Solubility of CO2 and H2S in the ionic liquid 1-ethyl-3-methylimidazolium trifluoromethanesulfonate

Accepted Manuscript Solubility of CO2 and H2S in the Ionic Liquid 1-Ethyl-3-methylimidazolium Trifluoromethanesulfonate Mohsen Nematpour, Amir H. Jalili, Cyrus Ghotbi, Davood Rashtchian PII:

S1875-5100(16)30051-8

DOI:

10.1016/j.jngse.2016.02.006

Reference:

JNGSE 1256

To appear in:

Journal of Natural Gas Science and Engineering

Received Date: 28 November 2015 Revised Date:

28 January 2016

Accepted Date: 13 February 2016

Please cite this article as: Nematpour, M., Jalili, A.H., Ghotbi, C., Rashtchian, D., Solubility of CO2 and H2S in the Ionic Liquid 1-Ethyl-3-methylimidazolium Trifluoromethanesulfonate, Journal of Natural Gas Science & Engineering (2016), doi: 10.1016/j.jngse.2016.02.006. This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customers we are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, and review of the resulting proof before it is published in its final form. Please note that during the production process errors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

ACCEPTED MANUSCRIPT

Solubility of CO2 and H2S in the Ionic Liquid 1-Ethyl-3-methylimidazolium Trifluoromethanesulfonate Mohsen Nematpour a, Amir H. Jalili b,*1, Cyrus Ghotbi a, Davood Rashtchian a a

Department of Chemical and Petroleum Engineering, Sharif University of Technology,

b

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Tehran, Iran. Gas Refining Technology Group, Gas Research Division, Research Institute of Petroleum

Industry (RIPI), P.O. Box: 14665–137, Tehran, Iran.

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ABSTRACT

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The solubility of carbon dioxide and hydrogen sulfide gases in the ionic liquid 1-ethyl-3methylimidazolium trifluoromethanesulfonate ([C2mim][OTf]) was measured at temperatures from (303.15 to 353.15) K and pressures up to about 3.0 MPa. The Henry’s law constants were determined from the new experimental data, which in turn were used to derive the change of some thermodynamic functions of dissolution of the gases in that particular ionic liquid. The new experimental data were correlated by a combination of the extended Henry’s law and Pitzer’s model for the excess Gibbs energy. The average relative percent deviation (ARD%) of correlated molality values from experimental data are within experimental uncertainties, which indicate quite good correlative accuracy of the Pitzer’s model for the systems under investigation. Results show that at the same temperature and pressure, the solubility of H2S in [C2mim][OTf], expressed on the molality scale, is more than four times that of CO2. Comparison of the obtained experimental data indicates that the solubility of H2S, expressed on the molality scale, in [C2mim][OTf] is much more than its magnitude in the high-capacity ionic liquids 1-ethyl-3-methylimidazolium tris(pentafluoroethyl) trifluorophosphate ([C2mim][eFAP]) and 1-ethyl-3-methylimidazolium bis(trifluoromethyl) sulfonylimide ([C2mim][Tf2N]). For example the molality of H2S at 303.15 K and 2.0 MPa is 2.5 times and 1.3 times that of [C2mim][eFAP] and [C2mim][Tf2N], respectively. Solubility of CO2 in the ionic liquids follows the order [C2mim][OTf] < [C2mim][Tf2N] < [C2mim][eFAP]. Compared with other ionic liquids, [C2mim][OTf] could potentially be used for separation of CO2 and H2S gases from each other. Keywords: hydrogen sulfide; carbon dioxide; H2S enrichment; gas separation; natural gas sweetening 1. Introduction

Carbon dioxide comprises one of the most abundant acid gas impurities in many fossil fuel energy resources such as natural gas and associated petroleum gas. Combustion processes *

Author to whom correspondence may be addressed. Tel: +98-21-4825 2466; Fax: +98-21-4473 9716. E–mail address: [email protected].

