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Equilibrium conditions for CO2 hydrate in porous medium
J. Chem. Thermodynamics 43 (2011) 334–338
Contents lists available at ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Equilibrium conditions for CO2 hydrate in porous medium Mingjun Yang a, Yongchen Song a,⇑, Yu Liu a, Wei-Haur Lam a, Qingping Li b a b
Key Laboratory of Ocean Energy Utilization and Energy Conservation of Ministry of Education, Dalian University of Technology, Dalian 116024, China CNOOC Research Center, Beijing 100027, China
a r t i c l e
i n f o
Article history: Received 2 March 2010 Received in revised form 26 September 2010 Accepted 9 October 2010 Available online 16 October 2010 Keywords: CO2 hydrate NaCl Porous medium Equilibrium condition
a b s t r a c t Experimental phase equilibrium conditions data for carbon dioxide (CO2) hydrate in porous medium with the presence of sodium chloride (NaCl) solution were investigated in this study. The experimental data were generated using graphic-method in presence of solutions contained (0, 0.2, 0.4, 0.6, and 0.8) mol/L NaCl. The results indicated the increase of NaCl concentration caused the enhancement in the equilibrium pressure of CO2 hydrate as the pore size and the temperature were kept the same. Effects of NaCl solutions on CO2 hydrate equilibrium conditions could be neglected when the temperature is lower than ice point. An improved model was used to predict CO2 hydrate equilibrium conditions, and the predictions showed good agreement with experimental measurements. Ó 2010 Elsevier Ltd. All rights reserved.
1. Introduction The concentration of greenhouse gas, such as CO2 and CH4, in atmosphere has gradually increased and reached to a high level, which leads to a significant climate warming and weather changes [1]. Since CO2 produced from the fossil fuel combustion is believed to be one of the most important greenhouse gases being responsible for about 64% of the enhanced ‘‘greenhouse effect”, the disposal of CO2 has become an issue of worldwide concern [2]. One of the proposed schemes is to sequester CO2 in form of gas hydrates in ocean and marine sediment. To make CO2 hydrate stabilized in these places, it is necessary to understand the thermodynamic characters for CO2 hydrate formation and dissociation. CO2 hydrate equilibrium conditions have been investigated in four kinds of environments. The first environment is bulk water. Wendland et al. [3] obtained three-phase lines, quadruple points, and critical points of the (carbon dioxide + water) binary system. After that, Yang et al. [4] studied the two phase equilibrium for (hydrate + water + CO2) systems with both of experiments and models. Due to the presence of salts in seawater and pores water of marine sediment, the second environment for CO2 hydrate equilibrium conditions was (CO2 + water + electrolytes) system. Englezos and Hall [5], Breland and Englezos [6], Dholabhai et al. [7,8] have measured equilibrium conditions data for CO2 hydrate in polymer and electrolyte solutions. After that, Kang et al. [9] tested equilibrium conditions data for CO2 hydrate in
⇑ Corresponding author. E-mail address: [email protected] (Y. Song). 0021-9614/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2010.10.007
(water + MgCl2 + CO2) system and predicted hydrate phase equilibrium with the improved van der Waals–Platteeuw model. Mohammadi et al. [10] measured equilibrium conditions data for CO2 hydrate in the presence of NaCl, KCl, and CaCl2 using a reliable isochoric technique. The third environment studied for CO2 hydrate equilibrium conditions is porous medium. Zatsepina and Buffet [11] measured CO2 hydrate stability in porous medium. They concluded that when the vapor phase of CO2 was absent, the volume of hydrate was limited by the transport of CO2 from solution. Smith et al. [12] measured equilibrium pressures for CO2 hydrate in silica gel pores with nominal radii (7.5, 5.0, and 3.0) nm, and observed they were higher than those for CO2 hydrate in bulk water. Following their studies, Kumar [13] collected experimental equilibrium conditions data for CO2 hydrate in porous medium and measured the permeability of the porous medium in the presence of hydrate by flowing through the system. When other gas components are present, the CO2 hydrate equilibrium conditions were exist as reported by Fan et al. [14,15] and Kumar et al. [16]. This can be treated as the fourth environment. Although lots of work has been done for CO2 hydrate equilibrium conditions, there is only limited data available in porous medium with the presence of electrolytes (the fifth environment), which can be used to simulate natural condition of carbon dioxide storage. Since the complexity of marine sediment which usually contains clays, organic matters, sand and other components, experiments are usually performed in glass beads which enable the study of the impact of porosity-related properties on the equilibrium conditions like capillary effects. In this research, the equilibrium conditions of CO2 hydrates with the presence of NaCl in
335
M. Yang et al. / J. Chem. Thermodynamics 43 (2011) 334–338
The experimental apparatus used in this study is shown in figure 1 and further detail of the experimental apparatus can be referred to the previous publications of our research team [17,18]. The high-pressure resistant vessel made of 316-stainless steel with a volume of 476 mL is used as the reactor. Thermocouples (produced by Yamari Industries, Japan) and two pressure transducers (produced by Nagano Keiki, Japan) are connected to the vessel. The estimated errors of temperature and pressure measurements are ±0.1 K and ±0.1 MPa, respectively. The glass beads (produced by As-One Co., Ltd., Japan) were used to form porous medium. CO2 (mass fraction 0.999) and NaCl (mass fraction 0.99) were provided solely by Dalian Guangming Special Gas Co., Ltd. and Shenyang Xinxing Reagent Factory, China, respectively. The amount of electrolytes was weighed using a high precision balance with a minimum reading up to 0.0001 g. All the chemicals were not purified and the de-ionized water was used in all the experiments. The graphic method was used to measure the phase equilibrium conditions by keeping one of the three parameters of pressure (p), volume (V) and temperature (T) constant and changing one of the resting parameters to form or decompose the hydrate. In this study, experiments were processed in constant volume. The temperature is changed during the experiment to form or decompose CO2 hydrate. The high-pressure resistant vessel containing water was pressurized with CO2 to the designed pressure. As the vessel temperature is lowered, the pressure decreases principally due to the gas contraction as well as increased gas solubility upon cooling at constant volume. Neither gas nor water is added to the system during the experiment. If hydrates begin to form, the temperature increases rapidly. When there is no pressure change, the vessel is heated to promote the dissociation of hydrate. The system pressure increases with hydrate dissociation. The hydrate phase equilibrium condition can be detected by measuring the intersection point of the cooling and heating isochors. The experimental procedures could be described as follows: dry glass beads were packed into the vessel tightly with configured solution. The vessel was kept vertically with both side valves opening to discharge solution until there was no liquid droplet. After the vessel was reconnected to the system, a vacuum pump was used to discharge the gas in the vessel. CO2 was then injected slowly into the vessel to a designed pressure and the pressure was kept constant. The amounts of residual solution and injected
