Effects of PVCap on Gas Hydrate Dissociation ... - ACS Publications

Jul 10, 2017 - 1500, 3000, and 6000 ppm (i.e., 0.075 to 0.6 wt %) of the KHI polyvinyl caprolactam (PVCap) present in the initial solution under isoch...
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Effects of PVCap on Gas Hydrate Dissociation Kinetics and the Thermodynamic Stability of the Hydrates Ann Cecilie Gulbrandsen*,†,# and Thor Martin Svartås† †

Department of Petroleum Engineering, Faculty of Science and Technology, University of Stavanger, 4036 Stavanger, Norway S Supporting Information *

ABSTRACT: Kinetic hydrate inhibitors (KHIs) are used in the oil and gas industry to prevent hydrate plug formation. Understanding the influence of KHIs on hydrate dissociation is important in order to evaluate the implication the additives have on gas hydrate remediation. The present study has examined decomposition of structure I, II, and H hydrates formed with 750, 1500, 3000, and 6000 ppm (i.e., 0.075 to 0.6 wt %) of the KHI polyvinyl caprolactam (PVCap) present in the initial solution under isochoric conditions. The hydrates were formed at various pressures ranging from 45 to 175 bara and dissociated by applying heating rates from 0.0125 to 0.2 °C/h. Elevated hydrate dissociation temperatures as compared to the uninhibited system were experienced during all experiments. The measured final dissociation temperature increased with increasing heating rate probably due to effects on dissociation kinetics, and the final dissociation temperature was significantly higher than the equilibrium temperature for the corresponding noninhibited system even at very low heating rates. In a previous study at pressure around 45 bara, we left the PVCap system 1.0 °C outside the hydrate region over a period of 13.7 days, and dissociation apparently decayed asymptotically toward a limit around 1.5 bar inside the hydrate region. This indicates that there exists a thermodynamic effect in addition to a kinetic effect for dissociation of hydrates formed with PVCap present. During dissociation, the amount of gas released as a function of increasing temperature was less for the PVCap systems as compared to the uninhibited reference system at similar pressure. This pointed in the direction of reduced fill fraction of gas in the PVCap treated system. Raman spectroscopy showed that structure I hydrates formed with PVCap had a lower large-to-small cage occupancy than hydrates formed without PVCap. Structure II hydrates formed with PVCap indicated that there was no methane in the large cages but only in the small cages. Hydrates formed without PVCap did contain methane in both small and large cages. These findings supported the observations from dissociation experiments that PVCap both increase the thermodynamic stability of the hydrates and affects the dissociation kinetics toward slower dissociation rates.

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INTRODUCTION Natural gas hydrates are non-stoichiometric, ice-like crystalline components formed from water and nonpolar small-sized hydrocarbon gas molecules such as methane, ethane, propane, and isobutane in addition to inorganic gases such as carbon dioxide, nitrogen, hydrogen sulfide, and others.1 Gas hydrates may form in oil and gas pipelines at temperatures well above the freezing point of water. The potential for gas hydrate formation in hydrocarbon transportation lines is a major issue which needs to be addressed when considering possible solutions for field production. The issue concerning hydrate formation is particularly important for situations like long tie backs, deep water, and shut-in situations in general, since the fluids then have time to cool down into the hydrate region. Hydrate inhibition and control is often the design basis for deepwater field developments.1 Even in shallow-water environments in the North Sea, it is common for hydrate management issues to dictate the system selection and topside design.2 Traditionally, the prevention of hydrate formation has been achieved with the addition of thermodynamic inhibitors, commonly methanol or glycols. However, in the last two decades, economic and environmental factors have motivated research and development to identify alternative low dosage (less than 1 wt %) hydrate inhibitors (LDHIs). Kinetic inhibitors do not lower the onset temperature of hydrate formation, but retard the process. Some LDHIs prolong the © 2017 American Chemical Society

induction time for hydrate nucleation and reduce growth (kinetic hydrate inhibitors, KHIs) while other LDHIs prevent hydrate crystal agglomeration (anti-agglomerants). The key ingredients in a KHI product are polymers or copolymers containing primarily vinyl lactam monomers, specifically the monomers vinylpyrrolidone and vinyl caprolactam.3 The kinetic inhibitor molecule retards the crystal growth by either adsorbing to the growth sites of the crystal or fitting into the crystal lattice.4 By interacting with the crystal in such a manner, the KHI distorts the crystal lattice or growth steps, thus preventing the crystal from growing rapidly in a regular crystal structure. The time required for hydrates to form depends on the effectiveness of the KHI, its dosage rate, and the driving force for hydrate formation.3 It has previously been observed that hydrates formed in the presence of polymeric kinetic inhibitors increased the dissociation temperature or the time required for complete hydrate decomposition.5−13 It has been proposed that the reason may be the fact that more hydrate formed in the presence of inhibitor, which then takes longer to decompose. On the other hand, increased hydrate stability in the presence of inhibitor, as reflected by complete decomposition at longer Received: December 27, 2016 Revised: July 6, 2017 Published: July 10, 2017 9863

DOI: 10.1021/acs.energyfuels.6b03478 Energy Fuels 2017, 31, 9863−9873

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Energy & Fuels

of the two peaks were dependent on the concentration of PVCap in the system. Multiple dissociation peaks were also observed by Daraboina et al.7 Multiple hydrate melting events, some indicating the formation of hydrate structures with high stability, were observed in the presence of the kinetic inhibitors PVP and H1W85281. Such multiple peaks suggest hydrate with different compositions, and thus nonuniform crystals are formed in the presence of PVP and H1W85281. Hydrates formed in the presence of any of the inhibitors started melting at a lower temperature than hydrate formed in water. The multiple hydrate melting peaks that were observed indicate hydrate structures of greater complexity compared to those formed in the presence of water. Increased hydrate stability, as reflected by decomposition at elevated temperatures, depends on hydrate composition. The presence of impurities can broaden the melting temperature range.23 It is not known if a hydrate of one composition can act as an “impurity” in the hydrate of another composition or if inhibitors can act as impurities. Studies have revealed that the crystalline morphology of the hydrate is substantially altered by the addition of kinetic inhibitors.24,25