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ACCEPTED MANUSCRIPT using coal and hydrocarbon fuels, e.g. natural gas, for the production of heat and electricity in power plants also produce carbon dioxide in large scale. Hydrogen sulfide is a highly toxic acid gas naturally occurring with methane, light hydrocarbons and CO2 in many oil and gas reservoirs. Carbon dioxide reduces the heating value of hydrocarbon fuel streams and like hydrogen sulfide causes corrosion in transmission pipelines and process equipment in

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presence of moisture. Furthermore, it’s a common scientific notion that the rise of the amount of carbon dioxide, as one of the most important green house gases, gives rise to global warming of the atmosphere. [1] One of the most versatile processes for removal of CO2 and H2S acid gases from industrial gas streams, especially natural gas, is processes based on

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physical and/or chemical absorption. [2] The former process makes use of physical dissolution of CO2 and/or H2S in a solvent such as methanol (Rectisol process). In the latter case, the acid gases are chemically dissolved in an aqueous solution of a reactive organic

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alkaline compound such as monoethanolamine (MEA), diethanolamine (DEA), and Nmethyldiethanolamine (MDEA). [2] At low concentrations of acid gas impurities processes based on chemical solvents are favored over those based on physical solvents. Chemical absorption processes have some disadvantages such as the loss of alkanolamine during desorption by vaporization, thermal/oxidative degradation to corrosive and toxic by-products

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and high power consumption. [3] The later issue is of extreme importance nowadays as the researchers strive to develop low energy consumption yet more economic processes, which at the same time aim to fulfill requirements set by world environmental organizations concerning global warming of the earth.

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The substitution of conventional alkanolamine solutions using ionic liquids (ILs) [4] in the removal of CO2 and H2S acid gases in gas sweetening processes is an area of ongoing

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research. Imidazolium-based ionic liquids are of special interest among other ionic liquids for the above-mentioned practical applications as well as in acid gas enrichment, CO2 capture from flue and post combustion gases and gas separation and purification for petrochemical industry. Reliable experimental data together with correlations/models, to reproduce the experimental data as accurate as possible, is a prerequisite for primary simulation and evaluation of performance of ionic liquids in industrial processes. Comprehensive reviews for solubility of gases, especially CO2, in ionic liquids and application of ionic liquids in CO2/H2S acid gas removal from industrial gas streams have been presented in the literature. [1,5,6] Experimental data for the solubility of CO2 and H2S in 1-alkyl-3-methylimidazolium trifluoromethanesulfonate ([Cnmim][OTf], where n indicates alkyl chain length) ionic liquids are summarized in Table 1.

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ACCEPTED MANUSCRIPT Table 1 here In previous works in our laboratory at R.I.P.I the solubility of CO2 and H2S in a variety of ionic liquids was measured at temperatures from (303 to 353) K at pressures up to about 2.0 MPa. [17-24] The present work, which is part of an ongoing study on the solubility of acid gases in ionic

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liquids, focuses on the binary systems CO2/H2S + [C2mim][OTf]. All measurements are carried out at temperatures from (303.15 to 353.15) K and pressures up to 3.0 MPa. In case of CO2 + [C2mim][OTf] there are two experimental data sets in the literature reported by other researchers. [7,8] The first data set from intermediate to high pressure regions (0.76 MPa –

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37.8 MPa) belongs to Shin and Lin, [8] and the second series from low to intermediate pressures (0.18 MPa – 5.9 MPa) are reported by Soriano et al. [7] They are both compared here with the corresponding data measured in this work. To the best of our knowledge, there

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isn’t any reported solubility data in the literature for H2S + [C2mim][OTf] system. The new solubility data are modeled using the extended Henry’s law combined with a modification of Pitzer’s activity coefficient model for electrolytes [25,26] on one side and an equation of state to describe the non-ideality of the vapor phase on the other side. Henry’s constant at zero pressure and some characteristic partial molar thermodynamic properties of CO2 and H2S

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dissolved at infinite dilution in [C2mim][OTf] were calculated from the solubility data. The new experimental solubility data are compared to literature data where other ionic liquids were employed for their acidic gases absorption capacity. The comparison allows examining the performance of ionic liquids for industrial applications such as natural gas sweetening and

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2. Experimental

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acid gas enrichment processes.

2.1. Chemicals

In Table 2, the specifications and sources of the chemicals used in this work are summarized. Moisture and other volatile impurities were removed by treating the ionic liquid in vacuum (below 1.0 kPa) at 70 oC for about 24 hours, before use. The water content of the ionic liquid was measured to be below (90 ± 10) ppm, by using a model DL-37 Karl-Fischer Mettler volumetric titrator. All the materials were used without further purification in the course of measurements. Table 2 here

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2.2. Apparatus and procedure The experimental apparatus (Figure 1) described by Jalili et al. [17,22] was used in the experimental part of this work and only a brief description is given here.

