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This chapter investigates three ionic liquids (ILs), namely butyl pyridinium acetate ([BPy][AC]), butyl pyridinium benzoate ([BPy][BZ]), and butyl pyridinium propionate ([BPy][PR]), applied as potential absorbents for acid gases (hydrogen sulfide and carbon dioxide) in natural gas. The molecular dynamics (MD) simulation results indicate that the ILs have a relatively low dynamic and compact structure, with high viscosity in their pure state. Consistent with the findings of other researchers, the qualitative analysis of the simulation data for the mixture of an IL with acid and methane gases suggests that the dynamics of the IL enhances in the presence of these gases. The radial distribution functions reveal strong interactions and structural compatibility between the ILs and hydrogen sulfide molecules, indicating their suitability for hydrogen sulfide absorption. The amount of carbon dioxide gas absorbed by these ILs was calculated to be in the range of 0.08–0.11, while the absorption of hydrogen sulfide gas ranged from 0.12 to 0.18. [BPy][PR] IL exhibited the highest percentage of absorption for carbon dioxide (0.1083) and hydrogen sulfide (0.177). Furthermore, a comparison of the interactions between acidic gases and [BPy][PR] with the results of methyldiethanolamine (MDEA) clearly demonstrates the superior physical absorption of these gases by [BPy][PR].
Faculty of Science, Department of Chemistry, Ferdowsi University of Mashhad, Mashhad, Iran
*Address all correspondence to: moosavibaigi@um.ac.ir
1. Introduction
Destruction of the environment has become one of the main obstacles to the further progress of society [1]. For example, large amounts of unwanted gases produced by various industries contribute to global warming, acid rain, and air pollution, posing serious threats to the human environment. Natural gas is a fossil fuel known for having fewer environmental effects compared to other fossil fuels. It primarily consists of methane, ethane, and trace amounts of propane and butane. Additionally, natural gas contains compounds like hydrogen sulfide and carbon dioxide, which must be removed due to their toxicity, corrosiveness, and ability to reduce the heat value of the fuel. Gas sweetening is a crucial process for purifying industrial gases, and various gas separation technologies, such as physical and chemical solvents, pressure swing adsorption, and membranes, are used in this field. However, the complexity of gas components and varying conditions often lead to high energy consumption, costs, and secondary pollution in most technologies. Two of the harmful gases that are detrimental to the environment and are classified as acid gases are carbon dioxide and hydrogen sulfide [2]. The most common method used in industries to absorb these gases is the use of aqueous amine solutions, such as monoethanolamine (MEA) [3, 4, 5, 6, 7], diethanolamine, methyldiethanolamine (MDEA) [8], and MDEA-piperazine mixtures [9]. However, issues such as corrosion, solvent degradation, high energy consumption, and pollution caused by volatility limit the effectiveness of these methods [10]. In recent years, a new type of solvents, known as green solvents or ionic liquids (ILs), have been introduced due to their exceptional properties for separating acid gases from natural gas [2, 11].
The process of removing unwanted compounds, such as hydrogen sulfide and carbon dioxide, from natural gas to reduce environmental and corrosion risks and increase the heat value of the fuel is known as natural gas sweetening. It is essential to separate these unwanted compounds to the fullest extent before using natural gas. The acceptable limits for hydrogen sulfide and carbon dioxide in natural gas are 4 ppm [12] and 2.5% [13], respectively.
In the oil and gas industry, amino solvents are used to separate acid gases from methane. These solvents have high absorption power due to chemical absorption but also drawbacks such as high vapor pressure, corrosion, and high energy consumption during solvent reclamation. For this reason, researchers are seeking solvents that can serve as alternatives to traditional solvents. One of the proposed alternatives are ILs, which have the ability to physically absorb significant amounts of acidic gases and can be considered as substitutes for amines. Compared to traditional solvents, ILs offer desirable properties such as high polarity; low volatility and inflammability; high thermal, chemical, and electrochemical stability; lower energy consumption; reduced absorbent and water use; the ability to dissolve many organic, inorganic, and organometallic compounds; modulated miscibility with various organic solvents; and adjustable structures [14].
Organic molten salts consisting of cations and anions with a melting point lower than 100°C are called ILs that are highly versatile as “design solvents.” Their valuable physical and chemical properties include a low melting point (less than 100°C), very low vapor pressure, high thermal stability, high electrical conductivity, low surface tension, and adjustable viscosity. In separation processes, important parameters include the melting point, viscosity, vapor pressure, thermal stability, and chemical stability [15]. If the melting point of an IL is around room temperature, it is referred to as a room temperature ionic liquid (RTIL). These compounds typically consist of a large asymmetric organic cation paired with an organic or inorganic anion [16]. The significant interest in ILs arises from their ability to be designed as non-volatile substances. Furthermore, the wide range of available cations and anions, and their ability to be combined in various ways, offers great flexibility in processes. ILs can contain both polar and non-polar components, making them ideal for dissolving substances with different polarities [2]. Various cations and anions can be used to create ILs resulting in a diverse array of ILs tailored for specific applications or with enhanced physicochemical properties.
Common cations include imidazolium, pyridinium, pyrrolidinium, ammonium, phosphonium, and sulfonium [17, 18], and common anions include bis(trifluoromethylsulfonyl) imide, trifluoromethyl sulfate, dicyanamide, tetrafluoroborate, and hexafluorophosphate. Additionally, there are simple anions like chloride, bromide, iodide, nitrate, perchlorate, formate, and acetate, which may have less stability or may not be liquid at room temperature.
Some physical properties of ILs, such as viscosity, boiling point, and solubility in water, can be adjusted based on the ion pair selection or through the introduction of special functional groups into the structure of the cation or anion. In general, the most important advantages of ILs are as follows:
These compounds have low vapor pressure, making them non-volatile and environmentally friendly compared to common organic solvents.
ILs are more conductive than organic compounds and can easily dissolve various enzymes.
These green solvents have high thermal stability and do not decompose at high temperatures, making product separation easy.
ILs have low flammability, making them safe for chemical reactions.
ILs can dissolve various compounds due to their structural diversity, and their thermodynamic properties can be improved by changing the structure.
These compounds have high density and polarity, remaining liquid over a wide range of temperatures.