3. Results and discussion The p–T curve during CO2 hydrate formation and dissociation was dependent on the initial condition of CO2. Two cases of CO2 hydrate formation process were obtained in experiments; one case was shown in figures 2 and 3. In this condition, the CO2 in the vessel was always gas–liquid phase coexistence before CO2 hydrate formation. After the bath temperature was decreased to the target value that was usually more than 4 K below the estimated equilibrium temperature (A–B), pressure and temperature of the vessel were suddenly rising simultaneously about 30 min later (B–C), which were caused by CO2 hydrate formation and phase change of CO2. When hydrate forming in the vessel, the temperature increased (due to the exothermic hydrate crystallization process) while the pressure decreased (due to
288
6.5
286
6
284
5.5
282
5
280
4.5
278
4
276
3.5
274
0
100
200
t/min
300
p/MPa
2. Experimental measurements
CO2 were all recorded. The vessel was kept at a steady ambient temperature which must be high enough to prevent CO2 hydrate from forming in 24 h. Subsequently, the bath temperature was decreased to a designed temperature which was usually lower than the predicted temperature for this pressure. When CO2 hydrates began to form in vessel, a rapid temperature increase was observed due to the formation of CO2 hydrate. The formation was considered to be completed when there was no pressure change in the system. After hydrate formation, the bath was warmed slowly to dissociate the CO2 hydrate. The pressure and temperature (p–T) condition at the end of the hydrate decomposition was considered to be CO2 hydrate phase equilibrium condition.
T/K
porous medium were investigated by experimental observations and numerical modeling. The purpose was to provide important data for understanding thermodynamic characters of CO2 hydrate in porous medium.
3 500
400
FIGURE 2. Pressure and temperature changes with the time for gas–liquid phase coexistence of CO2: the real line represent the value of pressure and other five broken lines indicate the temperature at different point.
6 A
5.75 5.5
H
5.25 G
p/MPa
5 4.75
F
4.5
C
4.25 4 3.75
B
3.5 274
FIGURE 1. Scheme of a gas hydrate experimental apparatus: (1) de-ionized water; (2) constant flow pump; (3) needle valve; (4) high pressure vessel; (5) glass beads; (6) glycol–water bath; (7) relief valve; (8) thermocouple; (9) pressure sensor; and (10) A/D module.
D
E 276
278
280
282
284
286
288
290
292
294
T/K FIGURE 3. Typical p–T curves for gas–liquid phase coexistence of CO2: N, p–T curves for CO2 in BZ-01 and O, (gas + liquid) equilibrium line of bulk CO2.
M. Yang et al. / J. Chem. Thermodynamics 43 (2011) 334–338
gas consumption as the gas is enclathrated into the hydrate lattice). At that time, the liquid CO2 changed to vapor CO2 and caused the pressure increase to counterbalance the pressure decrease and the temperature increase. Since the phase change rate of CO2 from liquid to vapor is higher than the rate of hydrate formation, the experimental phenomena showed that the pressure and temperature of the vessel increase at the same time. This can be explained according to the thermodynamic theories; the vapor pressure of CO2 was dominated by the temperature under our experimental condition. Therefore, the temperature increase caused by the hydrate formation promotes the change of CO2 from liquid to gas to achieve thermodynamic equilibrium. The temperature increases during hydrate formation were more than 5 K for figure 2. After the liquid CO2 changed to vapor CO2 completely, the pressure decreased rapidly with the formation of CO2 hydrate. Since the CO2 hydrate formation rate was low and the high temperature difference between the porous medium and bath, the temperature showed a slow decrease (C–D). When the CO2 hydrate formation process finished, the temperature decreased down to the initial setting value (D–E). Figure 3 shows the p–T curve when initial CO2 was coexisted of liquid–gas phase. After CO2 hydrate formation was completed, temperature of the vessel decreased gradually to the designed value (E) because of the temperature difference between the porous medium and bath. Then the vessel was warmed gradually about 200 min later, when the p–T condition reached to F, the hydrate began to decompose, which caused a significant pressure increase (F–G). Point G was considered as the end of hydrate decomposition, which implied the equilibrium condition for this case. After intersecting with A–B, the p–T curve was back to point A along the temperature reduction period. In figure 3, curves A–B, E–F, and (gas + liquid) equilibrium line were parallel each others. These curves were all indicating the (gas + liquid) equilibrium of CO2, and these different positions were caused by the capillary effects of the porous medium. Curve A–B indicates the (gas + liquid) equilibrium line of CO2 in the glass beads. In the E–F duration, the hydrate existed in porous medium and caused the pore sizes further reduced, which made E–F locate left to A–B. If the initial CO2 was vapor, the hydrate formation process was similar as that of methane hydrate, which has been discussed in the previous works [18]. This situation can be treated as the second case. The equilibrium condition of CO2 hydrate was affected by the presence of glass beads, as shown in figure 4. To make a comparison, the experimental results of CO2 hydrate equilibrium condition in bulk water, obtained by Deaton et al. and Larson et al., was cited. The experimental data of Smith et al. for CO2 hydrate equilibrium condition in porous medium (nominal radii: 5.0 nm) was also quoted. It can be observed that the presence of glass beads led to enhancement in CO2 hydrate equilibrium pressure as the salinity and the temperature are kept the same. This was mainly caused by additional resistance effect of capillary surface tension that leads to lower water activity and affects hydrate equilibrium condition [19]. Therefore, it is necessary to account for the additional forces that result from interactions with the porous medium, mainly the capillary forces. The effect of the capillary forces is to lower the activity of water in the pore. This, in turn, causes a depression of the freezing point of water in the pore [20]. The equilibrium condition for CO2 hydrate in porous medium can be calculated when the pore size was given, and the suppressed temperature caused by porous medium will decrease with the increase of pore size. When the pore radii are big enough, the suppressed temperature becomes very small. Turner et al. [21] reported that any shift in pores larger than 600 Å in radius cannot be distinguished from errors of the thermocouples in their equilibrium apparatus.