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Figure 1 here Known amounts of the gaseous solute and the degassed solvent were equilibrated at a preset temperature in an equilibrium cell of known volume. The difference between two pVT

l g n solute = n total − n solute

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l measurements was used to calculate the amount of gas present in the liquid solution, n solute :

(1)

where n total is the total number of moles of CO2 or H2S transferred from the gas container

ntotal =

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into the autoclave and calculated using Equation (2):

Vgc  pi pf   −  RTgc  Z i Z f 

(2)

where V gc is the gas container volume, Z i and Z f are compressibility factors corresponding to the initial and final pressures pi and pf , respectively, of the gas container, R is the well

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g known gas constant, and Tgc is the temperature of the gas container. n solute in Equation (1)

represents the number of moles of CO2/H2S gas in the gas phase. This was determined using

g nsolute =

pVg ZRT

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the following equation:

(3)

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where Vg is the volume of the gas-phase above the liquid phase, T is the equilibrium temperature of the cell, and Z is the compressibility factor of gas solute at T and p. The temperature of the equilibrium cell was controlled and monitored through connection with a LAUDA, model RE415 water recirculation bath, with temperature stability within ± 0.05 K. The temperature of the cell was measured by using a digital thermometer (Lutron model TM917) with a 0.01 K resolution connected to a Pt-100 sensor inserted into the cell. Keller model PA-33X pressure transmitter sensors were used to measure the pressure in equilibrium cell and the gas container. The pressure sensors, with ranges of (0 to 3) MPa and (0 to 4) MPa, were calibrated against a pneumatic dead-weight gauge (DH. Budenberg model 550). The uncertainty of the pressure measurements corresponds to 0.1 % of the full scale of the

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ACCEPTED MANUSCRIPT sensors. In each experiment equilibration was achieved within three hours when there was no detectable change in pressure of CO2/H2S in the equilibrium cell with time at a fixed temperature. The compressibility factors were calculated by pVT data presented by NIST for pure compounds. [27] The average relative error associated with the measurement of the solubility

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on the molality scale (100·∆m/m) of CO2 and H2S in [C2mim][OTf] is approximately ± 4.1 % and ± 3.0 %, respectively.

3. Results and discussion

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The measured experimental data in the form of pTmG, where mG represents the molality of CO2/H2S gas solute in the ionic liquid phase, in the temperature range from 303.15 K to 353.15 K (10 K steps) and pressures up to about 3.0 MPa, are summarized in Tables 3 and 4.

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The uncertainties associated with the measured T, p and mG values are also listed in Tables 3 and 4. The reliability and accuracy of the apparatus and method of measurement have been assessed previously in our investigations. [17]

Tables 3 and 4 here

Effect of temperature and pressure on solubility of CO2 and H2S gases in [C2mim][OTf] can

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be deduced from the pTmG values listed in Tables 2 and 3. As expected and typical for the solubility of a gas in a neutral solvent, the amount of dissolved gas increases with increasing pressure and decreasing temperature. Furthermore, it can be deduced that H2S is much more soluble in that particular ionic liquid than CO2.

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As was mentioned earlier, there are two other solubility data at intermediate and high pressure regions for CO2 + [C2mim][OTf] binary mixture. [7,8] They are compared with each other and with the data generated in this work in Figure 2 in the form of pressure-molality (P-

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m) isotherm at a typical temperature of about 40 oC. Figure 2 here

It can be observed that the data reported by Shin and Lee, [8] show a large positive deviation, in molality of dissolved gas, from the corresponding data of Soriano et al. [7] and this work. This difference can be deduced at all temperatures from (303 to 343) K (not shown in Figure 2). There is relatively good consistency between the data reported by Soriano et al. [7] and those obtained in this work.

The experimental molality values of this work and those

reported by Soriano et al. [7] are in average within 12 % (with both positive and negative deviations) of each other at the entire reported pressure and temperature ranges. The molality values, obtained from corresponding mole fraction values, reported by Shin and Lee [8]

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ACCEPTED MANUSCRIPT totally show positive deviations in the entire temperature and pressure ranges and in average are situated within about 39 % and 48 % of those of Soriano et al. [7] and this work, respectively. This is the case where at some instances this positive deviation extends to more than 100%. The relatively small difference between the solubility data of Soriano et al. [7] and this work is mainly due to different type of experimental methods and procedures and

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also to different source and purity of chemicals used by the two research groups to measure the solubility of gases in the ionic liquids. Soriano et al. [7] employed a high-pressure thermo-gravimetric microbalance to measure the solubility, while our group used an isochoric saturation method. Shin and Lee made use of a variable-volume bubble pressure

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measurement method for this purpose, [8] which has a proven accuracy mainly at high pressure regions but at low pressures the procedure needs consideration of special precautions. Furthermore, the data due to Shin and Lee [8] seems do not converge to zero

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molality/mole-fraction value in the zero pressure limit. Therefore, the solubility data of Shin and Lee [8] could be considered reliable at high-pressure regions as their data and those produced by Soriano et al. [7] seem to converge beyond 6.0 MPa.