ILs can be categorized into two important groups: conventional ILs and task-specific ILs [3]. The key difference is that task-specific ILs have chemical absorption capabilities, able to absorb acid gases with low concentrations, while conventional ILs primarily perform physical absorption. When the concentration of acid gas in the input feed is low, conventional ILs do not have a high absorption capacity [19]. Research has shown that functional ILs are capable of both physical and chemical absorption. In some functional ILs, certain atoms, groups, or structures of cations exist as functional groups to absorb acid gases. For example, 1-(2-diethylaminoethyl)-3-methylimidazolium hexafluorophosphate and 1-(2-diethylaminoethyl)-1-methylpyrrolidinium hexafluorophosphate are functional ILs used for the absorption of sulfur dioxide due to the amine group in the cation structure [20].
Although imidazolium-based ILs have disadvantages such as high manufacturing cost, viscosity, and toxicity [21], the low cost [3], low toxicity, high thermal stability, and biodegradability of ILs based on pyridinium cation make them favorable.
The ability to create a wide variety of ILs has led to the use of molecular dynamics (MD) simulation and computational chemistry methods to calculate the electronic, thermodynamic, and phase equilibrium properties of ILs [22, 23, 24, 25, 26, 27, 28, 29, 30, 31, 32, 33, 34, 35].
Additionally, the solubility of a mixture of CO2 and H2S in aqueous solutions of amines and ILs, such as diisopropanolamine (DIPA) blended with [bmim][acetate], has been measured [2]. The gas solubility is influenced by the concentration of DIPA, with the addition of IL leading to increased solubility. Pyridinium-based ILs have excellent properties including low cost, viscosity, and toxicity; high thermal stability and biodegradability; and relatively high CO2 solubility according to Frias et al. [11].
Sazanova et al. [1] focused on amphiphilic bis(2-ethylhexyl) sulfosuccinate anion and imidazolium, pyridinium, and pyrrolidinium cations to determine H2S and CO2 solubility in all prepared ILs. They demonstrated that van der Waals (vdW) interactions, in addition to electrostatic and hydrogen bonding, play a significant role in acid gas removal by these ILs, leading to the conclusion of a physical absorption mechanism. The protic nature of the cation results in a notable reduction in gas solubility, and the sorption capacity is influenced by the aprotic nature of the cation.
Wang et al. [36, 37] synthesized a series of pyridinium-based ILs and investigated their properties including density, viscosity, thermal decomposition, and gas separation performance for H2S, CO2, and SO2. The results showed that the solubility of gases in ILs increases significantly with pressure and decreases with temperature. Additionally, as the length of the alkyl chain attached to the cation increases, the solubility rises due to the intensity in vdW force between the ions and the acid gas. Among the studied ILs with the same cations, the [C4Py][SCN] exhibited the highest selectivity about 1.5–4 times larger than imidazolium-based ILs. These ILs showed a greater molar fraction of hydrogen sulfide gas absorbed compared to CO2. Furthermore, pyridinium-based ILs demonstrated stable absorption performance after five consecutive cycles, indicating their potential as effective absorbents for gas separation applications.
The direct emission of ammonia gas by the chemical industry severely pollutes the environment. Traditional methods for removing this gas include washing with water or acid. Nevertheless, absorption by ILs offers advantages such as no solvent wastage and easy recovery of NH3 [38, 39, 40, 41]. Yokozeki et al. [42, 43] introduced eight types of ILs for absorbing ammonia gas at different temperature and pressure ranges, suggesting a complex molecular interaction between ammonia gas and ILs. Yuan et al. [40] stated that NH3 gas interacts chemically and physically with dual-functionalized pyridinium-based ILs containing acidic protons and hydroxyl groups. Further studies on six-membered N-heterocyclic cations in protic ILs (PILs) confirmed thermal stability during NH3 absorption with high capacities and selectivity [41]. Pyridinium- and piperidinium-based PILs were identified as acceptable candidates for NH3 absorption due to their better biodegradability as well as lower cost compared to imidazolium-based ILs, with the acid-base interaction and hydrogen bonding [44] playing a crucial role in NH3 interaction with IL.
Sulfur dioxide is another acid gas present in exhaust gases from power plants. Various ILs have been designed and synthesized as efficient absorbents for the absorption and separation of SO2, offering opportunities for new separation processes with easy recovery, high capacity, and low enthalpy of absorption [20]. Wang et al. [45] have shown that small-sized ILs with moderate basicity of anions are more preferential for SO2 absorption.
Zeng et al. [36, 46, 47] investigated the effect of cation alkyl chain length and anion type on the solubility of sulfur gas in ILs, finding that anions play a significant role in the absorption of sulfur dioxide gas. Functionalized ILs with ether and nitrile groups exhibited high absorption capacity due to strong physical interactions with sulfur dioxide gas compared to conventional ILs [47, 48].
Shang et al. [10] reported that the chemical and physical absorption of SO2 gas affects the viscosity of pyridinium-based ILs, with chemically absorbed gas leading to an increase in viscosity. While most studies in the literature have focused on imidazolium and ammonium cations, investigations on other forms of cations are needed. Pyridinium-based ILs have shown promise for acidic gas capture, with ongoing research focusing on natural gas sweetening and selectivity of these ILs paired with carboxylate anions. The study aims to investigate the structures and properties of ILs and their efficiency in the natural gas sweetening process, highlighting their high absorption capacity and thermal stability, low vapor pressure, and adjustable structures that make them ideal for absorbing and separating acid gases.
In this study, MD simulations were conducted to investigate the removal of acid gases, including hydrogen sulfide and carbon dioxide from natural gas using ILs with butylpyridinium cation ([BPy+]) and acetate ([AC−]), benzoate ([BZ−]), and propionate ([PR−]) anions at 325 K and 1 atm. The temperature and pressure of the systems were maintained constant using the Nose-Hoover thermostat and barostat, respectively. The simulations were carried out using DL_POLY software version 2.17 [49], and the results were analyzed using VMD software version 1.9.1 [50]. MD simulations were performed for various systems, including ILs in a liquid and pure state, IL-carbon dioxide, IL-hydrogen sulfide, IL-CO2-H2S, and IL-CO2-H2S-CH4 mixtures. The simulation process consisted of three stages: preparation of input files, conducting the simulation, and extracting and analyzing the results.
