salt solution is injected to the cell and the system pressurized up to 151.6 MPa with methane gas at 288.15 K. After reaching steady

2. EXPERIMENTAL SECTION Figure 3 shows the experimental setup for the measurement of hydrate phase equilibrium and formation kinetics in high salinity systems and at very high pressures. The operating pressure limit for the system is 207 MPa, and the system can be temperature controlled from 238.15 to 333.15 K. For safety considerations, a hazard and operability (HAZOP) study was done on the system based on the worst what-if scenario (see Supporting Information). Double-polycarbonate shields surround the entire system so that any accidental pressure release is contained. The system has two pressure relief valves set to 207 MPa. Three kinds of pressure indicators are installed to ensure proper pressure monitoring: analogue type (Astra, 0.5% accuracy), digital type (Honeywell, 0.2% accuracy), and pressure transducer (Omega, 0.05% accuracy). The pressure cell, internal volume of 56 mL (manufactured by SejinYoungTech in Korea), was designed to sustain pressures up to 240 MPa. Because of the corrosive nature of salt on stainless steel, a protective cylinder made of Inconel 600 is placed in the pressure cell (item L in Figure 3) so that the stainless steel sections of the pressure vessel are not wetted by the brine solution. A baffle (item M in Figure 3) and a resistance temperature detector (RTD, ± 0.25 K uncertainty), both also made of Inconel 600, are placed in the cell. For each experiment, 25 mL of sample solution (0, 12, 23, 26, 27, and 30 wt % NaCl, 99% purity, Macron Fine Chemicals) is injected in the cell. The cell is immersed into the coolant bath containing a propylene glycol−water mixture, C



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Figure 4. Experimental procedure for the measurements of hydrate phase equilibrium in a high salinity system via the isochoric method. (a) Pressure and temperature trace for hydrate formation and dissociation in the 23 wt % NaCl system. Inset shows the slope change of the heating curve as the phase equilibrium point is reached. (b) Pressure (solid red) and temperature (solid blue) traces. Temperature steps are maintained for about 2 h until the pressure stabilizes.

conditions, the cell was cooled to 278.75 K. Temperature (liquid phase) and pressure in the cell are monitored to observe the formation of hydrates.

On the basis of the phase diagram in Figure 2, measurements in Figure 5a were made for three-phase equilibrium (CH4(g)− NaCl solution(l)−hydrate(s)) with NaCl concentrations of 0, 12, 23, and 26 wt %, and four-phase equilibrium (CH4(g)− NaCl solution(l)−NaCl precipitated(s)−hydrate(s)) with 27 and 30 wt % NaCl. The phases present for 27 and 30 wt % NaCl were verified through visual observations performed at atmospheric pressure and various temperatures between 263.75 and 293.15 K, see Figure 5b. On increasing the NaCl concentration from 12 to 27 wt %, the hydrate stable region shifted to harsher condition, that is, lower temperature and higher pressure. Interestingly, the data for 27 and 30 wt % NaCl coincide, demonstrating that the measured equilibria at these concentrations correspond to the same equilibrium conditions corresponding to the quadruple point: the excess salt above the saturation concentration is precipitated and the solution remains at saturation. 3.2. Growth Kinetics of Methane Hydrates at Extreme Conditions. Measurements of hydrate growth rate as shown in Figure 6 were performed to assess the effect of high salt

3. RESULTS AND DISCUSSION 3.1. Thermodynamics of Methane Hydrates at Extreme Conditions. Figure 5a shows the methane hydrate phase equilibria measured up to 200 MPa with different concentrations of NaCl. The actual values of the data points are listed in the Supporting Information. The measurements for each concentration were repeated at least twice with a fresh solution to verify the data reproducibility (see Supporting Information). The formation of gas hydrate is a physicalchemical combination between gas molecules and water, excluding impurities, such as salt, in the solid structure. The amount of water converted to hydrates in the experiments was intentionally controlled by monitoring the pressure drop (about 3 MPa) in order to prevent significant changes in the solution’s NaCl concentration. D