time and increased temperatures, may depend on hydrate compositional and structural changes.14 Kvamme et al.15 correlated hydrate inhibitor interactions to binding strengths. Alternatively, it has been suggested that very effective KHIs, once close to the hydrate surface, have a high probability of being “anchored”, using a chemical group that fills an empty cage, not only resulting in the inhibition of further crystal growth but also partially explaining the reduction in gas uptake. Lee et al.16 have studied methane hydrate formation in the presence of PVCap via Raman spectroscopy and found that PVCap prevents the rate of large cage cavity encapsulation at an early stage of hydrate formation. It was found that methane hydrates with PVCap reached a smaller value of the theoretical water to hydrate conversion compared to hydrates formed from pure water. This was seen as an indication that PVCap prevents further gas hydrate growth. It is assumed that kinetic inhibitors adsorb on the surface of hydrate microcrystals and significantly alter surface tension at the interface between the hydrate forming phases.4,9,17 Larsen et al.6 proposed that PVCap adsorbs onto hydrate surfaces through hydrogen bonding. As a result of surface pinning by the adsorbent, crystal surfaces have microcurvatures, resulting in crystal-growth inhibition due to the Gibbs−Thomson (Kelvin) effect. The adsorption of PVCap to hydrate surfaces was observed using small-angle neutron scattering18 which was also supported by molecular simulations.4,19 Daraboina et al.14 formed hydrates from a synthetic natural gas mixture consisting of methane, ethane, and propane. The system contained two different KHIs; the commercial chemical inhibitors polyvinylpyrrolidone (PVP) and H1W85281 (a proprietary commercial product of unknown composition). Raman spectroscopy confirmed that these hydrates formed with KHI were heterogeneous in contrast to the seemingly homogeneous hydrates formed in water. The large hydrate cages showed a reduction in methane compared to corresponding system without kinetic inhibitor. The methane molecules in the large cages appeared to be substituted by ethane, likely as a result of the preferential enclathration of the heavier hydrocarbon. When this occurs, the methane concentration in the gas phase rises, leading to an increase in the equilibrium pressure, and that, in turn, reduces the driving force for hydrate formation. Evidence for the inhomogeneous nature of hydrate formed in the presence of PVP and H1W85281 has been previously observed in DSC7 and gas uptake experiments.20 In the presence of inhibitors, metastability occurs during hydrate formation and hydrate dissociation.10 The performance of PVP and PVCap as kinetic gas hydrate inhibitors in saline solutions and with heptane has been evaluated. Hydrate decomposition took longer and proceeded in two steps.21 Makogon et al.9 found a hydrate metastability zone for hydrate decomposition for hydrates formed with kinetic inhibitors. The metastable zone was determined by the inhibitor’s activity, the magnitude of the water vapor pressure decrease, and the lowering of the water chemical potential in the presence of inhibitor adsorbed on the hydrate crystal surface. Lachance et al.22 measured the effect of kinetic inhibition for hydrates in emulsions (crude oil). In these particular experiments, the dissociation trends showed that instead of just one dissociation peak for sI hydrate, there were two peaks in each experiment with the PVCap polymer. The same trend was observed for a simple water/PVCap system, indicating that the trend was not a crude oil effect. Amplitudes and temperatures

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EXPERIMENTAL PROCEDURE FOR P-T MEASUREMENTS

The experimental setup for P-T (pressure -temperature) measurements is shown in Figure 1 and incorporates an autoclave cell

Figure 1. Outline of the experimental setup. consisting of a titanium cylinder with top and bottom end piece. A stirrer blade is connected to a magnet house in the bottom end piece via an axle. An outer rotating magnetic field created by a laboratory stirrer bar drive was used to regulate the stirrer speed. The stirrer motor can be regulated to maintain speed in the range 0 to 1200 rpm. The free volume between the top and end pieces is 141.3 mL. The free volume (dead volume) around the stirrer magnet inside the magnet house is 8 mL. The cell is equipped with a cooling/heating jacket, and temperature control is obtained by circulating water via a JULABO F34 HL refrigerated circulator. The desired temperature profile of the experiments is set on the JULABO F34 HL temperature control unit and the temperature is regulated within a set point stability of ±0.02 °C. The cell systems were equipped with 4-wire lead 1/10 DIN Pt-100 temperature sensors (accuracy ±0.03 °C) and Rosemount 3051 TA absolute pressure transmitters. Pressure is measured via in the inlet 9864

DOI: 10.1021/acs.energyfuels.6b03478 Energy Fuels 2017, 31, 9863−9873

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Energy & Fuels tubing and the temperature is measured inside the cell (in the vapor phase). The temperature was measured to an accuracy of ±0.10 °C and pressure was measured to an accuracy of ±0.25 bar. Data were sampled using a LabView data acquisition program. The experimental progress was continuously monitored on the PC screen during the experiments. At the end of the experiment, data were transferred to office PC for analysis and graphical presentation. Structure I, sII, and sH hydrates were produced using four different hydrocarbon mixtures: 1. Structure I hydrates were formed from pure methane (Scientific grade 5.5, 99.9995 mol % purity). 2. During autoclave cell experiments structure II hydrates were formed from a binary mixture of 92.5 mol % methane and 7.5 mol % propane (SNG2). 3. During Raman spectroscopy sII hydrates were formed in external lab by ternary mixture composed of 90.28 mol % methane, 4.95 mol % ethane, and 4.77 mol % propane. 4. Structure H hydrates were formed from a mixture of scientific grade 5.5 methane and methyl cyclohexane (GC quality C7H14, purity ≥99%). The experiments were carried out using PVCap concentrations ranging from 750 ppm to 6000 ppm in solution. Three different PVCap molecular weight fractions (Mw = 2500, 6000, and 10 000 Da) were examined. Polyvinyl caprolactam (PVCap): pure, dry powder of the polymer. The PVCap batch used in the experiments was originally supplied as 50% solution in isopropanol and the isopropanol was removed trough a two-stage process with vacuum distillation at 20 °C followed by a final vacuum treatment at approximately 10−5 Pa overnight to produce dry PVCap powder. Autoclave Cell Experiments. The same procedure for preparation of the experiment and filling of the cell was followed in all experiments. A description of the general experimental procedure is described below: 1. The desired PVCap solution was prepared for the experiment. 2. The magnet house was filled with the aqueous solution and any air residue was squeezed out of the magnet section during mounting of the magnet house into the bottom end piece. Any residues of solution on the top surface were removed prior to mounting the bottom end piece into the cell cylinder. 3. 50 mL of the aqueous solution was filled into the cell, and the top end piece was mounted. 4. The temperature of the heating/cooling unit was adjusted to 20 °C prior to cell pressurization. 5. Prior to loading the cell to the experimental pressure it was purged twice with the natural gas mixture to be used by pressuring the cell to 60 bara. This was done to remove (dilute) any residues of air in the cell. 6. At approximately 20 °C the cell was loaded to the desired pressure. 7. The stirring rate was kept constant at 750 rpm during the experiments. Hydrates were generally formed at fixed temperature conditions by cooling the system down to the desired formation temperature. Hydrate formation was induced by magnetic stirring. The system was kept at the formation temperature for a period of time to produce the required amount of hydrates for the experiment. The hydrates were then dissociated by gradually increasing the cell temperature at preset heating rates. In the first stage, the system was in most experiments heated relatively fast to a temperature 4−5 °C below the estimated equilibrium dissociation temperature applying a heating rate (∂T/∂t) of 1.0 °C/h. At this point the heating rate was reduced to 0.2 °C/h or less, and the system was left at that rate until all hydrates were dissociated. In the vicinity of the final hydrate equilibrium point a low heating rate is required to provide uniform temperature equilibrium throughout the hydrate mass and the cell volume during the melting process. Heating rates as low as 0.0125 °C/h were applied in some of the experiments. The pressure and temperature conditions in the cell were frequently sampled during the experiment. Hydrate formation was indicated by a sudden pressure drop in the cell due to gas leaving the vapor phase and entering empty cages in the hydrate structure. This can be seen in Figure 2, which describes a typical PT plot of sampled data during a run. The experiment is initiated at the right end of the baseline, from

Figure 2. Pressure versus temperature plot obtained from experiment where hydrate formation (pressure drop) is followed by hydrate dissociation.