To initiate the simulation, the initial configuration of the system was selected, and an initial cubic simulation box with an edge of 150 Å containing 125 ion pairs of IL and a specific number of gas molecules (90 molecules of methane, 4 molecules of H2S, and 6 molecules of carbon dioxide) was prepared. Information about the force fields for the cation [BPy]+ [51]; anions [AC]− [52], [BZ]− [53], and [PR]− [54]; and the target gases CO2 [53], H2S [55], and CH4 [39, 56] was obtained from reliable scientific sources. Simulations were conducted with a time step of 1 fs. The system first equilibrated in NPT ensemble for 1 ns, followed by production runs in the NVT ensemble for 1 ns.
After preparing the input files, simulations were conducted with the NPT statistical ensemble, where the system volume was equilibrated and then fluctuated around a constant average value. Subsequently, simulations were continued in the NVT ensemble to collect results. Once the simulations were completed, the results were analyzed to calculate the structural and dynamic properties of the system. Figure 1 shows the structure of each anion and cation studied in the present study.
Figure 1.
The structure of studied anions (from the left [AC]− (first), [BZ]− (second), and [PR]− (third)) and cation [BPy]+ (right).
After equilibration, the ratio of fluctuation to the root mean square (rms) of energy and volume quantities for pure ILs was computed. If the fluctuation in rms is less than 0.01, the system is stable, and the simulation time is sufficient. For [BPy][AC], [BPy][BZ], and [BPy][PR] ILs, the ratios of fluctuation to rms in energy values are 0.183, 0.178, and 0.247 J, respectively. The volume fluctuation ratios are 0.0045, 0.00153, and 0.00321 Å3, respectively. These results suggest that 1 ns of MD simulation is adequate.
Accurate thermophysical properties, like intermolecular energy, are essential for fluid modeling. Some properties are challenging to calculate in labs but can be easily evaluated through MD simulations [57]. Table 1 lists of some of these properties at 325 K and 1 atm.
IL
Intermolecular energy (Uint)
Electrostatic energy (Uele)
vdW energy (UvdW)
[BPy][AC]
−68,230
−62,516
−5714
[BPy][BZ]
−75,677
−67,788
−7889
[BPy][PR]
−73,810
−68,019
−5791
Table 1.
Comparison of intermolecular, electrostatic, and van der Waals (vdW) energies in terms of kJ/mol in each of the pure [BPy][AC], [BPy][BZ], and [BPy][PR] ILs.
Comparing intermolecular energy in three simulated pure ILs (Table 1) reveals that [BPy][BZ] has the strongest Uint. Electrostatic energy strongly influences the system’s energy, explaining ILs’ unique properties like low volatility and vapor pressure due to high electrostatic energy. [BPy][AC] has the lowest vdW energy, attributed to the acetate anion’s small size. Larger particles have considerable vdW energy, explaining why [BZ]− has the highest UvdW.
Mean square displacement (MSD) increases over time provide insight into species dynamics. MSD changes with time for anions and cations in each pure system at 325 K were calculated. The graphs in Figure 2 show almost linear behavior, with self-diffusion coefficients (D) calculated from the linear part’s slope. D values range from 100 to 400 picoseconds and are listed in Table 2.
Figure 2.
MSD variation with simulation time for studied pure ILs.
IL
ρSim (g/cm3)
D+ (m2/s)
D− (m2/s)
ηIL (cP)
[BPy][AC]
0.996
1.100 × 10−11
1.565 × 10−11
136
[BPy][BZ]
0.840
1.013 × 10−11
1.154 × 10−11
167
[BPy][PR]
0.968
1.008 × 10−11
1.351 × 10−11
153
Table 2.
Density, D of cations (D+), D of anions (D−), and viscosity of pure ILs at 325 K.
Self-diffusion coefficients for anions and cations are around 10−11 m2/s, consistent with experimental values [58, 59, 60]. Anions have faster dynamics and D values than cations due to their smaller sizes and mass, reflecting faster mobility [58]. Density comparisons validate the simulation, with [BPy][AC] showing acceptable agreement [57]. Viscosity was calculated using Stokes-Einstein Eq. (1) [61]:
DIL=DH2OηH2OηILE1
The density and viscosity values are given in Table 2.
Radial distribution functions (RDFs) provide detailed insight into the IL structure. RDFs of anions around the cation’s HA1 atom show strong hydrogen bonding (HB) (Figure 3). Benzoate anion has the most interaction with [BPy]+ cation, forming a compact structure with strong intermolecular forces. Center of mass (COM) RDFs for cation-cation, anion-cation, and anion-anion pairs (Figure 4) show different interactions. Anion-cation interactions are strong, with approximately six anions per cation in each IL, consistent with prior research [57].
Figure 3.
RDF of anion oxygen atom around cation (a) HA1 atom (shown in legend), (b) CA1 atom in each pure system of ionic liquids [BPy][AC] (blue), [BPy][BZ] (red), and [BPy][PR] (green).
Figure 4.
Diagram of RDFs at 325 K for cation-cation, cation-anion, and anion-anion of pure studied ILs.
3.2 Mixed systems
After studying the dynamic and structural properties of pure ionic liquids [BPy][AC], [BPy][BZ], and [BPy][PR], the systems of these ILs with CO2 and H2S were separately examined. The simulation systems reviewed in this section include a mixture of IL with CO2 gas, a mixture of IL with H2S gas, a mixture of IL with a combination of CO2 and H2S gases, and finally a mixture of IL with a combination of CO2, H2S, and CH4 gases. The dynamic and structural properties of these systems are presented and analyzed.
Figure 5 compares the MSD for [BPy]+ cation and [AC]−, [BZ]−, and [PR]− anions in pure and mixed systems with IL + CO2.
Figure 5.
Comparison of the MSD for cation and anions between the pure IL and mixed with CO2 at 325 K.
As shown in Figure 5, the dynamics of the cation and anions in the mixed systems increase significantly compared to the pure state, indicating a strong interaction between ions and carbon dioxide gas. This finding aligns with previous research by Morganti et al. [62] who studied the solubility of SO2 and CO2 gases in ammonium-based ILs and found that gas solubility increased the D of IL species.