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Figure 5. Methane hydrate phase equilibrium for high salinity systems. (a) Methane hydrate three-phase equilibria with 0,15 12, 23, and 26 wt % NaCl and four-phase equilibria with 27 and 30 wt % NaCl. Semifilled symbols are repeat experiments with fresh solution. The solid/dashed lines represent predictions by CSMGem, which is applicable under 68.9 MPa (solid and dashed are for pressures below and above 68.9 MPa, respectively). (b) Visual observations of NaCl precipitation at/near- or oversaturated concentration at 293.15, 263.75, and 278.75 K and atmospheric pressure. Images with a blue dashed outline represent clear solutions (no salt precipitate), while images with a red solid outline have salt precipitated (NaCl or NaCl·2H2O), which are analogous to the phase diagram in Figure 2.

precipitate at room temperature (∼295 K). In the case of the 27 wt % NaCl solution, the solution was supersaturated as all the salt added was fully dissolved at 353.15 K and subsequently cooled down to room temperature while mixing the solution in the cell, and as such, some amount of salt likely precipitated: NaCl saturation concentration is ∼26.4 wt % at 293.15 K. The solution with 26 wt % NaCl shows the fastest growth rate, which can be understood based on the driving force. On the basis of the phase equilibrium measured in Figure 5a, the subcooling for hydrate formation in the 26 wt % NaCl system is approximately 6.0 K, which is 1.6 K higher than that of 27 and 30 wt % NaCl system at 151.6 MPa. The initial gradient (∼3 h)

concentrations (26, 27, and 30 wt % NaCl) near or over the saturation concentration. All kinetic experiments started at 151.6 MPa and 288.15 K which are well outside the hydrate stable region. Subsequently, the system was cooled to 278.75 K to induce hydrate formation, which was detected from the continuous pressure decrease even after reaching the set temperature of 278.75 K. Each experiment was repeated at least two times (see Supporting Information). Figure 6a shows the growth rate measurements for solutions containing 26, 27, and 30 wt % NaCl. To start with 26 wt % solution (slightly undersaturated) had no solid salt, whereas the 30 wt % NaCl solution was saturated and the excess salt was present as salt E



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Figure 6. Measurements of growth rates for methane hydrates in 26, 27, and 30 wt % NaCl system. (a) Hydrate formation rates: (i) experiment started at the same temperature and pressure condition, (ii) cell temperature decreased, and (iii) constant temperature of 278.75 K maintained. t1 and t2 are the times when inflection occurred. (b) Cell temperature profile at inflection points. Temperature increase indicates a significant amount of hydrates formed at that time.

of the pressure drop is ∼5.2 MPa/hour, which is relatively steep, and then, gradually changes as the driving force eventually decreases to be the same for the 27 and 30 wt % NaCl systemdue to hydrate formation, the solution remains saturated at the saturation concentration of ∼26.4 wt % NaCl.13 While both the 27 and 30 wt % NaCl system had salt content above the saturation concentration, it is observed that the initial growth rate for the 30 wt % NaCl (∼0.2 MPa/hour) was lower than the 27 wt % NaCl (∼0.4 MPa/hour), but after about 5 h, a significant pressure drop happened, resulting in approximately double the amount of hydrate formed compared to the 27 wt % NaCl system, as shown in Table 1. One of the very interesting results from the hydrate formation data is the inflection points observed in the pressure traces, which are marked at time t1 and t2 in Figure 6a. These

inflection points are unusual and suspected to be related to competing solid precipitation between salt and hydrate. Table 1 also shows more detail on the estimated amount of NaCl precipitated (see Supporting Information for details on the calculated values) at around t1 and t2. Further evidence for increased hydrate formation at these inflection points is the cell temperature, shown in Figure 6b, indicating a slight spike in the temperature due to the exothermic nature for hydrate formation (ΔHformation = −54.2 kJ/mol10). In particular, for the 30 wt % NaCl case, the pressure drop and temperature increase at t1 in Figure 6b are very apparent. 3.3. Discussion on Gas Hydrate Formation. From the measurements of methane hydrate phase equilibrium in the NaCl system shown in Figure 5a, it is demonstrated that the hydrate phase equilibrium boundary moves to lower temperF



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the system tends to provide larger surface area, potentially increasing the growth rate and also the amount of hydrates. Ting and co-workers16 demonstrated that the induction time of salt crystallization is short when the amount of seed crystals is increased because of both the number of collisions and surface area are increased. To observe the hydrate formation process in these salt saturated systems, an experiment with 27 wt % NaCl was conducted in an autoclave cell with windows at approximately 10 MPa and 260.15 K with methane gas, as shown in Figure 8.