where the sample is cooled down. Hydrate formation is indicated by an increase in temperature and a decrease in pressure. This happens as gas molecules get enclathrated into the hydrate lattice and release of energy takes place. After the liquid has been transformed into hydrate, the temperature is increased, applying the fast heating rate up to the point indicated by the first step on the dissociation curve. From this point and upward the heating rate was reduced to 0.2 °C/h or less. The final hydrate dissociation point is where the dissociation curve intersects the cooling curve (baseline). The region to the right of this point represents PT conditions for system without hydrates present. The hydrate dissociation point is the hydrate equilibrium point, where the initial conditions are resumed after a hydrate formation/ decomposition cycle. Displacement of Dissociation Temperature. Hydrate dissociation temperatures were determined via PT plots as shown in Figure 2. The final hydrate dissociation point is found at the location where the PT dissociation curve (4−5 in Figure 2) intersects the “hydrate free” PT curve (1−2 in Figure 2). Measured hydrate dissociation temperatures were compared to calculated hydrate equilibrium temperatures, Teq Calc, for corresponding noninhibited systems using the Colorado School of Mines CMSHYD prediction program from the second edition of Sloan’s book on Natural Gas Hydrates.26 For our SNG2 methane−propane hydrate system, CSMHYD gave better agreement with experimental equilibrium data than the newer CSMGem version of this prediction program. Discrepancies (ΔT′) between experimental dissociation temperatures (Texpr) and predicted equilibrium temperatures (Teq Calc) at the experimental pressure were estimated as ΔT ′ = Texpr − Teq calc

(1)

Amount of Hydrates in System. The amount of hydrates in the isochoric system is estimated on the basis of the pressure drop due to hydrate formation. In a system without hydrates, the pressure is defined as a function of temperature via linear regression of PT points along the cooling line 1−2 in Figure 2. The pressure in the hydrate free system at any temperature T is then given by

PT,Reg = K1 + k 2· T

(2)

where subscript “T,Reg” means regression pressure, K1 is a regression constant for the line at given amount of gas in cell, and k2 is the slope of the regression line. Along the path 2−3 − 4−5 in Figure 2 the pressure drop due to hydrate formation at any temperature T is estimated by ΔPT = PT,Reg − PT,expr 9865

(3) DOI: 10.1021/acs.energyfuels.6b03478 Energy Fuels 2017, 31, 9863−9873

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Energy & Fuels where PT,expr is measured experimental pressure at temperature, Texpr, along the path. Finally, the amount of gas consumed (in moles) can be estimated by

Δn T =

ΔPT·Vg Z T· R· Texpr

(4)

where Vg is free gas volume in the cell (= Vcell − Vwater), ZT is the gas compressibility factor at the experimental temperature, R is the universal gas constant, and Texpr is absolute temperature in Kelvin. The fill fraction of gas in the hydrates can be estimated via CSMHYD. The amount of water converted into hydrates can now be estimated via the amount of gas consumed, the number of gas molecules per water molecule in lattice, and the initial amount of water loaded into the cell. For experiments at similar pressures, the amount of gas consumed is proportional with the pressure drop, ΔPT, and for simplicity, we based our evaluation of effects on pressure drop relations in the various pressure regions examined. Effects on Dissociation Kinetics. Hydrates were dissociated applying heating rates, ∂T/∂t, in the region 0.0125−0.2 °C/h. The rate is kept constant during the experiment; thus, (∂T/∂t)x = kx, where kx is heating rate of experimental series x. During isochoric dissociation, the rate can be described by the change in pressure as a function of time: ∂P ∂P ∂T ∂P = · = · kx ∂T ∂t ∂T ∂t

Figure 3. Slope of the PT curve during hydrate dissociation at heating rate ∂T/∂t = 0.2 °C/h. 2b and 2c between the two nonlinear parts 2a−2b and 2c−2d. The nonlinear parts of the dissociation curves with KHI are due to kinetic effects just past the start of slow heating and then just prior to complete dissociation of hydrates formed in the KHI system. The linear PT dissociation curve in the midsection is defined by the slope (∂P/∂T)2, while the nonlinear section is defined by (∂P/∂T)3 which is a function of T (cf., eq 7). As the temperature approaches the final dissociation point, (∂P/∂T)3 approaches a limit (∂P/∂T)lim. Comparison of slopes (∂P/∂T)1, (∂P/∂T)2, and (∂P/∂T)3 or (∂P/∂T)lim can then be used for the evaluation of the effects of KHI on the dissociation kinetics.

(5)

where ∂P/∂T is the slope of the PT curve during dissociation. If dissociation proceeds at equilibrium in the system, this curve describes the equilibrium between free water, hydrate, and the hydrate forming fluid at a given PT condition along this curve. Since hydrates are dissociated and the amount of hydrates decreases toward zero concentration at the final dissociation point, the ∂P/∂T curve describes equilibrium conditions inside the hydrate region as a function of the hydrate amount in the system. CSM Gem has an option for this type calculation, but unfortunately our version of this program fails due to convergence problems in the flash calculation. We have applied the slope of the ∂P/∂T curve over a section close to the final dissociation point to evaluate the effect of KHI on the kinetics of dissociation. For systems with pure water the ∂P/∂T curves can be approximated as a linear relation over a limited pressure region of approximately 5 to 10 bar around the final dissociation point. For systems with PVCap the ∂P/∂T curve was shown to be nonlinear over the last 1.5 to 1 bar prior to the final dissociation point. For analysis of this nonlinear section we have fitted the experimental data points along the ∂P/∂T curve to a polynomial relation

∂P = m0 + m1· T + m2 · T 2 + ··· + mn ·T n ∂T

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All Raman spectroscopic investigations were performed at the Helmholtz Centre Potsdam German Research Centre for Geosciences GFZ in Potsdam, Germany. The main component of the experimental setup was an optical pressure cell, which can be used in a temperature range between −27 and 80 °C and pressure range between 1 and 100 bar. The temperature of the sample cell was controlled by a thermostat and the cell temperature was determined with pt100 temperature sensor with a precision of ±0.1 °C. A pressure controller regulated the sample pressure with a precision of 2% rel. The system pressure was measured with a P3MB from Hottinger Baldwin Messtechnik with a precision of 0.01% rel. The small sample volume (0.393 cm3) and the cooling system of the cell minimized temperature gradients within the sample. The experiments were run under isobaric conditions. One important feature of the sample cell was the application of gas flow. The gas flow was measured and regulated with a commercial flowmeter F230-FA-11-Z from Bronkhorst. The gas flowed at 1 mL/min; therefore, it took 17 s for the incoming gas to pass the cell body and enter the sample camber of the cell: this time was sufficient to allow the gas to attain the cell temperature. A quartz window enabled the phases to be analyzed by Raman spectroscopy as well as visual observation and the recording of morphology changes during the hydrate formation and decomposition processes. A confocal Raman spectrometer (LABRAM, HORIBA JOBINYVON) was used, which allowed the laser beam to be focused on a precise point, e.g., the surface of a hydrate crystal, thus assuring that only the selected phase was analyzed. The composition and structure of these selected areas within hydrate phase can be determined with this method. All experiments were performed using a mixture of deionized water and 0.3 wt % PVCap 2.5k. 150 μL of the mixture was put into a pressure cell. Formation of structure I hydrate was obtained using pure methane (purity 99.995%) as hydrate forming gas. Formation of structure II hydrate was obtained using a mixture of methane (90.28

(6)

where m0, m1, etc., are regression fit constants. The slope of the ∂P/∂T curve is obtained by derivation of eq 6