Figure 6 illustrates the comparison of the MSD of cations and anions in all three mixed systems. The dynamics of anions and cations in the [BPy][AC] system with CO2 gas is higher than in the other two systems, likely due to the small size of acetate anion and the cluster movement of ILs.
Figure 6.
Comparison between MSD of cation and anion in each mixed system with CO2 and pure ILs at 325 K.
Figure 7 displays the MSD of CO2 in mixtures with CO2 gas and [BPy][AC], [BPy][BZ], and [BPy][PR] ILs.
Figure 7.
Comparison between MSD of CO2 in the target ILs [BPy][AC], [BPy][BZ], and [BPy][PR].
The dynamics of CO2 molecules in the [BPy][AC] + CO2 system is significantly different from the other systems, possibly due to the low absorption rate of CO2 molecules in [BPy][AC] IL, resulting in a separate gas phase within the simulation box; see Figure 8 for more details.
Figure 8.
Graphical representation of the CO2 absorption in (a) [BPy][AC], (b) [BPy][BZ], and (c) [BPy][PR]. Red, blue, and dark blue colors display oxygen, carbon, and nitrogen atoms, respectively.
The diffusion coefficients of each species were calculated using Einstein’s relation [63], with the values reported in Table 3.
IL + CO2
D+ (m2/s)
D− (m2/s)
DCO2 (m2/s)
[BPy][AC]
2.030 × 10−11
2.778 × 10−11
1.020 × 10−8
[BPy][BZ]
1.095 × 10−11
1.136 × 10−11
0.475 × 10−8
[BPy][PR]
1.721 × 10−11
2.024 × 10−11
0.485 × 10−8
Table 3.
Cation, anion, and CO2 gas diffusion coefficient in the mixture containing CO2 gas and each IL at 325 K.
According to Table 3, the self-diffusion coefficient of CO2 gas dissolved in IL is larger than the corresponding values of ions in good agreement with Morganti and the co-workers’ research [62]. The DCO2 in [BPy][BZ] IL is the lowest due to the large size of benzoate anion and strong intermolecular, electrostatic, and vdW forces in this system (see Table 1 for more details), leading to a compact structure with high viscosity. The role of anion in gas absorption is examined through the RDF of the COM of each species. Figure 9 compares the RDF of cation-CO2 and anion-CO2 to explore the strength of interaction with the absorbed gas.
Figure 9.
The RDF between CO2 and IL ion pairs at 325 K in [BPy][AC] + CO2, [BPy][BZ] + CO2, and [BPy][PR] + CO2 systems.
The RDF analysis shows that anion interacts more with the absorbed gas than the cation, with [PR]− anion displaying the strongest interaction with CO2 gas among the three ILs. From the other side of view, the terminal atoms related to the anion chain also show a strong interaction with CO2 molecules, that is, stronger than that of cations. This observation is in agreement with the literature [64, 65, 66].
According to the observations of Ren et al. [67], if the acidic pKa of the anion corresponding to the acids that make up the ILs is greater than 6.36, the dissolution of CO2 gas in this type of ILs takes place in the form of chemical absorption. Otherwise, CO2 absorption is done physically. In fact, the higher the pKa, the greater the alkalinity of ILs, the greater the adsorption capacity while the selectivity decreases. The nature of the anion mother is acidic. The pKa value for acetic acid is equal to 4.75, in the case of benzoic acid pKa = 4.202, and the value for propionic acid is 4.88 [68]. As can be seen, the pKa of these acids is less than 6.36. Therefore, it can be expected that the interactions of these ILs with carbon dioxide gas are weak, which is indicated by the presence of broad peaks with low intensity in Figure 9. It can also be expected that the [PR]− anion has a strongest interaction with CO2 gas in comparison with [AC]− and [BZ]− anions. The existence of a peak with high intensity compared to the other two anions confirms this result.
Another expectation is that the intensity of interaction between CO2 and [BZ]− anion is the lowest since it has the least alkaline property. However, it is observed that [BZ]− anion-CO2 intensity is almost similar to [AC]− anion-CO2 interaction. This is due to the large size of [BZ]− anion. In fact, physical absorption is influenced by vdW and electrostatic forces. The larger the size of the ions, the stronger the vdW force between the species [57].
The coordination number of CO2 gas around anion and cation is also reported in Table 4 showing CO2 molecules are most distributed around [BPy][PR] IL, indicating high absorption due to its alkalinity.
IL + CO2
N+
N−
[BPy][AC]
0.0404
0.0415
[BPy][BZ]
0.0423
0.0552
[BPy][PR]
0.0432
0.0645
Table 4.
Coordination number of CO2 around the cation (N+) and each anion (N−) in the mixed system including CO2 gas with ILs [BPy][AC], [BPy][BZ], and [BPy][PR].
Figure 10 compares the MSD of pure and mixed systems with H2S gas and ILs [BPy][AC], [BPy][BZ], and [BPy][PR].
Figure 10.
Comparison of the cation, anion, and H2S MSD in pure and mixed states with H2S gas in [BPy][AC], [BPy][BZ], and [BPy][PR] ILs.
The MSD analysis of H2S gas absorption in ILs presents increasing dynamics of ions compared to pure IL. [BPy][AC] IL shows the highest dynamics for the small size of the acetate anion. The coordination number of H2S molecules around each IL is reported in Table 5. The table reveals that hydrogen sulfide is most absorbed by [BPy][PR] IL, attributed to its higher alkalinity. The RDF analysis of H2S absorption in ILs (Figure 11) shows stronger interaction with the cation than the anion, with [PR]− anion exhibiting similar intensity to the cation due to its high alkalinity.
IL + H2S
D+ (m2/s)
D− (m2/s)
DH2S (m2/s)
N+
N−
[BPy][AC]
1.603 × 10−11
2.320 × 10−11
0.413 × 10−8
0.055
0.062
[BPy][BZ]
1.187 × 10−11
1.243 × 10−11
0.120 × 10−8
0.052
0.083
[BPy][PR]
1.270 × 10−11
1.665 × 10−11
0.353 × 10−8
0.056
0.121
Table 5.
Cation, anion, and H2S gas diffusion coefficient values and coordination number of H2S around [BPy]+, [AC]−, [BZ]−, and [PR]− in the mixture of IL and H2S at 325 K.