Table 1. Estimated Percentage of Salt Precipitated at Inflection Points in Pressure Trace during Kinetic Experiments and Final Conversion of Water to Hydrates total salt concentration initial

NaCl added water NaCl precipitation first inflection t1 beforea point total precipitation at t1 before t1 afterb total precipitation at t1 after second t2 beforea inflection Total precipitation at point t2 before t2 afterb total precipitation at t2 after final conversion of water to hydrates

26 wt % NaCl

27 wt % NaCl

30 wt % NaCl

7.83 g 22.30 g 0.00 g 2.3 h 0.02 g

8.24 g 22.28 g (∼0.28 g) 8.3 h 0.36 g

8.71 g 22.31 g 1.46 g 5.5 h 1.46 g

3.2 h 0.14 g

8.5 h 0.36 g

5.6 h 1.52 g 8.8 h 1.65 g 9.9 h 1.71 g

5.4%

2.2%

4.6%

a

t1 or t2 before stands for the time before starting the inflection. bt1 or t2 after means the time after finishing the inflection.

ature and higher pressure as the salt concentration increases up to the limit of salt saturation in solution. For systems with excess salt beyond the saturation concentration, the hydrate phase equilibrium will be unchanged from the conditions at saturation with four-phases (aqueous solution, hydrates, solid salt, and vapor) in equilibrium, as shown by the data measured for systems with 27 and 30 wt % NaCl. At the conditions with four phases present, there is competition for solid precipitationsalt or hydrate. The salt concentration in solution will always remain at the saturation concentration, as such, any amount of hydrate formed causes a depletion of free-water and thus precipitation of salt. On the other hand, hydrates can only form from unbound water not associated with the solvation of the dissolved salt, as shown in Figure 7. This competition is quite unusual and unique, as both involve the nucleation/growth of a solid phase from solution.

Figure 8. Visual observation of methane hydrate formation with scale deposition in 27 wt % NaCl system performed at 10 MPa and 260.15 K. Three separate phases are clearly observed due to the density difference between the phases: hydrate, NaCl solution, and solid NaCl.

The inflection points are also observed in this experiment (see details in the Supporting Information). Visual observations of the experiment confirmed that as more hydrates formed, more salt precipitated, implicitly indicating that the dissolved ions are excluded from hydrate structure during hydrate formation. Furthermore, phase separation of hydrate, NaCl solution, and NaCl precipitate was very clear due to the density difference of the phases (NaCl aqueous solution, 1.20 g/cm3;17 methane hydrate, 0.91 g/cm3;1 and solid NaCl, 2.16 g/cm3). Hydrates floated up to the gas−liquid interface and the salt precipitate settled down at the bottom of the liquid phase. The accumulation of hydrates at the gas−liquid interface can significantly slow hydrate growth as it forms a barrier in the contact of gas and liquid. Normally, hydrate formation depends on the formation conditions (e.g., subcooling temperature), state/history of the water, impurities, gas composition, mixing of the phases, and geometry of the system.18 The observed sudden pressure decrease, denoted as inflection points in Figure 6a, is suggested to correspond to the precipitation of salt crystals from supersaturated NaCl solution, consequently causing a sharp change in the hydrate formation rate. Kind et al.19 states that the supersaturation can be a driving force for crystallization from solution. Once the saturated solutions are cooled, the system can stay in a metastable state (even with mixing); consequently, the solution becomes supersaturated, potentially self-inhibiting hydrate formation due to strong electrostatic force between ions and water, causing the hydrate formation rate to slow down during this period. In addition, supersaturation of the salt solution may be further attained by the

Figure 7. Illustration of the competition between hydrate formation and salt precipitation.