( ∂∂TP ) = m

∂

∂T

1

+ 2· m2 · T + ··· + n·mn ·T n − 1

EXPERIMENTAL PROCEDURE FOR RAMAN MEASUREMENTS

(7)

We have analyzed the nonlinear part of the ∂P/∂T curves around the final dissociation point using this relation and searched for a minimum representing the final hydrate dissociation temperature. Figure 3 illustrates slopes of the hydrate dissociation curves in systems with and without KHI present in solution. The hydrates were dissociated at a fast heating rate (1 °C/h) up to points 1a and 2a followed by 1 h dwell time to equilibrate. The slow melting sequence was initiated by applying the desired and set heating gradient. Hydrates formed from pure water show one single and fairly linear dissociation rate over the whole slow heating region (1a to 1b in Figure 3) until hydrates are completely dissociated and the slope of the curve, (∂P/∂T)1, can be estimated by linear regression. The system with KHI may show PT curves with varying slopes over the dissociation regions, but linearization may be approximated over the midsection (2b to 2c in Figure 3) prior to complete dissociation (point 2d in Figure 3). In Figure 3 we have linearized the midsection between 9866

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Energy & Fuels Table 1. Summary Table for Performed Hydrate Dissociation Experiments

a

pressure [bara]

heating rate ∂T/∂t [°C/h]

displacement dissoc. temp ΔT′ [°C]

2500

90

3000

6000

45

sII

3000

6000

60

sII

3000

6000

90

sII

3000

10 000

90

sII

3000

6000

120

sII

3000

6000

175

sH

3000

6000

90

sH

3000

10 000

90

0.2a 0.025a 0.0125a 0.2 0.05 0.2 0.05 0.2 0.05 0.015 0.2 0.05 0.2 0.05 0.2 0.05 0.2 0.1 0.2 0.05

3.5 2.3 2.1 5.1 4.1 4.7 4.1 4.6 3.7 3.1 4.7 3.6 4.2 3.5 4.4 3.5 5.2 5.2 5.2 4.0

hydrate structure

PVCap concentration [ppm]

sI

3000

sII

molecular weight, Mw PVCap [Dalton]

First stage of the heating process went above T(CSMHYD).

mol %), ethane (4.95 mol %), and propane (4.77 mol %). The cells were pressurized to 40 bar at room temperature for both systems (sI and sII) and cooling of the cells was initiated until a temperature of −17.5 °C was reached. At this stage spontaneous formation of a solid phase occurred. Raman spectra were collected at this temperature as well as throughout the heating process, until dissociation occurred. For comparison, the same procedure was followed for the corresponding noninhibited systems. Additionally, microscopic observations of hydrate decomposition for sI and sII hydrates were obtained.

was significantly higher than the equilibrium temperature for the corresponding noninhibited system even at very low dissociation rates. For structure II hydrates, displacement of hydrate dissociation temperatures, ΔT′, was observed in the range 3−5 °C dependent on heating rate, system pressure, and molecular weight of the PVCap molecule. The lowest dissociation temperatures were obtained at the lowest dissociation rate, confirming kinetic effects. The observed effect on dissociation temperature appeared to be a function of inhibitor concentration at concentrations below 1500 ppm, but in the region 1500 to 6000 ppm the effect of concentration on dissociation temperature seemed marginal (cf., Table 2). For sI and sII hydrates formed without PVCap, all measured hydrate dissociation temperatures were in good agreement with predicted (CSMHYD) values provided that the applied heating rate was equal to or less than 0.2 °C/h. The ∂P/∂T relation followed a linear trend around the final dissociation point for both sI and sII systems without inhibitor present. This was not completely so for the pure sH hydrate, which showed a slightly nonlinear ∂P/∂T relation in the vicinity of the final dissociation point. Figure 4 shows the ∂P/∂T dissociation paths for the applied sII and sH systems without inhibitor present. The sII and sH hydrates in this figure were dissociated at heating rates of 0.2, 0.1, and 0.05 °C/h. The discrepancy between parallels measured at dissociation rates of 0.2 °C/h or less were negligible and normally within 0.1 °C. The discrepancy between experimental dissociation points and predicted hydrate equilibrium temperatures by CSMHYD were around ±0.1 °C for the sII system at pressures in the region 45−90 bara. At pressures above 90 bara the discrepancy between experimental dissociation points and those predicted by CSMHYD increased with increasing pressure up to around ±0.25 °C at 175 bara for sII. The discrepancy between experimental dissociation temperatures for the sH system at 90 bara and predicted hydrate equilibrium temperatures by CSMHYD was within 0.1 °C, at heating rates of 0.2 and 0.1 °C/h. The discrepancy between experimental parallel sH hydrate equilibrium temper-

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RESULTS Displacement of Dissociation Temperature. Elevated hydrate dissociation temperatures as compared to the noninhibited systems were experienced during all experiments on structure I, II, and H hydrates and the results are summarized in Tables 1 and 2. The measured final dissociation temperature increased with increasing heating rate probably due to effects on dissociation kinetics, but the final dissociation temperature Table 2. Effect of PVCap Concentration and Heating Rate on Displacement of Hydrate Dissociation Temperature PVCap concentration [ppm]

pressure [bara]

formation temperature [°C]

heating rate ∂T/∂t [°C/h]

displacement dissoc. temp ΔT′ [°C]

750 1500 3000 750 1500 3000 1500 3000 6000 1500 3000 6000

45 45 45 45 45 45 90 90 90 90 90 90

4 4 4 4 4 4 1 1 1 1 1 1

0.05 0.05 0.05 0.2 0.2 0.2 0.05 0.05 0.05 0.2 0.2 0.2

3.5 4.1 4.1 3.9 5 5 3.1 3.1 3.4 4 4 4.4 9867

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Figure 6. PT dissociation curves for SNG2 system with and without 3000 ppm PVCap-6k at 90 bara and heating rate of 0.05 °C/h.

Figure 4. PT dissociation curves for the pure SNG2 sII hydrate and methyl cyclohexane−methane sH system without inhibitor at 90 bara applying heating rates of 0.2, 0.1, and 0.05 °C/h.

atures measured at 90 bara was around 0.05 to 0.06 °C/h, and the deviation from the value predicted by CSMHYD was within 0.1 °C. SNG2 sII methane−propane hydrates were produced at pressures in the region between 45 and 175 bara and dissociated using heating rates of 0.2 and 0.05 °C/h. Methyl cyclohexane−methane sH hydrates were produced at pressures around 90 bara and dissociated at heating rates of 0.2, 0.1, 0.05, and 0.025 °C/h. The differences between the experimental and predicted dissociation temperatures (ΔT′) are listed in Table 1. The final dissociation temperatures in both sII and sH hydrates were significantly displaced toward higher temperatures, indicating increased stability of the hydrates. Furthermore, results for sII and sH hydrates at 90 bara indicated that the PVCap polymer length did not significantly influence ΔT′. Figures 5 and 6 show dissociation of SNG2 hydrates formed at

Figure 7. PT dissociation curves for methyl cyclohexane sH system with and without 3000 ppm PVCap-6k at around 90 bara and heating rate of 0.2 °C/h.