Figure 11.
The RDF of H2S around cation and anion of [BPy][AC], [BPy][BZ], and [BPy][PR].
Radial distribution function (RDF) investigation shows that H2S gas accumulates around the cation alkyl chain and anion, with strong interactions observed at the terminal atoms of this chain. The results of other researchers confirm this finding [69]. The polarity of hydrogen sulfide is lower than that of water, resulting in higher solubility in organic solvents compared to aqueous solvents [69]. This can be attributed to the less polar nature of alkyl chain if compared to water, leading to better absorption in organic solvents. Therefore, hydrogen sulfide molecules are absorbed in the organic compartment of cation and anion.
As depicted in Figure 11, the interaction between the cation and H2S is stronger than the anion and hydrogen sulfide. This is due to the larger size of the cation relative to the anion. According to the literature [69, 70], the large molecular volume is responsible for gas solubility. Therefore, as the [BPy]+ cation is larger than the [AC]− and [BZ]− anions, H2S molecules are more concentrated around the cation, resulting in stronger interactions with the cation than with the anion. Analysis of the RDF related to the mixed system containing hydrogen sulfide and [BPy][PR] IL reveals that despite the cation’s larger size compared to the [PR]− anion, the intensity of the molecular interaction of hydrogen sulfide with [PR]− anion is nearly equivalent to its interaction with the cation. This could be attributed to the high alkalinity of the [PR]− anion.
Careful examination of the graphical representation of each IL + H2S system (Figure 12) reveals that hydrogen sulfide molecules are more readily absorbed on the surface of the IL. This results in a reduction of ion-ion interaction due to the interaction of ions with the gas molecules, ultimately leading to increased mobility and movement of ions within the mass. The fast diffusion of ions in mixed systems, as opposed to pure ones, is attributed to the strong interaction between ions and hydrogen sulfide. Specifically, H2S penetrates more into the [BPy][BZ] IL compared to the other two systems, resulting in these molecules entering the bulk of the IL, decreasing mobility more than the other two systems. In summary, the larger the size of the ion, the stronger the interaction with the absorbed gas, highlighting the importance of anion properties in gas absorption in ILs.
Figure 12.
Graphical representation of the H2S absorption in (a) [Bpy][AC], (b) [Bpy][BZ], and (c) [Bpy][PR]. The colors red, light blue, dark blue, and yellow represent oxygen, carbon, nitrogen, and sulfur atoms, respectively.
3.3 Mixed systems and selectivity (mixed gases)
The investigation of the selectivity and structural properties of the mixed system containing H2S and CO2 gases with each IL [Bpy][AC], [Bpy][BZ], and [Bpy][PR] was conducted. Selective separation of H2S and CO2 gases is crucial in natural gas sweetening. Therefore, it is essential to use solvents with high selectivity for H2S gas over CO2 in natural gas sweetening. As mentioned above, the larger the size of the ions and the greater the alkalinity of ILs, the more acidic gases (H2S and CO2) are absorbed, but the selectivity of H2S decreases. Therefore, ILs with small ions and low alkalinity may exhibit high selectivity.
In Figure 13, the RDF diagram of H2S and CO2 molecules around cations and anions in each of the three-component mixed systems of H2S + CO2 + [Bpy][AC], H2S + CO2 + [Bpy][BZ], and H2S + CO2 + [Bpy][PR] is depicted. The intensity of H2S…cation interaction is greater than the intensity of H2S…anion interaction. The absorption of hydrogen sulfide gas is more influenced by the size of the ions, with larger ions absorbing more hydrogen sulfide molecules. Increasing the length of the cation chain or using larger anions can enhance the absorption of H2S gas. However, high alkalinity is also an important factor in the absorption of these molecules [69], as seen in the strong interaction of propionate anion with hydrogen sulfide (Figure 11).
Figure 13.
RDF related to the S atom of H2S and the O atom of CO2 (OC) around [Bpy]+, [AC]−, [BZ]−, and [PR]− in each of the mixed systems containing both gases and IL.
Figure 13 also indicates that the cation-hydrogen sulfide interaction is stronger than the cation-carbon dioxide interaction. Enlarging the length of the alkyl chain improves the absorption of both gases. The interaction of anion-H2S is stronger than that of anion-CO2, with increased alkalinity leading to decreased selectivity [62]. The benzoate anion shows the highest selectivity due to its lower alkalinity compared to the other anions.
Comparing RDFs related to hydrogen sulfide and carbon dioxide around cations and anions in each of the mixed systems with ILs [BPy][AC], [BPy][BZ], and [BPy][PR] reveals consistency in each system. The intensity of the interaction between benzoate anion and hydrogen sulfide is stronger than that of acetate anion, attributed to the larger size of benzoate anion. The alkalinity of the anion influences the absorption of CO2, with higher alkalinity leading to greater interaction and absorption.
Figure 14 presents the selectivity of H2S gas over CO2 and methane in each of the mixed systems.
Figure 14.
The qualitative comparison of the selectivity of H2S gas over CO2 and CH4 in each of the mixed systems including [BPy][AC] + H2S + CO2 + CH4, [BPy][BZ] + H2S + CO2 + CH4, and [BPy][PR] + H2S + CO2 + CH4.
The interactions of cation and anion with H2S molecules are stronger than with other gases, CO2 and CH4, reflecting the higher solubility of H2S in ILs. The increase in anion alkalinity enhances interaction with acid gases, that is, the strongest in the case of [BPy][PR]. However, increasing alkalinity can reduce selectivity [69].
Figure 15 depicts the MD process in a comprehensive global view. The first step involves simulating the pure IL. Next, CO2 acidic gas is introduced into the IL, and the simulation begins. It is evident from the demonstration that CO2 is fully dissolved in the [BPy][PR] IL. In the third step, a mixture of CO2 and H2S acidic gases is introduced into the solvent for MD simulation to determine if the gases are absorbed. The final step (4) includes not only the IL but also a mixture of CO2, H2S, and CH4 gases. The illustration indicates that the CO2 gas has successfully dissolved completely.
Figure 15.
The MD process for gas absorption by [BPy][PR] as a typical sample of the present study.