It is not unusual that seawater and pores water in marine sediments contain electrolytes, which can also inhibit CO2 hydrate formation. Accurate knowledge of the thermodynamic stability for CO2 hydrates as a function of both electrolyte and methanol concentrations are crucial to the success of CO2 storage in marine sediments. Since NaCl is the most popular electrolyte in seawater and pores water of marine sediment, it was chosen as the electrolyte in this study. The glass beads were used to make up of the skeleton structure of porous medium. NaCl was added into de-ionized water to configure solution with different concentrations of NaCl. The effects of different concentrations of NaCl solution on the CO2 hydrate equilibrium condition were experimental investigated. The CO2 hydrate equilibrium conditions with different concentrations of NaCl in glass beads were shown in figure 5. It could be concluded that for the same glass beads, the presence of NaCl in the solution caused the equilibrium condition to move to a higher pressure and a lower temperature, while the increase of concentrations led to the increase in the extent of equilibrium curve movement. This is mainly due to the ‘‘ion effect” caused by the ions come from the electrolytes ionization in water. The ions interact with the dipoles of water molecules with Coulombic bond, which is much stronger than hydrogen bond and van der Waals forces. The stronger bonds of water with ions inhibit CO2 hydrate formation, that is to say water is attracted to ions rather than to the hydrate structure. As a result, the equilibrium temperature of CO2 hydrate was reduced [22–25].
0.8 0.7 0.6 0.5
lg(p/MPa)
336
0.4 0.3 0.2 0.1 0 268
270
272
274
276
278
280
282
284
286
T/K FIGURE 5. Measurement of equilibrium condition data for CO2 hydrate in NaCl solutions: N, pure water; 4, 0.2 mol/L NaCl; ., 0.4 mol/L NaCl; 5, 0.6mol/L NaCl; d, 0.8 mol/L NaCl; s, bulk CO2 hydrate (Deaton et al.); r , bulk CO2 hydrate (Larson et al.).
6
0.8
5.5 0.7
5 4.5
0.6
3.5
lg(p/MPa)
p/MPa
4
3 2.5 2
0.5 0.4 0.3
1.5 1
0.2
0.5 0 264
266
268
270
272
274
276
278
280
282
284
286
T/K FIGURE 4. Comparison of CO2 hydrate equilibrium condition data for in bulk water and porous medium reported in literature and in glass beads measured in this study: 4, bulk CO2 hydrate (Deaton et al.); h, bulk CO2 hydrate (Larson et al.); N, CO2 hydrate in BZ-01; O, CO2 hydrate in 5.0 nm nominal radii porous media (Smith et al.).
0.1 271
272
273
274
275
276
277
278
279
280
281
282
283
T/K FIGURE 6. Comparison of measured and calculated equilibrium condition data for CO2 hydrate in BZ-01 with different NaCl solutions. Symbols denote experimental data: 4, pure water; N, 0.2mol/L NaCl; }, 0.4 mol/L NaCl; r, 0.6 mol/L NaCl; s, 0.8 mol/L NaCl; lines represent the calculated results: —, pure water; – – –, 0.2 mol/L NaCl; ––, 0.4 mol/L NaCl; ––, 0.6 mol/L NaCl; - - - -, 0.8 mol/L NaCl.
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M. Yang et al. / J. Chem. Thermodynamics 43 (2011) 334–338 TABLE 1 Absolute average deviations of predicted CO2 hydrate formation conditions for NaCl solutions.
a
ca/(mol/L)
T/K Minimum
Maximum
Minimum
Maximum
0 0.2 0.4 0.6 0.8
273.3 273.1 273.7 274.2 277.5
281.9 281.8 282.9 282.3 281.7
1.5 1.6 1.7 1.8 2.1
3.4 3.6 3.9 3.8 3.9
p/MPa
Np
DAADT/%
DAADP/%
8 10 12 12 11
0.09 0.14 0.26 0.26 0.30
2.82 4.1 7.58 6.72 8.49
Molar fraction of NaCl in configured solution.
When temperature was lower than ice point, the experimental data indicated that the CO2 hydrate equilibrium points were nearly at the same curve. This may be caused by the change of water activity with temperature. Water activity is a dimensionless quantity used to represent the energy status of the water in a system. It is defined as the vapor pressure of water above a sample divided by that of pure water at the same temperature. Therefore, pure distilled water has a water activity of exactly one. The definition of water activity can be expressed as:
high pressure and high salt concentration conditions. However, this is not surprising due to the complexity of the system. The result can be concluded without doubt that the improved model gave good predictions for the CO2 hydrate in porous medium.
aw ¼ pw =p0w ;
The effects of NaCl solution on CO2 hydrate phase equilibrium conditions were investigated in porous medium. Two kinds of hydrate formation cases were obtained due to the different initial p–T conditions of CO2. The results indicated the present of glass bead leads to the enhancement in the equilibrium pressure as the salinity and the temperature are kept the same. The increase of NaCl concentration caused the enhancement in the equilibrium pressure as the pore size and the temperature are kept the same. The effects of NaCl solution on hydrate equilibrium can be neglected when temperature was lower than ice points. This can be explained that the vapor pressure of solution was equal to that of solid phase (ice) when the equilibrium temperature reaches ice point. The improved model proposed by Song et al. was used to predict the phase equilibrium condition, the CO2 solubility in pore water was calculated using an empirical modification of K–K equation. The predictions showed good agreement with the experimental measurements.