4. Modeling

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4.1. Evaluation of Henry’s constants

Henry’s law constants on the molality scale for dissolution of CO2/H2S gases in the ionic (0) , were evaluated from the new experimental data liquid at zero pressure limits, i.e. k H,m

measured in this work. m / m0 → 0

f (T , p) m / m0

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k H,( 0m) (T ) = lim

(4)

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where f is the fugacity of pure gas (CO2/H2S) at temperature T and pressure p, m is the molality of gas G in the ionic liquid IL and m0 = 1 mol·kg-1 IL. The software package “Thermofluids” was used to calculate the fugacities. [28] Thermofluids uses the equation of state proposed by Span and Wagner for CO2 [29] and the equation of state due to Lemmon and Span for H2S. [30] Figures 3a and 3b show the fugacity f divided by (m/m0) plotted versus (m/m0). The procedure of that evaluation was adopted from Jalili et al., [24] therefore, it is not described here. Figure 3a here Figure 3b here

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ACCEPTED MANUSCRIPT (0) (0) The resulting numbers for k H,m and the estimated uncertainty ∆k H,m (for its estimation also

see Jalili et al. [24]) are given in Table 5. Table 5 here (0) The uncertainty ∆k H,m results mainly from the uncertainty of the extrapolation and the

scattering of the experimental data. In most cases the scattering is well below one percent. To

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account for non-identified systematic errors as well, we assume that the uncertainty of the (0) (0) reported numbers for k H,m corresponds approximately to the maximum number for ∆k H,m ,

i.e., the estimated relative uncertainty of the reported values for Henry’s constant is about ln(k H,( 0m) /MPa) = Ah, m + Bh, m /(T/K) + C h, m ⋅ (T/K) resulting in

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1%. The influence of temperature on Henry’s constant was described by

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ln(k H,( 0m) /MPa) = 21.2756 − 4057 .67 /(T/K) − 0.0243383(T/K) and

ln( k H,( 0m) /MPa) = 8.21030 − 2465 .37 /(T/K) − 0.00368088 (T/K)

(5)

CO2

(5a)

H2 S

(5b)

Equations (5a) and (5b) represent the experimental results for Henry’s constants of CO2 and H2S with an average deviation of 0.10 % (maximum deviation of 0.23 % at 343.15 K) and

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0.22 % (maximum deviation of 0.35 % at 343.15 K), respectively.

4.2. Correlation of data with Pitzer’s model

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The Pitzer’s model was described in detail in previous publications, [24] therefore, only an outline is given here.

As the saturation pressure of the ionic liquid can be neglected, and the pressure does not

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increase beyond 3 MPa the solubility of a single gas in a pure ionic liquid is described by the extended Henry’s law on the molality scale

k H( 0,m) (T )

m γ m = f (T , p) m0

(6)

In Equation (6), γ m stands for the activity coefficient of gas G on the molality scale and it was calculated from a simplified version of Pitzer’s virial expansion for the molality-scale excess Gibbs energy 2

m  m  ln γ m = 2 ⋅ 0 ⋅ β + 3 ⋅  0  ⋅ µ m m 

(7)

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ACCEPTED MANUSCRIPT where β and µ are binary and ternary parameters, respectively, describing interaction between gas molecules in the ionic liquid. The influence of temperature on β and µ was approximated here by

X = AX + B X /(T / K )

(8)

where X represents β or µ . The experimental T and p values, listed in Tables 3 and 4, were

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used to calculate the fugacity of CO2/H2S pure gases as described in section 4.1. The parameters of Equation (8) ( AX and B X ) were obtained by minimizing the sum of squared differences between the calculated gas fugacity values (from the software package Thermofluids [28] at temperature T and pressure p) and the values from extended Henry’s

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law. The correlation equations for Henry’s constants (Equations (5a) and (5b)) were used in all those calculations. The results of the correlations (interaction parameters) are given in

5. Model verification and application

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Table 5.