3.4 Comparison of the solubility of acid gases in [BPy][PR] and MDEA
In Figure 16, the solubility of H2S and CO2 gases in [BPy][PR] IL and MDEA amine solvent is compared. The results show that [BPy][PR] IL is the most effective solvent among the studied ILs for acid gas absorption, leading to its selection for comparison with MDEA. The MD simulation details for MDEA are consistent with those of the ILs. Force field data were taken from the CHARAM force field [71]. While MD simulation cannot explore chemical absorption, it does determine the physical absorption of H2S and CO2 gases, as depicted in Figure 16. The intensity of H2S interaction with the IL is notably higher than with amine solvent. The peak in the RDF of IL with H2S occurs at a shorter distance, with greater intensity and a narrower width, indicating a strong interaction with H2S.
Figure 16.
Comparing the solubility of acid gases in MDEA and [BPy][PR].
The dynamics and structure of ILs [BPy][AC], [BPy][BZ], and [BPy][PR] were studied in their pure state and mixed with H2S, CO2, and CH4 gases. To analyze the dynamics of the systems, the MSD was utilized, while the RDF was used to examine the structure. The MSD quantity for the COM of the cation and each of the anions at 325 K was calculated to investigate the dynamics of the species. In all systems, the anion exhibited faster dynamics than the cation, which can be attributed to the lower molecular mass of the cation. The self-diffusion coefficient values for the cation and anion of pure ILs were calculated in each system at the specified temperature with values around 10−11 m2/s for ion pairs of ILs. Results from the MSD analysis indicated that the presence of acid gases significantly increased the dynamics of ions due to interactions between ions and gas molecules. The MSD of H2S molecules was notably lower than that of CO2 molecules, likely due to the strong interaction of ions with hydrogen sulfide.
Analysis of the RDF for pure ILs revealed that interactions between ions depended on their size, with larger ions leading to a more compact structure and higher viscosity. When examining RDFs in mixed systems, it was observed that larger cation and anion sizes resulted in increased absorption of H2S, CO2, and CH4 gases in ILs, leading to decreased selectivity. The solubility of poorly soluble gases in ILs, such as methane, was found to be more dependent on the volume and size of the cation and anion in the IL. On the other hand, the solubility of H2S and CO2 gases depended on electrostatic interactions in addition to the volume and size of ions. Higher alkalinity in ILs led to increased solubility of acid gases (H2S and CO2) in the liquids. Due to the higher acidity and polarity of H2S molecules compared to CO2, the solubility of H2S gas in IL was higher than that of CO2.
This research project was financially supported by Ferdowsi University of Mashhad, Iran (grant No. 3/48448). The computations were partly carried out in the High-Performance Computing (HPC) Center at Ferdowsi University of Mashhad. The author would like to appreciate the HPC cooperation.
1.Sazanova TS, Akhmetshina AI, Petukhov AN, Vorotyntsev AV, Suvorov SS, Barysheva AV, et al. The cation effect on the free volume and the solubility of H2S and CO2 in ionic liquids based on Bis(2-ethylhexyl) sulfosuccinate anion. Membranes. 2023;13(2):238
2.Afsharpour A, Haghtalab A. Simultaneous measurement absorption of CO2 and H2S mixture into aqueous solutions containing diisopropanolamine blended with 1-butyl-3-methylimidazolium acetate ionic liquid. International Journal of Greenhouse Gas Control. 2017;58:71-80
3.Hussain SMS, Adewunmi AA, Alade OS, Murtaza M, Mahboob A, Khan HJ, et al. A review of ionic liquids: Recent synthetic advances and oilfield applications. Journal of the Taiwan Institute of Chemical Engineers. 2023;153:105195
4.Budzianowski WM. Single solvents, solvent blends, and advanced solvent systems in CO2 capture by absorption: A review. International Journal of Global Warming. 2015;7(2):184-225
5.Dutcher B, Fan M, Russell AG. Amine-based CO2 capture technology development from the beginning of 2013 - A review. ACS Applied Materials & Interfaces. 2015;7(4):2137-2148
6.Yamada H. Amine-based capture of CO2 for utilization and storage. Polymer Journal. 2021;53(1):93-102
7.Julio AAV, Castro-Amoedo R, Maréchal F, González AM, Escobar Palacio JC. Exergy and economic analysis of the trade-off for design of post-combustion CO2 capture plant by chemical absorption with MEA. Energy. 2023;280:128004
8.Kopitha K, Elakneswaran Y, Kitagaki R, Saito R, Tsujino M, Nishida A, et al. N-methyldiethanolamine (MDEA) as an effective CO2 absorbent for direct air capture (DAC) in cement-based materials. Chemical Engineering Journal. 2023;475:146067
9.Li F, Hemmati A, Rashidi H. Industrial CO2 absorption into methyldiethanolamine/piperazine in place of monoethanolamine in the absorption column. Process Safety and Environmental Protection. 2020;142:83-91
10.Shang D, Liu X, Bai L, Zeng S, Xu Q , Gao H, et al. Ionic liquids in gas separation processing. Current Opinion in Green and Sustainable Chemistry. 2017;5:74-81
11.Frias RM, Follegatti-Romero LM, Follegatti-Romero LA. Modelling mixtures of ammonium and pyridinium-based ionic liquids and carbon dioxide with the Cubic Plus Association Equation of state. Fluid Phase Equilibria. 2023;574:113910
12.Farzaneh A, Saghatoleslami N, Feyzi Y. Transient H2S content rise in the effluent of a natural gas treating unit: Role of COS hydrolysis and heat stable amine salts. Process Safety and Environmental Protection. 2021;153:84-93