ð1Þ
where pw is the vapor pressure of water in the substance, and p0w is the vapor pressure of pure water at the same temperature. The vapor pressure of water is the vapor pressure exerted by water at a specific temperature. When NaCl is present in solution, the vapor pressure will decrease due to cohesive forces between the electrolyte and water. Seen from periodic table of electronegativities, we will find that sodium chloride is considerably more-polar than water. Due to that extra polarity, water can make much stronger bonds to sodium chloride than it can to itself. Stronger intramolecular forces in solution always cause the vapor pressure to decrease. So the water activity decreases with the adding of NaCl. When the equilibrium temperature gets to ice point, the vapor pressure of solution will be equal to that of solid phase (ice). Since the ice vapor pressure is constant, the presence of electrolyte has no effect on it. The following equation can be used to express this situation:
pw ¼ p0w :
ð2Þ
That was to say, the water activity was equal to 1 when temperature got closer to ice point. The presence of NaCl had no effects on CO2 hydrate equilibrium condition. The improved model of Song et al. [26] was used for predicting the equilibrium conditions for CO2 hydrates in bulk water and in glass beads with/without electrolytes, in which the mechanical equilibrium of force between the interfaces in (hydrate + liquid + vapor) system was considered. The improved model was based on the traditional model of van der Waals and Plateeuw [27]. In this model, Gibbs– Thomson equation with correct parameters for hydrate–water interface modified by Henry et al. [28] was used to account for the capillary effect in porous sediments, and mechanical equilibrium of force between the interfaces in (hydrate + liquid + vapor) system was considered. According to the mechanical equilibrium relations to (H + L + V) system, interfacial energy between hydrate and liquid was corrected by the function that is expressed with temperature and electrolyte concentration when electrolyte was in pore water. The activity of water is calculated using the Pitzer model and the interfacial energy between liquid and gas is solved using the Li’s method. The gas fugacity was calculated using modified Patel–Teja equation of state. When hydrate former is CO2, the solubility of former gas cannot be neglected. An empirical modification of Krichevsky–Kasarnovsky (K–K) equation was used to calculate Xgas [29]. This equation also includes the effect of pressure on the solubility. Having calculated the Xgas, Xw is calculated from:
X W ¼ 1 X gas ¼ 1 fgas =½g3 expðpg1 g2 =RTÞ;
ð3Þ
where g1, g2, and g3 are equal to 0.3411, 22.33, and 899683.5 for CO2, which can be found in the study of Nasrifar [29]. Figure 6 shows the comparison between the experimental results and the thermodynamic model of Song et al. [26] for CO2 hydrate phase equilibrium condition. The predictions show a good agreement with the experimental data. The absolute average deviation of predicted temperature (DAADT) and pressure (DAADP) are defined as follows: [30]
DAADT ¼ ð1=N p Þ
Np X ½jT cal T exp j=T exp j 100;
ð4Þ
j¼1
DAADP ¼ ð1=N p Þ
Np X ½jpcal pexp j=pexp j 100;
ð5Þ
j¼1
where Np denotes the number of data points. Error analysis of the predicted CO2 hydrate equilibrium conditions for different NaCl concentration solution is shown in table 1. The overall DAADT value for the improved model was 0.3% respectively; and DAADP were usually less 8.5%. There is a significant deviation between the measurements and the modeling results at both
4. Conclusions
Acknowledgements This work was supported by the Key program of National Natural Science Foundation of China (50736001), the High-tech Research and Development Program of China (2006AA09A209-5) and the Major State Basic Research Development Program of China (2009CB219507) and the Fundamental Research Funds for the Central Universities of China. References [1] Y.H. Ji, X.Y. Ji, X. Feng, C. Liu, L.H. Lv, X.H. Lu, Chin. J. Chem. Eng. 15 (2007) 439– 448. [2] S. Bachu, Energy Convers. Manage. 41 (2000) 953–970. [3] M. Wendland, H. Hasse, G. Maurer, J. Chem. Eng. Data 44 (1999) 901–906. [4] S.O. Yang, I.M. Yang, Y.S. Kim, C.S. Lee, Fluid Phase Equilib. 175 (2000) 75–89. [5] P. Englezos, S. Hall, Can. J. Chem. Eng. 72 (1994) 887–893. [6] E. Breland, P. Englezos, J. Chem. Eng. Data 41 (1996) 11–13. [7] P.D. Dholabhai, J. Scott Parent, P. Raj Bishnoi, Ind. Eng. Chem. Res. 35 (1996) 819–823. [8] P.D. Dholabhai, J. Scott Parent, P. Raj Bishnoi, Fluid Phase Equilib. 41 (1997) 235–246. [9] S.P. Kang, M.K. Chun, H. Lee, Fluid Phase Equilib. 147 (1998) 229–238. [10] A.H. Mohammadi, W. Afzal, D. Richon, J. Chem. Thermodyn. 40 (2008) 1693– 1697. [11] O.Y. Zatsepina, B.A. Buffet, Fluid Phase Equilib. 200 (2002) 263–275. [12] D. Smith, J. Wilder, K. Seshadri, Environ. Sci. Technol. 36 (2002) 5192–5198. [13] A. Kumar, Formation and Dissociation of Gas Hydrates in Porous Media, University of Calgary, 2005. [14] S.S. Fan, T.M. Gu, J. Chem. Eng. Data 44 (1999) 829–832. [15] S.S. Fan, G.J. Chen, Q.L. Ma, T.M. Guo, Chem. Eng. J. 78 (2000) 173–178. [16] R. Kumar, H.J. Wu, P. Englezos, Fluid Phase Equilib. 244 (2006) 167–171. [17] Y.C. Song, M.J. Yang, Y. Liu, Q.P. Li, CIESC J. 60 (2009) 1362–1366. [18] M.J. Yang, Y.C. Song, Y. Liu, Y.J. Chen, Q.P. Li, Chin. J. Chem. Eng. 18 (2010) 292– 296. [19] J.W. Wilder, K. Seshadri, D.H. Smith, Langmuir 17 (2001) 6729–6735. [20] M.A. Clarke, M.P. Darvish, P.R. Bishnoi, Ind. Eng. Chem. Res. 38 (1999) 2485– 2490.