systems containing PVCap, indicating increased stability of the hydrates. To check whether this was a kinetic effect or an effect on equilibrium property of system or both, we examined dissociation behavior for the SNG2 structure II forming system at two different temperatures outside the hydrate region of the noninhibited system. At a temperature of 3.9 °C outside the hydrate region of the noninhibited system at 45 bara (T = 21.6 °C) it took approximately 3.5 to 4 days to obtain complete dissociation of hydrates formed in the presence of 1500 and 3000 ppm PVCap6k. At a temperature of 1.0 °C outside the hydrate region (18.2 °C at 45 bara) complete dissociation was obtained neither in the 1500 ppm system nor the 3000 ppm system when the experiments were interrupted by power failure 19.8 and 13.7 days after start of experiments, respectively. The latter two dissociation experiments are shown in Figure 8a,b, respectively. During the experiment with 1500 ppm PVCap the final amount of gas in the hydrate (ΔPT) apparently approached a limit corresponding to a pressure drop of 1.44 ± 0.02 bar (Figure 9). This showed that the displacement of the final dissociation point was not only due to kinetic effects but also most probably due to effects on the thermodynamic stability of the hydrates. In addition, extrapolation of the ΔPT curve at 3000 ppm PVCap apparently approaches a constant level around 1.6 bar as a function of increasing time. In Figures 5 and 6, it was observed that the PT dissociation curves for the system with PVCap showed a region with a linear ∂P/∂T relation (cf., the midsection 2b−2c in Figure 3) prior to the

Figure 5. PT dissociation curves for SNG2 system with and without 3000 ppm PVCap-6k at 45 bara and heating rate of 0.05 °C/h.

45 and 90 bara, respectively, with and without PVCap present in the solution. Figure 7 shows dissociation of sH hydrates formed at pressure around 90 bara with and without PVCap present in the solution. For the noninhibited sI and sII systems examined, the ∂P/∂T curves followed a linear path up to the final dissociation point, while the sH system showed nonlinear behavior over a region of 0.2 to 0.3 °C prior to the final dissociation point. The final dissociation temperatures were significantly displaced toward higher temperatures in all 9868

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Figure 8. Gas release curves at pressure of approximately 45 bara for systems with 1500 ppm (a) and 3000 ppm (b) PVCap-6k at constant temperature of 18 °C.

Gulbrandsen and Svartaas27 observed a relation between the start of the nonlinear dissociation path and a surface adsorption factor, SPVCap, defining the amount of PVCap adsorbed/bound to the hydrate. Figures 5 and 6 also show that the slopes of the PT dissociation curves (∂P/∂T) were different in systems without and with PVCap. For the linear part of the PVCap systems, the slope (∂P/∂T)2 was reduced as compared to that of the pure system, (∂P/∂T)1. Similar trends with reduced slopes (∂P/∂T)2 in the linear section were observed at all pressures examined for the PVCap systems. It was observed that a reduction of the heating gradient resulted in a displacement of the ∂P/∂T curve toward lower temperatures, while the slope remained fairly constant (see Supporting Information sections S1 to S3). The constant slope most probably reflects a correlation with equilibrium conditions for the PVCap containing system inside the hydrate region, while the displacement most probably is due to kinetic effect (e.g., dissociation rate). This means that PVCap affects not only the dissociation temperature, but also the gas release during dissociation. This could be due to some interaction with the PVCap molecule reducing the fill fraction of gas in the hydrate lattice and thus affecting the equilibrium properties of the system. This assumption is supported by results from Raman spectroscopy on the systems described in the section below. Daraboina et al.14 observed some changes in cage occupancy of methane, ethane, and propane for sII hydrate with PVP and H1W85281, but their Raman, NMD, and PXRD measurements were conducted well inside the hydrate region and not during dissociation and closer to equilibrium conditions. In theory, cage occupancy is analogous to surface adsorption via Langmuir constants. Thus, it could be reasonable to assume that substances adsorbed on the hydrate surface (e.g., PVCap) could interact with the cage occupancy and hydrate stability. Additional information on analysis of ∂P/∂T dissociation paths in systems with PVCap is given as Supporting Information (sections S1 to S3). Raman Measurements. Structure I System. The symmetric C−H stretching vibration of methane encased in the structure I methane hydrate split into two peaks. These indicate the location of the methane molecules in the hydrate, namely, in the pentagonal dodecahedral cavities (512) and the tetrakaidecahedral cavities (51262). Structure I is constituted by two different cages and the unit cell consists of six large

Figure 9. Gas release curves at end of dissociation experiment with 1500 PVCap-6k at constant temperature of 18 °C and approximately 45 bara pressure.

nonlinear section around the final dissociation point. The nonlinear ∂P/∂T relation started approximately 1 to 1.5 bar below the final dissociation point. Analysis of the slope of the ∂P/∂T curve in this nonlinear region indicated that the slope approached the slope of the (∂P/∂T) baseline for system without hydrate at the final dissociation point. Such dissociation behavior was not observed for sI and sII hydrates with pure water, but the sH system with pure water gave a nonlinear ∂P/ ∂T curve over the last 0.2 to 0.3 °C prior to the final dissociation point. Lachance et al.22 observed two dissociation peaks for the sI hydrate system with PVCap, and Daraboina et al.7,20 reported similar observations for PVP and commercial inhibitor H1W85281 on sII hydrates. The two ∂P/∂T relations for PVCap hydrates in Figures 5 and 6 could be related to similar phenomena. For the sH system without inhibitor the nonlinear ∂P/∂T path in the vicinity of the final dissociation point could be related to a nongaseous hydrate former (e.g., liquid methyl cyclohexane) that will not contribute to a similar pressure response during hydrate dissociation as released gaseous hydrate formers would do. The nonlinear ∂P/∂T path for the inhibited sH system could be due to both nongaseous hydrate former and the PVCap being released from the hydrate (cavity and/or surface) during final dissociation. 9869

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at 40 bar occurred at 18.4 °C. This corresponds to a displacement of the dissociation temperature of 3.6 °C outside the hydrate region at 40 bar. During the Raman experiments the noninhibited hydrates dissociated at a temperature in agreement with predicted equilibrium temperature by CSMHYD as observed during cell experiments at all pressures.