The Henry’s constants obtained from this work are compared with other [C2mim] – based ionic liquids with different anions in Figure 4.

Figure 4 here

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A monotonic steady increase in Henry’s constants with temperature, which corresponds to a decrease in solubility of CO2 / H2S in the solvents, can be observed from the curves of Figure 4 for all the ionic liquids.

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The three partial molar thermodynamic functions of solution at infinite dilution, i.e., ∆ sol Gm∞ , ∆ sol H m∞ and ∆ sol S m∞ were estimated from the correlation Equations (5a) and (5b). These

functions represent the process of transferring gas G from the ideal gas state at temperature T

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and standard pressure p = p0 = 0.1 MPa to a one molal solution of the gas in the ionic liquid (reference state) where the dissolved gas experiences the same interactions as in infinite dilution. [25] These properties are also given in Table 5. As expected, ∆ sol Gm∞ is positive for both gases and increases with temperature. Also, as expected, ∆ sol H m∞ and ∆ sol S m∞ are negative. ∆ sol H m∞ amounts to about (−10 to −20) kJ·mol-1, as is typical for the solubility of CO2 and H2S in imidazolium-based ionic liquids. Based on the results obtained in this work, the solubility of CO2, expressed in mole fraction scale, in the ionic liquids follows the order: [C2mim][EtSO4] < [C2mim][OTf] < [C2mim][Tf2N] < [C2mim][eFAP], i.e., the solubility increases with increasing number of –

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ACCEPTED MANUSCRIPT CF3 groups in the anion. Baltus et al. also pointed out this same conclusion. [32] Figure 4 shows that the solubility of H2S follows the order: [C2mim][EtSO4] < [C2mim][eFAP] ~ [C2mim][Tf2N] ~ [C2mim][OTf], i.e., although the presence of –CF3 group(s) in the anion results in a profound increase on solubility of H2S in imidazolium – based ionic liquids; however, the solubility does not depend on the number of –CF3 groups. This implies that the

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solubility of CO2 and H2S gases in imidazolium – based ionic liquids obeys different mechanisms. The higher affinity of ionic liquids towards H2S compared to CO2 can be ascribed to a stronger intermolecular attraction, through hydrogen bonding, between H2S and the anion of the ionic liquids, where H2S acts as hydrogen-bond donor and the anion as

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hydrogen-bond acceptor. [22,33] However, in case of CO2 intermolecular attraction, which is weaker than that of H2S, is provided by a different mechanism. In this case, the Lewis acidbase interactions, with CO2 as the acid (electron acceptor) and the anion as the base (electron

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donor), dominates. [34] The above discussions concerns theoretical bases relevant to solubility behavior of acid gases in ionic liquids. However, for practical applications we need to consider molality-scale solubility values/Henry’s constants. Molality scale Henry’s constants, k H,( 0m) , and mole fraction scale Henry’s constants, k H,( 0x) , are related to each other by the simple relation k H,( 0m) = 0.001· k H,( 0x) ·MW, where MW stands for molar mass of the ionic

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liquid in g·mol-1. Comparison of k H,( 0m) values reveals that solubility of H2S in [C2mim] – based ionic liquids follows the order [C2mim][EtSO4] < [C2mim][eFAP] < [C2mim][Tf2N] < [C2mim][OTf]. That is 1 kg [C2mim][OTf] has the highest absorption capacity for H2S gas

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among the studied ionic liquids. However, in case of CO2 the order changes to [C2mim][Tf2N] < [C2mim][EtSO4] < [C2mim][OTf] ≤ [C2mim][eFAP], i.e. CO2 shows the lowest solubility in 1 kg of [C2mim][Tf2N] and the highest solubility in [C2mim][eFAP].

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CO2 shows comparable solubility in [C2mim][OTf] and [C2mim][eFAP]. The quality of the Pitzer’s correlation is discussed here using the average relative deviation ARD % and the maximum relative deviation MRD % in the molality of gas (for preset temperature and solubility pressure)

100 N micor (T , p) − miexp (T , p) ARD % = ∑ N i =1 miexp (T , p)

(9)

 m cor (T , p) − miexp (T , p)  MRD % = max i ⋅ 100  exp mi (T , p)  

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ACCEPTED MANUSCRIPT The results are given in Table 6 together with the corresponding experimental uncertainties ARD %exp and MRD %exp: ARD % exp =

1 N

N

∑ abs(∆m i =1

(11a)

)

exp i

N

MRD%exp = max (∆mexp )i

(11b)

i =1

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where (∆mexp ) i = 100 ⋅ (∆ exp m(T , p ) / m(T , p ) exp )i and ∆ exp m is the relative experimental uncertainty of the molality of gas G. Comparing ARD% and MRD% of correlation with ARD%exp and MRD%exp values (Table 6) shows that the correlation represents the new gas

Table 6 here

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solubility data within experimental uncertainty for both gases.