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[21] D.J. Turner, R.S. Cherry, E.D. Sloan, Fluid Phase Equilib. 228–229 (2005) 505– 510. [22] G.M. Marion, D.C. Catling, J.S. Kargel, Computer Coupling of Phase Diagrams and Thermochemistry 30 (2006) 248–259. [23] K. Nasrifar, M. Moshfeghian, Fluid Phase Equilib. 146 (1998) 1–13. [24] K. Aasberg-Petersen, E. Stenby, A. Fredenslund, Ind. Eng. Chem. Res. 30 (1991) 2180–2185. [25] J. Javanmardi, M. Moshfeghian, Fluid Phase Equilib. 168 (2000) 135–148.
[26] [27] [28] [29] [30]
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JCT 09-414
Contents lists available at ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Equilibrium conditions for CO2 hydrate in porous medium Mingjun Yang a, Yongchen Song a,⇑, Yu Liu a, Wei-Haur Lam a, Qingping Li b a b
Key Laboratory of Ocean Energy Utilization and Energy Conservation of Ministry of Education, Dalian University of Technology, Dalian 116024, China CNOOC Research Center, Beijing 100027, China
a r t i c l e
i n f o
Article history: Received 2 March 2010 Received in revised form 26 September 2010 Accepted 9 October 2010 Available online 16 October 2010 Keywords: CO2 hydrate NaCl Porous medium Equilibrium condition
a b s t r a c t Experimental phase equilibrium conditions data for carbon dioxide (CO2) hydrate in porous medium with the presence of sodium chloride (NaCl) solution were investigated in this study. The experimental data were generated using graphic-method in presence of solutions contained (0, 0.2, 0.4, 0.6, and 0.8) mol/L NaCl. The results indicated the increase of NaCl concentration caused the enhancement in the equilibrium pressure of CO2 hydrate as the pore size and the temperature were kept the same. Effects of NaCl solutions on CO2 hydrate equilibrium conditions could be neglected when the temperature is lower than ice point. An improved model was used to predict CO2 hydrate equilibrium conditions, and the predictions showed good agreement with experimental measurements. Ó 2010 Elsevier Ltd. All rights reserved.
1. Introduction The concentration of greenhouse gas, such as CO2 and CH4, in atmosphere has gradually increased and reached to a high level, which leads to a significant climate warming and weather changes [1]. Since CO2 produced from the fossil fuel combustion is believed to be one of the most important greenhouse gases being responsible for about 64% of the enhanced ‘‘greenhouse effect”, the disposal of CO2 has become an issue of worldwide concern [2]. One of the proposed schemes is to sequester CO2 in form of gas hydrates in ocean and marine sediment. To make CO2 hydrate stabilized in these places, it is necessary to understand the thermodynamic characters for CO2 hydrate formation and dissociation. CO2 hydrate equilibrium conditions have been investigated in four kinds of environments. The first environment is bulk water. Wendland et al. [3] obtained three-phase lines, quadruple points, and critical points of the (carbon dioxide + water) binary system. After that, Yang et al. [4] studied the two phase equilibrium for (hydrate + water + CO2) systems with both of experiments and models. Due to the presence of salts in seawater and pores water of marine sediment, the second environment for CO2 hydrate equilibrium conditions was (CO2 + water + electrolytes) system. Englezos and Hall [5], Breland and Englezos [6], Dholabhai et al. [7,8] have measured equilibrium conditions data for CO2 hydrate in polymer and electrolyte solutions. After that, Kang et al. [9] tested equilibrium conditions data for CO2 hydrate in
⇑ Corresponding author. E-mail address: [email protected] (Y. Song). 0021-9614/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2010.10.007
(water + MgCl2 + CO2) system and predicted hydrate phase equilibrium with the improved van der Waals–Platteeuw model. Mohammadi et al. [10] measured equilibrium conditions data for CO2 hydrate in the presence of NaCl, KCl, and CaCl2 using a reliable isochoric technique. The third environment studied for CO2 hydrate equilibrium conditions is porous medium. Zatsepina and Buffet [11] measured CO2 hydrate stability in porous medium. They concluded that when the vapor phase of CO2 was absent, the volume of hydrate was limited by the transport of CO2 from solution. Smith et al. [12] measured equilibrium pressures for CO2 hydrate in silica gel pores with nominal radii (7.5, 5.0, and 3.0) nm, and observed they were higher than those for CO2 hydrate in bulk water. Following their studies, Kumar [13] collected experimental equilibrium conditions data for CO2 hydrate in porous medium and measured the permeability of the porous medium in the presence of hydrate by flowing through the system. When other gas components are present, the CO2 hydrate equilibrium conditions were exist as reported by Fan et al. [14,15] and Kumar et al. [16]. This can be treated as the fourth environment. Although lots of work has been done for CO2 hydrate equilibrium conditions, there is only limited data available in porous medium with the presence of electrolytes (the fifth environment), which can be used to simulate natural condition of carbon dioxide storage. Since the complexity of marine sediment which usually contains clays, organic matters, sand and other components, experiments are usually performed in glass beads which enable the study of the impact of porosity-related properties on the equilibrium conditions like capillary effects. In this research, the equilibrium conditions of CO2 hydrates with the presence of NaCl in
335
M. Yang et al. / J. Chem. Thermodynamics 43 (2011) 334–338
The experimental apparatus used in this study is shown in figure 1 and further detail of the experimental apparatus can be referred to the previous publications of our research team [17,18]. The high-pressure resistant vessel made of 316-stainless steel with a volume of 476 mL is used as the reactor. Thermocouples (produced by Yamari Industries, Japan) and two pressure transducers (produced by Nagano Keiki, Japan) are connected to the vessel. The estimated errors of temperature and pressure measurements are ±0.1 K and ±0.1 MPa, respectively. The glass beads (produced by As-One Co., Ltd., Japan) were used to form porous medium. CO2 (mass fraction 0.999) and NaCl (mass fraction 0.99) were provided solely by Dalian Guangming Special Gas Co., Ltd. and Shenyang Xinxing Reagent Factory, China, respectively. The amount of electrolytes was weighed using a high precision balance with a minimum reading up to 0.0001 g. All the chemicals were not purified and the de-ionized water was used in all the experiments. The graphic method was used to measure the phase equilibrium conditions by keeping one of the three parameters of pressure (p), volume (V) and temperature (T) constant and changing one of the resting parameters to form or decompose the hydrate. In this study, experiments were processed in constant volume. The temperature is changed during the experiment to form or decompose CO2 hydrate. The high-pressure resistant vessel containing water was pressurized with CO2 to the designed pressure. As the vessel temperature is lowered, the pressure decreases principally due to the gas contraction as well as increased gas solubility upon cooling at constant volume. Neither gas nor water is added to the system during the experiment. If hydrates begin to form, the temperature increases rapidly. When there is no pressure change, the vessel is heated to promote the dissociation of hydrate. The system pressure increases with hydrate dissociation. The hydrate phase equilibrium condition can be detected by measuring the intersection point of the cooling and heating isochors. The experimental procedures could be described as follows: dry glass beads were packed into the vessel tightly with configured solution. The vessel was kept vertically with both side valves opening to discharge solution until there was no liquid droplet. After the vessel was reconnected to the system, a vacuum pump was used to discharge the gas in the vessel. CO2 was then injected slowly into the vessel to a designed pressure and the pressure was kept constant. The amounts of residual solution and injected