13.Mazzetti MJ, Skagestad R, Mathisen A, Eldrup NH. CO2 from natural gas sweetening to kick-start EOR in the north sea. Energy Procedia. 2014;63:7280-7289
14.Ochędzan-Siodłak W, Dziubek K, Siodłak D. Densities and viscosities of imidazolium and pyridinium chloroaluminate ionic liquids. Journal of Molecular Liquids. 2013;177:85-93
15.Singh SK, Savoy AW. Ionic liquids synthesis and applications: An overview. Journal of Molecular Liquids. 2020;297:112038
16.Hallett JP, Welton T. Room-temperature ionic liquids: Solvents for synthesis and catalysis. 2. Chemical Reviews. 2011;111(5):3508-3576
17.Moosavi F. The structure of supported ionic liquids at the interface. In: Kadokawa J-i, editor. Ionic Liquids-New Aspects for the Future. London, UK: IntechOpen; 2013
18.Poole CF, Poole SK. Extraction of organic compounds with room temperature ionic liquids. Journal of Chromatography A. 2010;1217(16):2268-2286
19.Babamohammadi S, Shamiri A, Aroua MK. A review of CO2 capture by absorption in ionic liquid-based solvents. Reviews in Chemical Engineering. 2015;31(4):383-412
20.Ren S, Hou Y, Zhang K, Wu W. Ionic liquids: Functionalization and absorption of SO2. Green Energy & Environment. 2018;3(3):179-190
21.Flieger J, Flieger M. Ionic liquids toxicity-benefits and threats. International Journal of Molecular Sciences. 2020;21(17):1-41
22.Siami H, Razmkhah M, Moosavi F. Cation functional group effect on SO2 absorption in amino acid ionic liquids. Frontiers in Chemistry. 2023;11:1-12
23.Siami H, Razmkhah M, Moosavi F. Does side chain group of anion affect absorption of SO2 in amino acid ionic liquid? Journal of Molecular Liquids. 2023;376:121479
24.Taji K, Moosavi F. Molecular simulation of CO2 capturing by dual functionalized phosphonium-based amino acid ionic liquids. Materials Today Communications. 2023;37:106945
25.Ghatee MH, Moosavi F, Zolghadr AR, Jahromi R. Critical-point temperature of ionic liquids from surface tension at liquid-vapor equilibrium and the correlation with the interaction energy. Industrial and Engineering Chemistry Research. 2010;49(24):12696-12701
26.Ghatee MH, Zare M, Moosavi F, Zolghadr AR. Temperature-dependent density and viscosity of the ionic liquids 1-alkyl-3-methylimidazolium iodides: Experiment and molecular dynamics simulation. Journal of Chemical and Engineering Data. 2010;55(9):3084-3088
27.Ghatee MH, Zare M, Zolghadr AR, Moosavi F. Temperature dependence of viscosity and relation with the surface tension of ionic liquids. Fluid Phase Equilibria. 2010;291(2):188-194
28.Ghatee MH, Moosavi F. Physisorption of hydrophobic and hydrophilic 1-alkyl-3-methylimidazolium ionic liquids on the graphenes. Journal of Physical Chemistry C. 2011;115(13):5626-5636
29.Ghatee MH, Zolghadr AR, Moosavi F, Ansari Y. Studies of structural, dynamical, and interfacial properties of 1-alkyl-3-methylimidazolium iodide ionic liquids by molecular dynamics simulation. Journal of Chemical Physics. 22012;136(12):124706(1-14)
30.Moosavi F, Razmkhah M. Structural view of hydrophobic ionic liquid on graphene: Comparing static and ab initio computer simulations. ECS Journal of Solid State Science and Technology. 2015;4(12):M77-M87
31.Sadeghi Moghadam B, Razmkhah M, Hamed Mosavian MT, Moosavi F. Molecular dynamics simulation of amino acid ionic liquids near a graphene electrode: Effects of alkyl side-chain length. Physical Chemistry Chemical Physics. 2016;18(48):33053-33067
32.Ebrahimi M, Moosavi F. The effects of temperature, alkyl chain length, and anion type on thermophysical properties of the imidazolium based amino acid ionic liquids. Journal of Molecular Liquids. 2018;250:121-130
33.Razmkhah M, Hamed Mosavian MT, Moosavi F. Structural analysis of an amino acid ionic liquid: Bulk and electrical double layer. Journal of Molecular Liquids. 2018;268:506-516
34.Razmkhah M, Hamed Mosavian MT, Moosavi F. What is the effect of polar and nonpolar side chain group on bulk and electrical double layer properties of amino acid ionic liquids? Electrochimica Acta. 2018;285:393-404
35.Nadimi H, Housaindokht MR, Moosavi F. The effect of anion on aggregation of amino acid ionic liquid: Atomistic simulation. Journal of Molecular Graphics and Modelling. 2020;101:107733
36.Wang J, Zeng S, Bai L, Gao H, Zhang X, Zhang S. Novel ether-functionalized pyridinium chloride ionic liquids for efficient SO2 capture. Industrial & Engineering Chemistry Research. 2014;53(43):16832-16839
37.Wang X, Zeng S, Wang J, Shang D, Zhang X, Liu J, et al. Selective separation of hydrogen sulfide with pyridinium-based ionic liquids. Industrial & Engineering Chemistry Research. 2018;57(4):1284-1293
38.Duan H, Liu F, Hussain S, Dong H, Zhang X, Cheng Z. Degradation processes of protic ionic liquids for NH3 separation. New Journal of Chemistry. 2023;47(23):11275-11284
39.Kapoor U, Banerjee A, Shah JK. Evaluation of the predictive capability of ionic liquid force fields for CH4, CO2, NH3, and SO2 phase equilibria. Fluid Phase Equilibria. 2019;492:161-173
40.Yuan L, Gao H, Jiang H, Zeng S, Li T, Ren B, et al. Experimental and thermodynamic analysis of NH3 absorption in dual-functionalized pyridinium-based ionic liquids. Journal of Molecular Liquids. 2021;323:114601
41.Li P, Shang D, Tu W, Zeng S, Nie Y, Bai L, et al. NH3 absorption performance and reversible absorption mechanisms of protic ionic liquids with six-membered N-heterocyclic cations. Separation and Purification Technology. 2020;248:117087
42.Yokozeki A, Shiflett MB. Ammonia solubilities in room-temperature ionic liquids. Industrial & Engineering Chemistry Research. 2007;46(5):1605-1610