Figures 5 and 6 graphically compare the correlated values (continuous lines) with the

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experimental data (symbols). It can be seen that the Pitzer’s model is able, both quantitatively and qualitatively, to represent the solubility behavior of CO2 and H2S in [C2mim][OTf].

Figure 5 here Figure 6 here

As was mentioned before (sections 3 and earlier in this section), H2S shows a much more

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solubility in [C2mim][OTf] than CO2. For example it can be deduced, by making use of the Pitzer’s model, that at a pressure of 2.0 MPa the amount of H2S dissolved in one kilogram of [C2mim][OTf] decreases from 8.9 mol·kg-1 IL, by a factor of about 3.5, to 2.5 mol·kg-1 IL when the temperature is increased from 303.15 K to 353.15 K. In case of CO2, the molality

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decreases from (1.2 to 0.63) mol CO2·kg-1 IL (about two-fold decrease) at the same conditions. Figure 7 illustrates the situation graphically by comparing the variation of

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concentrations, expressed in molality, of dissolved CO2 and H2S in some [C2mim] – based ionic liquids as a function of temperature at constant pressure of 2.0 MPa. It can be observed that the solubility of CO2 in [C2mim][OTf] follows a monotonic behavior, which decreases steadily with temperature. The solubility of H2S decreases sharply with increasing temperature. Comparison of the two curves show that at low temperature of 303.15 K H2S is more than 7.5 times as soluble as CO2 in [C2mim][OTf] and this ratio decreases to 4.0 at T = 353.15 K. This observation together with previous discussions concerning molality scale Henry’s constants implies that [C2mim][OTf] could potentially be a good candidate for separation of CO2 and H2S gases from each other for such applications as H2S – enrichment process. Furthermore, Figure 7 reveals that the solubility of H2S, expressed on the molality

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ACCEPTED MANUSCRIPT scale, in [C2mim][OTf] is 2.5 times its solubility in [C2mim][eFAP] at 303.15K and 2 MPa and this ratio gradually decreases to 1.8 as temperature increases to 353.15 K.

6. Conclusions New experimental data for the solubility of CO2 and H2S gases in the ionic liquid

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[C2mim][OTf] are presented in this work. The results show that for both gases [C2mim][OTf] is a “physical” solvent. [C2mim][OTf] has high affinity toward H2S, which is comparable in magnitude with the high capacity ionic liquids [C2mim][eFAP] and [C2mim][Tf2N]. The solubility of CO2 in [C2mim][OTf] is much lower than that of H2S. The obtained

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experimental data were used to evaluate the Henry’s law constants and accompanying thermodynamic functions for dissolution of CO2 and H2S gases in [C2mim][OTf]. They were also correlated by the extended Henry’s law in which the non-ideality of the liquid phase was

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accounted for by using the Pitzer’s model for the excess Gibbs energy. It was shown that the Pitzer’s model has a quite good correlative accuracy for the two binary systems studied in this work. The new experimental results provided further evidence to support the previous notion that the solubility of CO2 increases when the number of –CF3 groups in the anion of the ionic liquid increases. Furthermore, they were used to assess the accuracy of previously

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reported data for solubility of CO2 in [C2mim][OTf]. The large difference between solubilities of H2S and CO2 in [C2mim][OTf] suggests that this solvent could potentially be used (either in pure form or in a mixture with other solvents) for separation of H2S and CO2

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gases from each other or for H2S gas enrichment process.

Acknowledgements

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We are thankful to the Research Council of the Research Institute of Petroleum Industry (RIPI) for their support of this work.

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[15] J.-H. Yim, J. S. Lim, Fluid Phase Equilibr. 2013, 352, 67. [16] I. Mejia, K. Stanley, R. Canales, J. F. Brennecke, J. Chem. Eng. Data 2013, 58, 2642.