3. Results and discussion The p–T curve during CO2 hydrate formation and dissociation was dependent on the initial condition of CO2. Two cases of CO2 hydrate formation process were obtained in experiments; one case was shown in figures 2 and 3. In this condition, the CO2 in the vessel was always gas–liquid phase coexistence before CO2 hydrate formation. After the bath temperature was decreased to the target value that was usually more than 4 K below the estimated equilibrium temperature (A–B), pressure and temperature of the vessel were suddenly rising simultaneously about 30 min later (B–C), which were caused by CO2 hydrate formation and phase change of CO2. When hydrate forming in the vessel, the temperature increased (due to the exothermic hydrate crystallization process) while the pressure decreased (due to
288
6.5
286
6
284
5.5
282
5
280
4.5
278
4
276
3.5
274
0
100
200
t/min
300
p/MPa
2. Experimental measurements
CO2 were all recorded. The vessel was kept at a steady ambient temperature which must be high enough to prevent CO2 hydrate from forming in 24 h. Subsequently, the bath temperature was decreased to a designed temperature which was usually lower than the predicted temperature for this pressure. When CO2 hydrates began to form in vessel, a rapid temperature increase was observed due to the formation of CO2 hydrate. The formation was considered to be completed when there was no pressure change in the system. After hydrate formation, the bath was warmed slowly to dissociate the CO2 hydrate. The pressure and temperature (p–T) condition at the end of the hydrate decomposition was considered to be CO2 hydrate phase equilibrium condition.
T/K
porous medium were investigated by experimental observations and numerical modeling. The purpose was to provide important data for understanding thermodynamic characters of CO2 hydrate in porous medium.
3 500
400
FIGURE 2. Pressure and temperature changes with the time for gas–liquid phase coexistence of CO2: the real line represent the value of pressure and other five broken lines indicate the temperature at different point.
6 A
5.75 5.5
H
5.25 G
p/MPa
5 4.75
F
4.5
C
4.25 4 3.75
B
3.5 274
FIGURE 1. Scheme of a gas hydrate experimental apparatus: (1) de-ionized water; (2) constant flow pump; (3) needle valve; (4) high pressure vessel; (5) glass beads; (6) glycol–water bath; (7) relief valve; (8) thermocouple; (9) pressure sensor; and (10) A/D module.
D
E 276
278
280
282
284
286
288
290
292
294
T/K FIGURE 3. Typical p–T curves for gas–liquid phase coexistence of CO2: N, p–T curves for CO2 in BZ-01 and O, (gas + liquid) equilibrium line of bulk CO2.
M. Yang et al. / J. Chem. Thermodynamics 43 (2011) 334–338
gas consumption as the gas is enclathrated into the hydrate lattice). At that time, the liquid CO2 changed to vapor CO2 and caused the pressure increase to counterbalance the pressure decrease and the temperature increase. Since the phase change rate of CO2 from liquid to vapor is higher than the rate of hydrate formation, the experimental phenomena showed that the pressure and temperature of the vessel increase at the same time. This can be explained according to the thermodynamic theories; the vapor pressure of CO2 was dominated by the temperature under our experimental condition. Therefore, the temperature increase caused by the hydrate formation promotes the change of CO2 from liquid to gas to achieve thermodynamic equilibrium. The temperature increases during hydrate formation were more than 5 K for figure 2. After the liquid CO2 changed to vapor CO2 completely, the pressure decreased rapidly with the formation of CO2 hydrate. Since the CO2 hydrate formation rate was low and the high temperature difference between the porous medium and bath, the temperature showed a slow decrease (C–D). When the CO2 hydrate formation process finished, the temperature decreased down to the initial setting value (D–E). Figure 3 shows the p–T curve when initial CO2 was coexisted of liquid–gas phase. After CO2 hydrate formation was completed, temperature of the vessel decreased gradually to the designed value (E) because of the temperature difference between the porous medium and bath. Then the vessel was warmed gradually about 200 min later, when the p–T condition reached to F, the hydrate began to decompose, which caused a significant pressure increase (F–G). Point G was considered as the end of hydrate decomposition, which implied the equilibrium condition for this case. After intersecting with A–B, the p–T curve was back to point A along the temperature reduction period. In figure 3, curves A–B, E–F, and (gas + liquid) equilibrium line were parallel each others. These curves were all indicating the (gas + liquid) equilibrium of CO2, and these different positions were caused by the capillary effects of the porous medium. Curve A–B indicates the (gas + liquid) equilibrium line of CO2 in the glass beads. In the E–F duration, the hydrate existed in porous medium and caused the pore sizes further reduced, which made E–F locate left to A–B. If the initial CO2 was vapor, the hydrate formation process was similar as that of methane hydrate, which has been discussed in the previous works [18]. This situation can be treated as the second case. The equilibrium condition of CO2 hydrate was affected by the presence of glass beads, as shown in figure 4. To make a comparison, the experimental results of CO2 hydrate equilibrium condition in bulk water, obtained by Deaton et al. and Larson et al., was cited. The experimental data of Smith et al. for CO2 hydrate equilibrium condition in porous medium (nominal radii: 5.0 nm) was also quoted. It can be observed that the presence of glass beads led to enhancement in CO2 hydrate equilibrium pressure as the salinity and the temperature are kept the same. This was mainly caused by additional resistance effect of capillary surface tension that leads to lower water activity and affects hydrate equilibrium condition [19]. Therefore, it is necessary to account for the additional forces that result from interactions with the porous medium, mainly the capillary forces. The effect of the capillary forces is to lower the activity of water in the pore. This, in turn, causes a depression of the freezing point of water in the pore [20]. The equilibrium condition for CO2 hydrate in porous medium can be calculated when the pore size was given, and the suppressed temperature caused by porous medium will decrease with the increase of pore size. When the pore radii are big enough, the suppressed temperature becomes very small. Turner et al. [21] reported that any shift in pores larger than 600 Å in radius cannot be distinguished from errors of the thermocouples in their equilibrium apparatus.