cavities (51262) and two small cavities (512). Hence, the ratio of large to small cavities theoretically should be 3.0.26 Raman spectra were collected from both inhibited and noninhibited hydrates. Spectra revealed a ratio between methane in the large and small cages which did not comply with a normal sI spectrum. With time one could observe a transformation into “normal” structure I hydrate for the systems. This indicates that in the initial stage of hydrate formation, small cages were excessively formed, and in turn, the amount of large cages was not sufficient to form perfect unit cells of sI hydrate. A similar observation was reported by Lee et al.16 for a stirred system with PVCap and Subramanian et al.,28 who suggested that the formation of the large 51262 cavity is the rate-limiting step for hydrate formation. For sI hydrates the large to small cage occupancy ratios were biased from the theoretical values at an early stage of hydrate formation. There were indications that this effect was somewhat stronger for hydrates formed with PVCap versus hydrates formed without PVCap. It has been reported with real time Raman spectroscopy that there is a difference regarding development of cavity ratio with time for sI hydrates formed with and without PVCap.28 The large to small cage occupancy ratio for hydrates formed with PVCap eventually resulted in a value of 2.5. This is in good agreement with what others have observed for methane hydrates formed with PVCap.28 During dissociation of the hydrates formed with PVCap, at a temperature of 0.3 °C, a morphologic alteration from wet to powdery was observed. Pre and post the morphologic alteration, the Raman spectra showed the same gas composition, but at the same time, the spectra showed higher water to gas ratio after the alteration had taken place, indicating reduced total fill fraction in the cavities. Hydrates without inhibitor dissociated at 5 °C and the PVCap containing hydrates dissociated at 6 °C at 40 bar. Structure II System. Raman spectra were collected and some spectra revealed normal sII cage occupancy, whereas other spectra revealed a ratio between methane, ethane, and propane that did not comply with normal sII cage occupancy. With time, one could observe a transformation into normal structure II for the hydrates formed without PVCap. The cage occupancy for hydrates in the system containing PVCap differed from that without PVCap. Changes in cage occupancy were observed by changes in peak heights (signal intensity) at the wavenumber corresponding to the respective cage. Raman spectra from these measurements are given in Supporting Information (section S4, Figures S4.1 and S4.2). The hydrates which had been formed without PVCap showed signals for methane, ethane, and propane in the large cavities of structure II. In the hydrates that had been formed with PVCap, almost no methane was detectable in the large cavities, only ethane and propane. As the system temperature was increased, a transformation of the hydrates in the PVCap containing system could be observed at 8 °C. The transformation did not take place in hydrates formed without PVCap, but a similar transformation was observed for the PVCap containing methane sI hydrate. First, the hydrate appearance was altered from solid to wet, followed by a second step where the appearance changed from wet to a powdery, compact state. Spectral analysis revealed that no change in guest cage occupancy had occurred during the visual hydrate transformation. At the same time results revealed that an alteration regarding the water to gas ratio had taken place. Hydrate dissociation for the PVCap containing hydrates

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DISCUSSION Gas hydrate decomposition has been investigated in the presence of the kinetic inhibitor PVCap. We have quantitatively described the influence of the melting rate on the final dissociation temperature for structure I, II, and H hydrates formed with PVCap. Results indicated the melting rate to be a parameter influential for the dissociation temperature. Even for very slow melting rates such as 0.0125 °C/h, the final dissociation temperature was significantly higher than the equilibrium temperature for the corresponding noninhibited system. This indicates that there exists a thermodynamic effect in addition to a kinetic effect for hydrates formed with PVCap. Experiments regarding the influence of PVCap concentration on dissociation temperature for sII hydrates showed that low inhibitor concentrations (750 ppm) resulted in lower dissociation temperatures than intermediate (1500 and 3000 ppm) concentrations. 1500 and 3000 ppm resulted in exactly the same dissociation temperatures. Higher concentrations of 6000 ppm had a further increase in dissociation temperature. Results indicate the existence of a threshold concentration level of the inhibitor, above which the hydrate stabilizing effect of the PVCap was not further influenced until another threshold level is encountered. Previous studies have reported on the influence of VC713 concentrations on the temperature of dissociation for THI hydrates.29 Concentrations of 5000 ppm resulted in higher dissociation temperatures than at 250 ppm. Increasing the inhibitor concentration to 10 000 ppm did not elevate the temperature of dissociation any further. Sharifi et al.30 examined propane hydrates in their experiments and found that increasing the concentration of PVCap increased the amount of hydrate which dissociated at higher temperatures and decreased the amount of hydrates that dissociated at the equilibrium temperature. For the sII hydrate it was found that for PVCap concentration of 0.02 wt % the dissociation temperature nearly coincided with the equilibrium temperature. Experimental results indicated that hydrates formed in the presence of PVCap obtained increased stability. Observations revealed morphologic alterations taking place as dissociation proceeded in the hydrates formed with PVCap versus hydrates formed without PVCap. An increased thermodynamic barrier or a reduced kinetic transfer rate may be attributed to a difference in microstructure. Studies31 have been performed where THF hydrate was produced with PVCap amounts in the range 0.05−0.25 wt %. Results from these experiments revealed striking morphological changes when PVCap was added to the system. Observations of dendritic growth were obtained, indicating that heat transfer limitations were enhanced by the presence of the inhibitor. It was not found whether inhibition and adsorption were specific to the {110} faces, but it appeared that there was adsorption on {110}, probably interfering with step growth across these faces and leading to the development of the branching structure.18 Starting from an ice phase, Małolepsza and Keyes32 show by molecular dynamics simulations using a replica exchange algorithm that the formation of methane hydrate is nucleated by the incorporation of methane molecules into the ice phase which catalyzes the 9870

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Energy & Fuels transformation of ice to a metastable β-phase which then transforms to a methane hydrate. This work describes a route to hydrate formation at low temperatures. Experimentally, hydrate formation from ice is a two-stage process with a relatively fast conversion of the ice surface33−36 and a much slower process for bulk ice. The structure I hydrates without PVCap turned into normal large to small cage occupancy (3.0) as time progressed and the hydrates with PVCap had a somewhat lower value (2.5). Regarding structure II hydrates, it turned out that the composition for hydrates with PVCap differed from hydrates without PVCap. Hydrates without PVCap showed signals for methane, ethane, and propane in the large cavities (51264) of structure II. In hydrates with PVCap only ethane and propane were detectable in the large cavities, and almost no methane was detectable in the large cavities. Earlier, it was reported that the cage occupancy of methane in large cages of sII hydrates decreases in the presence of PVP and PVCap.37 In hydrates with PVCap, it seems that methane prefers the small cages and only propane and ethane stabilize the large cages. This might be an explanation why this hydrate is more stable to higher temperatures compared to those hydrates which have been synthesized without PVCap, where methane occupies large and small cavities. Uchida et al.38 reported that mixed gases can form sI and sII in multiple steps, with heavier components favored for enclathration, resulting in an enriched light gas phase, or a fractionation effect. In natural gas mixtures with three components, methane, ethane, and propane, both sI and sII hydrates can form simultaneously depending on the gas composition.39 Daraboina et al.14 reported that heavier hydrocarbons tend to participate more in hydrate formation in the presence of inhibitors causing fractionation of the gas. They also observed that hydrate formed in the presence of inhibitor had different local compositions and structures within the same sample. These findings coincide with what we observed for our systems. Ripmeester et al.40 analyzed the viscoelastic properties of the adsorbed PVP and PVCap molecules revealed by QCM-D. The ratio of the change of dissipation factor and adsorption mass, R, calculated as R = ΔD/Δm, is a measure of the adlayer status. A large absolute R value indicates a porous, flexible adlayer with considerable trapped liquid.41,42 PVCap and PVP showed two distinct steps of adsorption, indicating a rearrangement of the adsorbed layer as adsorption progressed, and R2 represents the final status of the adlayer. Our microscopic investigations revealed that for PVCap hydrates there was release of countless gas bubbles at a certain temperature well into the hydrate stability region. This only occurred for the PVCap containing hydrates and not for the corresponding hydrates formed with gas and water. After the incident had taken place the appearance of the hydrates had altered. The hydrate crystal altered its appearance, indicating changes in thickness gradient across the surface. Raman spectra were recorded pre and post the event taking place. The obtained spectra showed the same composition, indicating that the apparent morphological change could not be attributed to a change in composition. Previously, reports have stated that the presence of kinetic inhibitors can lead to increased order at the surface, and consequently the water molecules will be more efficiently packed.43 Lee and Englezos suggest that a proper term for the hydrate formed in the presence of kinetic inhibitors would be “hard” versus “soft” plugs formed in noninhibited systems.8 It

should be noted that the change in appearance in our experimental system did not occur during hydrate formation, but during the heating of the sample. Raman spectra collected before and after the transformation showed no difference regarding the guest molecule occupancy, but displayed higher water to gas ratio after the alteration had taken place. Morphological alterations may influence the stability of the different cages that constitute the various hydrate structures. The hydrates with PVCap may become stabilized to a higher degree by alterations in the binding strengths. This can alter the total energy of the cavity, and as a result, it may be able to withstand a higher temperature before dissociation takes place.