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[17] A. H. Jalili, M. Rahmati-Rostami, C. Ghotbi, M. Hosseini-Jenab, A. N. Ahmadi, J. Chem. Eng. Data 2009, 54, 1844. [18] M. Rahmati-Rostami, C. Ghotbi, M. Hosseini-Jenab, A. N. Ahmadi, A. H. Jalili, J. Chem. Thermodyn. 2009, 41, 1052–1055. [19] M. Shokouhi, M. Adibi, A. H. Jalili, M. Hosseini-Jenab, A. Mehdizadeh, J. Chem. Eng. Data 2010, 55, 1663. [20] A. H. Jalili, A. Mehdizadeh, M. Shokouhi, A. N. Ahmadi, M. Hosseini-Jenab, F. Fateminassab, J. Chem. Thermodyn. 2010, 42, 1298. [21] H. Sakhaeinia, A. H. Jalili, V. Taghikhani, A. A. Safekordi, J. Chem. Eng. Data 2010, 55, 5839.

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ACCEPTED MANUSCRIPT [22] A. H. Jalili, M. Safavi, C. Ghotbi, A. Mehdizadeh, M. Hosseini-Jenab, V. Taghikhani, J. Phys. Chem. B 2012, 116, 2758. [23] M. Safavi, C. Ghotbi, V. Taghikhani, A. H. Jalili, A. Mehdizadeh, J. Chem. Thermodyn.

2013, 65, 220. [24] A. H. Jalili, M. Shokouhi, G. Maurer, M. Hosseini-Jenab, J. Chem. Thermodyn. 2013,

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67, 55. [25] K. S. Pitzer, J. Phys. Chem. 1973, 77, 268.

[26] K. S. Pitzer, Activity Coefficients in Electrolyte Solution, Ed.; CRC Press: Boca Raton, FL, 1991.

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[27] NIST Scientific and Technical Databases, Thermophysical Properties of Fluid Systems. http://webbook.nist.gov/chemistry/fluid/ (accessed August 2015).

Heidelberg (Germany), 2006.

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[28] W. Wagner, U. Overhoff, ThermoFluids (Version 1.0, Build 1.0.0), Springer: Berlin,

[29] R. Span, W. Wagner, J. Phys. Chem. Ref. Data 1996, 25, 1509. [30] E. W. Lemmon, R. Span, J. Chem. Eng. Data 2006, 51, 785.

[31] D. Camper, C. Becker, C. Koval, R. Noble, Ind. Eng. Chem. Res. 2006, 45, 445.

108, 721.

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[32] R. E. Baltus, B. H. Culbertson, S. Dai, H. Luo, D. W. DePaoli, J. Phys. Chem. B 2004,

[33] C. S. Pomelli, C. Chiappe, A. Vidis, G. Laurenczy, P. J. Dyson, J. Phys. Chem. B 2007, 111, 13014.

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[34] B. L. Bhargava, S. Balasubramanian, Chem. Phys. Lett. 2007, 444, 242.

13

ACCEPTED MANUSCRIPT Figure Captions

Figure 1. A schematic of the apparatus used for measuring solubility of single gases CO2 and H2S in the ionic liquid [C2mim][OTf].

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Figure 2. Comparison of experimental p versus m data of this work with Soriano et al. [7] and Shin and Lee [8] for solubility of CO2 in [C2mim][OTf] at a temperature of about 40 oC; continuous lines are added to help guiding eyes.

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Figure 3. Plot of the ratio f /m/m0 versus m/m0 at various temperatures for evaluation of

Henry’s law constants for (a) CO2 + [C2mim][OTf] and (b) H2S + [C2mim][OTf]: ◊, T =

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303.15 K; ▲, 313.15 K; □, 323.15 K; •, 333.15 K; +, 343.15 K; ∆, 353.15 K; continuous lines, straight line fitting.

Figure 4. Comparison of mole fraction scale Henry’s constants k H,( 0x) as a function of

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temperature for solubility of CO2 and H2S gases in [C2mim]+– based ionic liquids with different anions; continuous lines are added to help guiding eyes.

Figure 5. Experimental results for the solubility of CO2 in [C2mim][OTf]: ◊, T = 303.15 K; □, 323.15 K; •, 333.15 K; +, 343.15 K; ∆, 353.15 K; continuous lines,

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▲, 313.15 K;

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correlations from Pitzer’s model.

Figure 6. Experimental results for the solubility of H2S in [C2mim][OTf]: ◊, T = 303.15 K; ▲, 313.15 K;

□, 323.15 K; •, 333.15 K; +, 343.15 K; ∆, 353.15 K; continuous lines,

correlations from Pitzer’s model.