It is not unusual that seawater and pores water in marine sediments contain electrolytes, which can also inhibit CO2 hydrate formation. Accurate knowledge of the thermodynamic stability for CO2 hydrates as a function of both electrolyte and methanol concentrations are crucial to the success of CO2 storage in marine sediments. Since NaCl is the most popular electrolyte in seawater and pores water of marine sediment, it was chosen as the electrolyte in this study. The glass beads were used to make up of the skeleton structure of porous medium. NaCl was added into de-ionized water to configure solution with different concentrations of NaCl. The effects of different concentrations of NaCl solution on the CO2 hydrate equilibrium condition were experimental investigated. The CO2 hydrate equilibrium conditions with different concentrations of NaCl in glass beads were shown in figure 5. It could be concluded that for the same glass beads, the presence of NaCl in the solution caused the equilibrium condition to move to a higher pressure and a lower temperature, while the increase of concentrations led to the increase in the extent of equilibrium curve movement. This is mainly due to the ‘‘ion effect” caused by the ions come from the electrolytes ionization in water. The ions interact with the dipoles of water molecules with Coulombic bond, which is much stronger than hydrogen bond and van der Waals forces. The stronger bonds of water with ions inhibit CO2 hydrate formation, that is to say water is attracted to ions rather than to the hydrate structure. As a result, the equilibrium temperature of CO2 hydrate was reduced [22–25].
0.8 0.7 0.6 0.5
lg(p/MPa)
336
0.4 0.3 0.2 0.1 0 268
270
272
274
276
278
280
282
284
286
T/K FIGURE 5. Measurement of equilibrium condition data for CO2 hydrate in NaCl solutions: N, pure water; 4, 0.2 mol/L NaCl; ., 0.4 mol/L NaCl; 5, 0.6mol/L NaCl; d, 0.8 mol/L NaCl; s, bulk CO2 hydrate (Deaton et al.); r , bulk CO2 hydrate (Larson et al.).
6
0.8
5.5 0.7
5 4.5
0.6
3.5
lg(p/MPa)
p/MPa
4
3 2.5 2
0.5 0.4 0.3
1.5 1
0.2
0.5 0 264
266
268
270
272
274
276
278
280
282
284
286
T/K FIGURE 4. Comparison of CO2 hydrate equilibrium condition data for in bulk water and porous medium reported in literature and in glass beads measured in this study: 4, bulk CO2 hydrate (Deaton et al.); h, bulk CO2 hydrate (Larson et al.); N, CO2 hydrate in BZ-01; O, CO2 hydrate in 5.0 nm nominal radii porous media (Smith et al.).
0.1 271
272
273
274
275
276
277
278
279
280
281
282
283
T/K FIGURE 6. Comparison of measured and calculated equilibrium condition data for CO2 hydrate in BZ-01 with different NaCl solutions. Symbols denote experimental data: 4, pure water; N, 0.2mol/L NaCl; }, 0.4 mol/L NaCl; r, 0.6 mol/L NaCl; s, 0.8 mol/L NaCl; lines represent the calculated results: —, pure water; – – –, 0.2 mol/L NaCl; ––, 0.4 mol/L NaCl; ––, 0.6 mol/L NaCl; - - - -, 0.8 mol/L NaCl.
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M. Yang et al. / J. Chem. Thermodynamics 43 (2011) 334–338 TABLE 1 Absolute average deviations of predicted CO2 hydrate formation conditions for NaCl solutions.
a
ca/(mol/L)
T/K Minimum
Maximum
Minimum
Maximum
0 0.2 0.4 0.6 0.8
273.3 273.1 273.7 274.2 277.5
281.9 281.8 282.9 282.3 281.7
1.5 1.6 1.7 1.8 2.1
3.4 3.6 3.9 3.8 3.9
p/MPa
Np
DAADT/%
DAADP/%
8 10 12 12 11
0.09 0.14 0.26 0.26 0.30
2.82 4.1 7.58 6.72 8.49
Molar fraction of NaCl in configured solution.
When temperature was lower than ice point, the experimental data indicated that the CO2 hydrate equilibrium points were nearly at the same curve. This may be caused by the change of water activity with temperature. Water activity is a dimensionless quantity used to represent the energy status of the water in a system. It is defined as the vapor pressure of water above a sample divided by that of pure water at the same temperature. Therefore, pure distilled water has a water activity of exactly one. The definition of water activity can be expressed as:
high pressure and high salt concentration conditions. However, this is not surprising due to the complexity of the system. The result can be concluded without doubt that the improved model gave good predictions for the CO2 hydrate in porous medium.
aw ¼ pw =p0w ;
The effects of NaCl solution on CO2 hydrate phase equilibrium conditions were investigated in porous medium. Two kinds of hydrate formation cases were obtained due to the different initial p–T conditions of CO2. The results indicated the present of glass bead leads to the enhancement in the equilibrium pressure as the salinity and the temperature are kept the same. The increase of NaCl concentration caused the enhancement in the equilibrium pressure as the pore size and the temperature are kept the same. The effects of NaCl solution on hydrate equilibrium can be neglected when temperature was lower than ice points. This can be explained that the vapor pressure of solution was equal to that of solid phase (ice) when the equilibrium temperature reaches ice point. The improved model proposed by Song et al. was used to predict the phase equilibrium condition, the CO2 solubility in pore water was calculated using an empirical modification of K–K equation. The predictions showed good agreement with the experimental measurements.