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CONCLUSION Recently, the application of kinetic hydrate inhibitors in the oil and gas industry has become increasingly important because of their strong inhibitory effect in relatively small doses. In this study, gas hydrate decomposition has been investigated in the presence of the kinetic inhibitor PVCap. We have quantitatively described the influence of the melting rate on the final dissociation temperature for structure I, II, and H hydrates formed with PVCap. Results indicated the melting rate to be a parameter influential for the dissociation temperature. Even for very slow melting rates such as 0.0125 °C/h, the final dissociation temperature was significantly higher than the equilibrium temperature for the corresponding noninhibited system. This indicates that there exists a thermodynamic effect in addition to a kinetic effect for hydrates formed with PVCap. Furthermore, experiments regarding the influence of PVCap concentration on dissociation temperature for sII hydrates showed that low inhibitor concentrations (750 ppm) resulted in lower dissociation temperatures than intermediate (1500 and 3000 ppm) concentrations. These resulted in exactly the same dissociation temperatures. Higher concentrations of 6000 ppm had a further increase in dissociation temperature. Results indicate the existence of a threshold concentration level of the inhibitor, above which the hydrate stabilizing effect of the PVCap is not further influenced until another threshold level is encountered. For sI hydrates the large to small cage occupancy ratio were biased from the theoretical values at an early stage of hydrate formation. Results revealed some differences for PVCap systems and non-inhibitor-containing systems. The large to small cage occupancy ratio for sI hydrates formed without PVCap and with PVCap eventually resulted in a value of 3.0 and 2.5, respectively. This indicates that for hydrates formed with PVCap, it is not only the kinetic effect for how fast the large cages are formed that is influenced, but also the final product which again can have an impact on hydrate stability. The composition for sII hydrates with PVCap differed from the composition for hydrates without PVCap. Hydrates without PVCap showed signals for methane, ethane, and propane in the large cavities (51264) of structure II. In hydrates with PVCap, only ethane and propane were detectable in the large cavities, and almost no methane was detectable in the large cavities. In hydrates with PVCap, methane prefers the small cages and only propane and ethane stabilize the large cages. This might be an explanation why hydrates with PVCap are stable to higher temperatures compared to hydrates which have been synthesized without PVCap, where methane occupies large and small cavities. Microscopic investigations revealed that for PVCap hydrates there was a morphological transformation taking place, well 9871

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(6) Larsen, R.; Knight, C. A.; Sloan, E. D. Clathrate hydrate growth and inhibition. Fluid Phase Equilib. 1998, 150, 353−360. (7) Daraboina, N.; Ripmeester, J.; Walker, V. K.; Englezos, P. Natural gas hydrate formation and decomposition in the presence of kinetic inhibitors. 1. High pressure calorimetry. Energy Fuels 2011, 25, 4392− 4397. (8) Lee, J. D.; Englezos, P. Unusual kinetic inhibitor effects on gas hydrate formation. Chem. Eng. Sci. 2006, 61, 1368−76. (9) Makogon, Y. E.; Holditch, S. A. Gas hydrates-conclusion: Experiments illustrate hydrate morphology, kinetics. Oil & Gas Journal 2001, 99, 45−50. (10) Makogon, Y. F.; Holditch, S. A. Lab work clarifies gas hydrate formation, dissociation. Oil & Gas Journal 2001, 99, 47−48 50−52. (11) Svartaas, T. M.; Gulbrandsen, A. C.; Husebø, S. B. R.; Sandved, O. An experimental study on “un-normal” dissociation properties of structure II hydrates formed in presence of PVCAP at pressures in the region 30 to 175 bars − dissociation by temperature increase. Proceedings of the 6th International Conference on Gas Hydrates; Vancouver, Canada, 2008; p 5696. (12) Kelland, M. A. History of the development of low dosage hydrate inhibitors. Energy Fuels 2006, 20, 825−847. (13) Lederhos, J. P.; Long, J. P.; Sum, A.; Christiansen, R. L.; Sloan, E. D., Jr. Effective Kinetic Inhibitors for Natural Gas Hydrates. Chem. Eng. Sci. 1996, 51, 1221−1229. (14) Daraboina, N.; Ripmeester, J.; Walker, V. K.; Englezos, P. Natural gas hydrate formation and decomposition in the presence of kinetic inhibitors. 3. Structural and compositional changes. Energy Fuels 2011, 25, 4398−4404. (15) Kvamme, B.; Kuznetsova, T.; Aasoldsen, K. Molecular dynamics simulations for selection of kinetic hydrate inhibitors. J. Mol. Graphics Modell. 2005, 23, 524−536. (16) Hong, S. Y.; Lim, J. I.; Kim, J. H.; Lee, J. D. Kinetic studies on methane hydrate formation in the presence of kinetic inhibitor via in situ Raman spectroscopy. Energy Fuels 2012, 26, 7045−7050. (17) Kumar, R.; Lee, J. D.; Song, M.; Englezos, P. Kinetic inhibitor effects on methane/propane clathrate hydrate-crystal growth at the gas/ water and water/n-heptane interfaces. J. Cryst. Growth 2008, 310, 1154−1166. (18) King, H. E., Jr.; Hutter, J. L.; Lin, M. Y.; Sun, T. J. Polymer conformations of gas-hydrate kinetic inhibitors: A small angle neutron scattering study. J. Chem. Phys. 2000, 112, 2523−2532. (19) Carver, T. J.; Drew, M. G. B.; Rodger, P. M. Inhibition of crystal growth in methane hydrate. J. Chem. Soc., Faraday Trans. 1995, 91, 3449−3460. (20) Daraboina, N.; Ripmeester, J. A.; Walker, V. K.; Englezos, P. Natural gas hydrate formation and decomposition in the presence of kinetic inhibitors. 2. Stirred reactor experiments. Energy Fuels 2011, 25, 4384−4391. (21) Sharifi, H.; Ripmeester, J.; Walker, V. K.; Englezos, P. Kinetic inhibition of natural gas hydrates in saline solutions and heptane. Fuel 2014, 117, 109−117. (22) Lachance, J. W.; Sloan, E. D.; Koh, C. A. Determining gas hydrate kinetic inhibitor effectiveness using emulsions. Chem. Eng. Sci. 2009, 64, 180−184. (23) Mccullough, J. P.; Waddington, G. Melting-point purity determinationslimitations as evidenced by calorimetric studies in the melting region. Anal. Chim. Acta 1957, 17, 80−96. (24) Tetervak, A.; Hudson, S.; Zhang, J.; Hutter, J. L. Continuum model for mesh crystallization. Phys. Rev. E 2005, 71, 051606 DOI: 10.1103/PhysRevE.71.051606. (25) Sakaguchi, H.; Ohmura, R.; Mori, Y. H. Effects of kinetic inhibitors on the formation and growth of hydrate crystals at a liquidliquid interface. J. Cryst. Growth 2003, 247, 631−641. (26) Sloan, E. D. Clathrate Hydrates of Natural Gases, 2nd ed.; CRC Press, Marcel Dekker, 1998. (27) Gulbrandsen, A.C.; Svartås, T. M. Effect of Poly Vinyl Caprolactam Concentration on the Dissociation Temperature for Methane Hydrates. Energy Fuels, 2017. DOI: 10.1021/acs.energyfuels.6b03487