43.Yokozeki A, Shiflett MB. Vapor–liquid equilibria of ammonia+ionic liquid mixtures. Applied Energy. 2007;84(12):1258-1273
44.Qiu R, Luo X, Yang L, Li J, Chen X, Peng C, et al. Regulated threshold pressure of reversibly sigmoidal NH3 absorption isotherm with ionic liquids. ACS Sustainable Chemistry & Engineering. 2020;8(3):1637-1643
45.Wang L, Zhang Y, Liu Y, Xie H, Xu Y, Wei J. SO2 absorption in pure ionic liquids: Solubility and functionalization. Journal of Hazardous Materials. 2020;392:122504
46.Zeng S, Gao H, Zhang X, Dong H, Zhang X, Zhang S. Efficient and reversible capture of SO2 by pyridinium-based ionic liquids. Chemical Engineering Journal. 2014;251:248-256
47.Zeng S, He H, Gao H, Zhang X, Wang J, Huang Y, et al. Improving SO2 capture by tuning functional groups on the cation of pyridinium-based ionic liquids. RSC Advances. 2015;5(4):2470-2478
48.Hong SY, Im J, Palgunadi J, Lee SD, Lee JS, Kim HS, et al. Ether-functionalized ionic liquids as highly efficient SO2 absorbents. Energy & Environmental Science. 2011;4(5):1802-1806
49.Smith W, Forester TR. DL_POLY_2.0: A general-purpose parallel molecular dynamics simulation package. Journal of Molecular Graphics. 1996;14(3):136-141
50.Humphrey W, Dalke A, Schulten K. VMD: Visual molecular dynamics. Journal of Molecular Graphics. 1996;14(1):33-38
51.Bandrés I, Alcalde R, Lafuente C, Atilhan M, Aparicio S. On the viscosity of pyridinium based ionic liquids: An experimental and computational study. The Journal of Physical Chemistry B. 2011;115(43):12499-12513
52.Shi W, Myers CR, Luebke DR, Steckel JA, Sorescu DC. Theoretical and experimental studies of CO2 and H2 separation using the 1-ethyl-3-methylimidazolium acetate ([emim][CH3COO]) ionic liquid. The Journal of Physical Chemistry B. 2012;116(1):283-295
53.Aparicio S, Atilhan M. A computational study on choline benzoate and choline salicylate ionic liquids in the pure state and after CO2 adsorption. The Journal of Physical Chemistry B. 2012;116(30):9171-9185
54.Aparicio S, Atilhan M. Nanoscopic vision on fuel dearomatization using ionic liquids: The case of piperazine-based fluids. Energy & Fuels. 2013;27(5):2515-2527
55.Kamath G, Lubna N, Potoff JJ. Effect of partial charge parametrization on the fluid phase behavior of hydrogen sulfide. The Journal of Chemical Physics. 2005;123(12):124505(1-7)
56.Martin MG, Siepmann JI. Transferable potentials for phase equilibria. 1. United-atom description of n-alkanes. The Journal of Physical Chemistry B. 1998;102(14):2569-2577
57.Chandran A, Prakash K, Senapati S. Structure and dynamics of acetate anion-based ionic liquids from molecular dynamics study. Chemical Physics. 2010;374(1):46-54
58.Cadena C, Zhao Q , Snurr RQ , Maginn EJ. Molecular modeling and experimental studies of the thermodynamic and transport properties of pyridinium-based ionic liquids. The Journal of Physical Chemistry B. 2006;110(6):2821-2832
59.Larive CK, Lin M, Piersma BJ, Carper WR. Diffusion ordered spectroscopy of room temperature chloroaluminate melts. The Journal of Physical Chemistry. 1995;99(33):12409-12412
60.Tsimpanogiannis IN, Moultos OA, Franco LFM, Spera MBM, Erdős M, Economou IG. Self-diffusion coefficient of bulk and confined water: A critical review of classical molecular simulation studies. Molecular Simulation. 2019;45(4-5):425-453
61.Morrow TI, Maginn EJ. Molecular dynamics study of the ionic liquid 1-n-butyl-3-methylimidazolium hexafluorophosphate. The Journal of Physical Chemistry B. 2002;106(49):12807-12813
62.Morganti JD, Hoher K, Ribeiro MCC, Ando RA, Siqueira LJA. Molecular dynamics simulations of acidic gases at interface of quaternary ammonium ionic liquids. The Journal of Physical Chemistry C. 2014;118(38):22012-22020
63.Kowsari MH, Alavi S, Ashrafizaadeh M, Najafi B. Molecular dynamics simulation of imidazolium-based ionic liquids. I. Dynamics and diffusion coefficient. The Journal of Chemical Physics. 2008;129(22):224508(1-13)
64.Anthony JL, Anderson JL, Maginn EJ, Brennecke JF. Anion effects on gas solubility in ionic liquids. The Journal of Physical Chemistry B. 2005;109(13):6366-6374
65.Bhargava BL, Balasubramanian S. Probing anion–carbon dioxide interactions in room temperature ionic liquids: Gas phase cluster calculations. Chemical Physics Letters. 2007;444(4):242-246
66.Seki T, Grunwaldt J-D, Baiker A. In situ attenuated total reflection infrared spectroscopy of imidazolium-based room-temperature ionic liquids under “supercritical” CO2. The Journal of Physical Chemistry B. 2009;113(1):114-122
67.Ren S, Hou Y, Tian S, Chen X, Wu W. Correction to “what are functional ionic liquids for the absorption of acidic gases?”. The Journal of Physical Chemistry B. 2013;117(17):5394-5395
68.Lide DR. CRC Handbook of Chemistry and Physics. Vol. 85. Boca Raton: CRC Press LLC; 2004
69.Wang L-y, Xu Y-l, Li Z-d, Wei Y-n, Wei J-p. CO2/CH4 and H2S/CO2 selectivity by ionic liquids in natural gas sweetening. Energy & Fuels. 2018;32(1):10-23
70.Bara JE, Gabriel CJ, Lessmann S, Carlisle TK, Finotello A, Gin DL, et al. Enhanced CO2 separation selectivity in oligo(ethylene glycol) functionalized room-temperature ionic liquids. Industrial & Engineering Chemistry Research. 2007;46(16):5380-5386
71.Brooks BR, Bruccoleri RE, Olafson BD, States DJ, Swaminathan S, Karplus M. CHARMM: A program for macromolecular energy, minimization, and dynamics calculations. Journal of Computational Chemistry. 1983;4(2):187-217
Written By
Fatemeh Moosavi
Submitted: 22 February 2024Reviewed: 11 April 2024Published: 31 July 2024
Open access peer-reviewed chapter - ONLINE FIRST