Figure 6 shows the measurements for the hydrate growth rate in the 26, 27, and 30 wt % NaCl systems. From the comparison of the formation rates after the inflection points (0.2 MPa/hour and 0.7 MPa/hour for 27 and 30 wt % NaCl, respectively) and the amount of hydrates formed in 27 and 30 wt % NaCl systems, the larger amount of salts precipitated in G



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localized heat generated from the exothermic heat for hydrate formation (ΔHformation = −54.2 kJ/mol, CH41), which is relatively large compared to the heat of dissolution of NaCl in solution (ΔHprecipitation = −3.88 kJ/mol, NaCl20). The supersaturated solution will eventually spontaneously crystallize once the supersaturation limit is reached, and then, the hydrate formation rate may be accelerated from the increased solid content (larger surface area) and the availability of water as the solution returns to the saturation concentration, thus resulting in the inflection point in the pressure trace. Even though for the 26 wt % NaCl system there was no salt precipitate at the start of the experiment, the larger driving force induces faster hydrate formation, leading to the accumulation of heat in the cell, and consequently, the supersaturation of the solution. This interplay between salt and hydrates demonstrates the competition of solid precipitation in saturated salt solutions. Another consideration in the competing processes between hydrate formation and salt precipitation is the enthalpy change for the phase change. The heat of hydrate formation is ΔHformation = −54.2 kJ/mol gas = −9.0 kJ/mol H2O and the heat of salt precipitation is ΔHprecipitation = −3.88 kJ/mol NaCl = −0.71 kJ/mol H2O, assuming that the hydration number for methane hydrate and NaCl are 61 and 5.5,21 respectively. Comparing the enthalpy change on a per mol of water basis, the hydrate phase is much more stable than solid NaCl phase. Therefore, as long as there is a driving force for hydrate formation (e.g., subcooling), hydrates will tend to form more favorably than the salt precipitation, which will then only happen when the solution is supersaturated, leading to the observed inflection point in pressure trace.

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ASSOCIATED CONTENT

S Supporting Information *

. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b00292. HAZOP study for high pressure apparatus (Figures S1 and S2); reproducibility in measurements of hydrate phase equilibrium (Figure S3), methane hydrate phase equilibrium data measured (Table S1), calculations for hydrate formation rate/conversion ratio/salt precipitation (Tables S2 to S7); and kinetic experiments with 30 and 27 wt % NaCl systems (Figures S4 to S7) (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. ORCID

Amadeu K. Sum: 0000-0003-1903-4537 Funding

This work was funded and performed for DeepStar Phase XII 12202 Deepwater Technology Development Project by Colorado School of Mines. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors would like to acknowledge technical guidance and contributions of DeepStar Flow Assurance Committee chaired by Siva Subramanian, Jeff Creek, and Sally Thomas. A special recognition to DeepStar A Global Offshore R&D Consortium for the financial support to complete the work highlighted within this paper, in particular Joseph Gomes and Greg Kusinski from DeepStar for permission to publish the results of the study.

4. CONCLUSIONS Gas hydrate phase equilibrium and kinetics at high NaCl concentrations (near and at saturation in solution) and very high pressure (up to ∼200 MPa) were investigated to study the effect of total salt content on the phase equilibria and growth rate of hydrates. The measurements of hydrate phase equilibrium show the equilibrium boundary moves to lower temperature and higher pressure as the NaCl concentration increased up to the limit of NaCl saturation in the solution. In the systems with NaCl concentration above the saturation concentration, the hydrate phase equilibrium is unchanged from the conditions at saturation with four phases (liquid− hydrates−solid NaCl−gas) in equilibrium. In the kinetic experiments for hydrate formation with high salinity systems (above saturation concentration), inflection points were observed in the pressure trace as a factor of time, which demonstrates that (i) hydrates can still form even with salt precipitated and (ii) the formation of hydrates and salt precipitation are competing effects. On the basis of experimental data, a mechanism is proposed to describe the competing effects of hydrate formation and NaCl precipitation. Supersaturation of the solutions potentially self-inhibits hydrate formation due to strong electrostatic forces between ions and water, causing the hydrate formation rate to slow down. As the supersaturation limit is reached, the supersaturated solution eventually spontaneously crystallizes, resulting in an accelerated hydrate formation rate from the increased solid content (larger surface area) and the availability of water as the solution returns to the saturation concentration.

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REFERENCES

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Gas Hydrates Phase Equilibria and Formation from High

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