into the hydrate stability region. This only occurred for the PVCap containing hydrates and not for the corresponding hydrates formed with gas and water. Raman spectra were recorded pre and post the transformation where the hydrate altered its appearance. The obtained spectra showed the same composition, indicating that the apparent morphological change could not be attributed to a change in composition. However, Raman spectra collected before and after the transformation showed higher water to gas ratio after the alteration had taken place. Morphological alterations may influence the stability of the different cages that constitute the various hydrate structures. The hydrates with PVCap may become stabilized to a higher degree by alterations in the binding strengths. This can alter the total energy of the cavity and as a result may be able to withstand a higher temperature before dissociation takes place.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.energyfuels.6b03478. Dissociation kinetics sII and sH hydrates; Raman spectroscopy (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Phone: + 47 971 15 592. E-mail: [email protected]. ORCID

Ann Cecilie Gulbrandsen: 0000-0001-6231-4839 Present Address #

Statoil ASA, Forusbeen 50, 4035 Stavanger, Norway

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors thank The Norwegian Ministry of Education and Research, University of Stavanger, StatoilHydro, and British Petroleum for their support to this work. The authors also thank Dr. Judith Schicks, GFZ German Research Centre for Geosciences, for the joint execution of the Raman studies at the laboratories in Potsdam. Furthermore, the authors thank Sandve & Husebø for having performed at the University of Stavanger hydrate PT experiments that have been of use.

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REFERENCES

(1) Sloan, E. D; Koh, C. A. Clathrate Hydrates of Natural Gases, 3rd ed.; CRC Press, Taylor & Francis Group: Boca Raton, 2008. (2) Fu, B. The development of advanced kinetic hydrate inhibitors. In Chemistry in the oil industry VII; Royal Society of Chemistry, ACS: Cambridge, UK, 2002. (3) Mehta, A. P.; Hebert, P. B.; Cadena, E. R.; Weatherman, J. P. Fulfilling the promise of low-dosage hydrate inhibitors: Journey from academic curiosity to successful field implementation. SPE Prod. Facil. 2003, 18, 73−79. (4) Anderson, B. J.; Tester, J. W.; Borghi, G. P.; Trout, B. L. Properties of inhibitors of methane hydrate formation via molecular dynamics simulations. J. Am. Chem. Soc. 2005, 127, 17852−17862. (5) Bruusgaard, H.; Lessard, L. D.; Servio, P. Morphology Study of Structure I Methane Hydrate Formation and Decomposition of Water Droplets in the Presence of Biological and Polymeric Kinetic Inhibitors. Cryst. Growth Des. 2009, 9, 3014−3023. 9872

DOI: 10.1021/acs.energyfuels.6b03478 Energy Fuels 2017, 31, 9863−9873

Article

Energy & Fuels (28) Subramanian, S.; Sloan, E. D., Jr. Molecular measurements of methane hydrate formation. Fluid Phase Equilib. 1999, 158, 813−820. (29) Habetinova, E.; Lund, A.; Larsen, R. Hydrate dissociation under the influence of low-dosage kinetic inhibitors; 4th International Conference on Natural Gas Hydrates, 2002, Yokohama, Japan. (30) Sharifi, H.; Ripmeester, J.; Englezos, P. Recalcitrance of gas hydrate crystals formed in the presence of kinetic hydrate inhibitors. J. Nat. Gas Sci. Eng. 2016, 35, 1573−1578. (31) Larsen, R.; Knight, C. A.; Rider, K. T.; Sloan, E. D., Jr. Growth and Inhibition of Ethylene Oxide Clathrate Hydrate. Ann. N. Y. Acad. Sci. 2000, 912, 441−451. (32) Małolepsza, E.; Keyes, T. Pathways through equilibrated states with coexisting phases for gas hydrate formation. J. Phys. Chem. B 2015, 119, 15857−15865. (33) Takeya, S.; Hondoh, T.; Uchida, T. In situ observation of CO2 hydrate by X-ray diffraction. Ann. N. Y. Acad. Sci. 2000, 912, 973−982. (34) Henning, R. W.; Schultz, A. J.; Thieu, V.; Halpern, Y. Neutron diffraction studies of CO2 clathrate hydrate: formation from deuterated ice. J. Phys. Chem. A 2000, 104, 5066−5071. (35) Moudrakovski, I. L.; Sanchez, A. A.; Ratcliffe, C. I.; Ripmeester, J. A. Nucleation and growth of hydrates on ice surfaces: new insights from 129Xe NMR experiments with hyperpolarized xenon. J. Phys. Chem. B 2001, 105, 12338−12347. (36) Staykova, D. K.; Kuhs, W. F.; Salamatin, A. N.; Hansen, T. Formation of Porous Gas Hydrates from Ice Powders: Diffraction Experiments and Multistage Model. J. Phys. Chem. B 2003, 107, 10299−10311. (37) Seo, Y.; Cha, M.; Kang, S.-P.; Shin, J.-Y.; Shin, K. Catastrophic growth of gas hydrates in the presence of kinetic hydrate inhibitors. J. Phys. Chem. A 2013, 117, 13988−13995. (38) Uchida, T.; Moriwaki, M.; Takeya, S.; Ikeda, I. Y.; Ohmura, R.; Nagao, J.; et al. Two-step formation of methane-propane mixed gas hydrates in a batch-type reactor. AIChE J. 2004, 50, 518−523. (39) Schicks, J. M.; Naumann, R.; Erzinger, J.; Hester, K. C.; Koh, C. A.; Sloan, E. D., Jr. Phase transitions in mixed gas hydrates: experimental observations versus calculated data. J. Phys. Chem. B 2006, 110, 11468−11474. (40) Ripmeester, J. A.; Zeng, H.; Lu, H.; Huva, E.; Walker, V. K. Differences in nucleator adsorption may explain distinct inhibition activities of two gas hydrate kinetic inhibitors. Chem. Eng. Sci. 2008, 63, 4026−4029. (41) Rodahl, M.; Hook, F.; Krozer, A.; Brzezinski, P.; Kasemo, B. Quartz crystal microbalance setup for frequency and Q-factor measurements in gaseous and liquid environments. Rev. Sci. Instrum. 1995, 66, 3924−3930. (42) Rodahl, M.; Hook, F.; Fredriksson, C.; Keller, C.; Krozer, A.; Brzezinski, P.; Voinova, M.; Kasemo, B. Simultaneous frequency and dissipation factor QCM measurements of biomolecular adsorption and cell adhesion. Faraday Discuss. 1997, 107, 229−246. (43) Storr, M. T.; Rodger, P. M. A molecular dynamics study of the mechanism of kinetic inhibition. Ann. N. Y. Acad. Sci. 2000, 912, 669− 677.

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