ð1Þ
where pw is the vapor pressure of water in the substance, and p0w is the vapor pressure of pure water at the same temperature. The vapor pressure of water is the vapor pressure exerted by water at a specific temperature. When NaCl is present in solution, the vapor pressure will decrease due to cohesive forces between the electrolyte and water. Seen from periodic table of electronegativities, we will find that sodium chloride is considerably more-polar than water. Due to that extra polarity, water can make much stronger bonds to sodium chloride than it can to itself. Stronger intramolecular forces in solution always cause the vapor pressure to decrease. So the water activity decreases with the adding of NaCl. When the equilibrium temperature gets to ice point, the vapor pressure of solution will be equal to that of solid phase (ice). Since the ice vapor pressure is constant, the presence of electrolyte has no effect on it. The following equation can be used to express this situation:
pw ¼ p0w :
ð2Þ
That was to say, the water activity was equal to 1 when temperature got closer to ice point. The presence of NaCl had no effects on CO2 hydrate equilibrium condition. The improved model of Song et al. [26] was used for predicting the equilibrium conditions for CO2 hydrates in bulk water and in glass beads with/without electrolytes, in which the mechanical equilibrium of force between the interfaces in (hydrate + liquid + vapor) system was considered. The improved model was based on the traditional model of van der Waals and Plateeuw [27]. In this model, Gibbs– Thomson equation with correct parameters for hydrate–water interface modified by Henry et al. [28] was used to account for the capillary effect in porous sediments, and mechanical equilibrium of force between the interfaces in (hydrate + liquid + vapor) system was considered. According to the mechanical equilibrium relations to (H + L + V) system, interfacial energy between hydrate and liquid was corrected by the function that is expressed with temperature and electrolyte concentration when electrolyte was in pore water. The activity of water is calculated using the Pitzer model and the interfacial energy between liquid and gas is solved using the Li’s method. The gas fugacity was calculated using modified Patel–Teja equation of state. When hydrate former is CO2, the solubility of former gas cannot be neglected. An empirical modification of Krichevsky–Kasarnovsky (K–K) equation was used to calculate Xgas [29]. This equation also includes the effect of pressure on the solubility. Having calculated the Xgas, Xw is calculated from:
X W ¼ 1 X gas ¼ 1 fgas =½g3 expðpg1 g2 =RTÞ;
ð3Þ
where g1, g2, and g3 are equal to 0.3411, 22.33, and 899683.5 for CO2, which can be found in the study of Nasrifar [29]. Figure 6 shows the comparison between the experimental results and the thermodynamic model of Song et al. [26] for CO2 hydrate phase equilibrium condition. The predictions show a good agreement with the experimental data. The absolute average deviation of predicted temperature (DAADT) and pressure (DAADP) are defined as follows: [30]
DAADT ¼ ð1=N p Þ
Np X ½jT cal T exp j=T exp j 100;
ð4Þ
j¼1
DAADP ¼ ð1=N p Þ
Np X ½jpcal pexp j=pexp j 100;
ð5Þ
j¼1
where Np denotes the number of data points. Error analysis of the predicted CO2 hydrate equilibrium conditions for different NaCl concentration solution is shown in table 1. The overall DAADT value for the improved model was 0.3% respectively; and DAADP were usually less 8.5%. There is a significant deviation between the measurements and the modeling results at both
4. Conclusions
Acknowledgements This work was supported by the Key program of National Natural Science Foundation of China (50736001), the High-tech Research and Development Program of China (2006AA09A209-5) and the Major State Basic Research Development Program of China (2009CB219507) and the Fundamental Research Funds for the Central Universities of China. References [1] Y.H. Ji, X.Y. Ji, X. Feng, C. Liu, L.H. Lv, X.H. Lu, Chin. J. Chem. Eng. 15 (2007) 439– 448. [2] S. Bachu, Energy Convers. Manage. 41 (2000) 953–970. [3] M. Wendland, H. Hasse, G. Maurer, J. Chem. Eng. Data 44 (1999) 901–906. [4] S.O. Yang, I.M. Yang, Y.S. Kim, C.S. Lee, Fluid Phase Equilib. 175 (2000) 75–89. [5] P. Englezos, S. Hall, Can. J. Chem. Eng. 72 (1994) 887–893. [6] E. Breland, P. Englezos, J. Chem. Eng. Data 41 (1996) 11–13. [7] P.D. Dholabhai, J. Scott Parent, P. Raj Bishnoi, Ind. Eng. Chem. Res. 35 (1996) 819–823. [8] P.D. Dholabhai, J. Scott Parent, P. Raj Bishnoi, Fluid Phase Equilib. 41 (1997) 235–246. [9] S.P. Kang, M.K. Chun, H. Lee, Fluid Phase Equilib. 147 (1998) 229–238. [10] A.H. Mohammadi, W. Afzal, D. Richon, J. Chem. Thermodyn. 40 (2008) 1693– 1697. [11] O.Y. Zatsepina, B.A. Buffet, Fluid Phase Equilib. 200 (2002) 263–275. [12] D. Smith, J. Wilder, K. Seshadri, Environ. Sci. Technol. 36 (2002) 5192–5198. [13] A. Kumar, Formation and Dissociation of Gas Hydrates in Porous Media, University of Calgary, 2005. [14] S.S. Fan, T.M. Gu, J. Chem. Eng. Data 44 (1999) 829–832. [15] S.S. Fan, G.J. Chen, Q.L. Ma, T.M. Guo, Chem. Eng. J. 78 (2000) 173–178. [16] R. Kumar, H.J. Wu, P. Englezos, Fluid Phase Equilib. 244 (2006) 167–171. [17] Y.C. Song, M.J. Yang, Y. Liu, Q.P. Li, CIESC J. 60 (2009) 1362–1366. [18] M.J. Yang, Y.C. Song, Y. Liu, Y.J. Chen, Q.P. Li, Chin. J. Chem. Eng. 18 (2010) 292– 296. [19] J.W. Wilder, K. Seshadri, D.H. Smith, Langmuir 17 (2001) 6729–6735. [20] M.A. Clarke, M.P. Darvish, P.R. Bishnoi, Ind. Eng. Chem. Res. 38 (1999) 2485– 2490.
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JCT 09-414