Figure 7. Comparison of the molality of dissolved CO2 and H2S gases in [C2mim][OTf] as a function of temperature at p = 2.0 MPa.

14

ACCEPTED MANUSCRIPT Table 1. Experimental data reported in the literature for CO2/H2S + [Cnmim][OTf] (1-alkyl3-methyl imidazolium trifluoromethanesulfonate, n indicates alkyl chain length) binary mixtures. Ionic Liquid

Gas

No. Data

T range

p range

Points

K

MPa 0.18–5.884

References

CO2

30

303.2–343.2

[C2mim][OTf]

CO2

19

303.85–344.55

[C2mim][OTf]

CO2

1

303.15

[C2OHmim][OTf] a

H2 S

41

303.15–353.15

[C2OHmim][OTf] a

CO2

33

303.15–353.15

[C4mim][OTf]

CO2

49

303.85–344.55

0.8–16.8

[8]

[C4mim][OTf]

CO2

26

298.2–333.3

1–9

[12]

[C4mim][OTf]

CO2

17

303.85–344.55

8.5–37.5

[8]

[C4mim][OTf]

CO2

4

294.85–312.05

0.048–0.052

[13]

[C4mim][OTf]

CO2

3

318.15

9–17

[14]

[C6mim][OTf]

CO2

54

303.85–344.55

1.2–18.2

[8]

[C6mim][OTf]

CO2

32

303.15–373.15

1.4–12.3

[15]

[C6mim][OTf]

CO2

16

303.85–344.55

8.5–36.3

[8]

[C6mim][OTf]

CO2

11

297.8–313.1

1.2–8

[16]

[C8mim][OTf]

CO2

44

303.85–344.55

0.68–13.2

[8]

[C8mim][OTf]

CO2

21

303.85–344.55

10–34

[8]

[8]

1

[9]

0.1–2

[10]

0.1–2

[11]

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SC

7.6–37.8

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1-(2-hydroxyethyl)-3-methylimidazolium trifluoromethanesulfonate

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a

15

[7]

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[C2mim][OTf]

ACCEPTED MANUSCRIPT Table 2. Specifications and sources of chemicals used in this work. CAS registry

Mass fraction

Formula

number

purity

Hydrogen sulfide

H2 S

[7783-06-4]

0.995

Carbon dioxide

CO2

[124-38-9]

C7H11F3N2O3S

[145022-44-2]

1-Ethyl-3methylimidazolium

0.995

> 0.99

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trifluoromethanesulfonate

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([C2mim][OTf])

16

Source Roham Gas Company Roham Gas

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Molecular

Chemical name

Company

Merck Chemical Company

ACCEPTED MANUSCRIPT Table 3. Experimental data for CO2 + [C2mim][OTf] measured in this work: T, equilibrium temperature; p, equilibrium pressure; mG, molality of CO2 in liquid phase.a p/

mG /

p/

mG /

p/

mG /

MPa

mol CO2·kg-1IL

MPa

mol CO2·kg-1IL

MPa

mol CO2·kg-1IL

T = 313.15 K

T = 323.15 K

0.077 ± 0.005

0.1477

0.073 ± 0.004

0.4409

0.268 ± 0.015

0.4690

0.236 ± 0.014

0.7682

0.458 ± 0.021

0.8193

0.405 ± 0.019

1.1972

0.728 ± 0.027

1.2806

0.648 ± 0.025

1.7070

1.029 ± 0.034

1.8276

0.914 ± 0.032

2.2405

1.317 ± 0.039

2.3995

1.169 ± 0.036

T = 343.15 K

0.066 ± 0.004

0.4965

0.214 ± 0.013

0.8680

0.367 ± 0.018

1.3546

0.594 ± 0.023

1.9394

0.827 ± 0.029 1.059 ± 0.034

2.5476

T = 353.15 K

0.1644

0.061 ± 0.003

0.1722

0.057 ± 0.003

0.1791

0.054 ± 0.003

0.5221

0.196 ± 0.011

0.5460

0.185 ± 0.011

0.5684

0.175 ± 0.010

0.9137

0.338 ± 0.017

0.9564

0.318 ± 0.016

0.9960

0.303 ± 0.015

1.4265

0.548 ± 0.023

1.4944

0.513 ± 0.022

1.5586

0.487 ± 0.020

2.0441

0.762 ± 0.028

2.1426

0.714 ± 0.026

2.2375

0.677 ± 0.025

2.6867

0.978 ± 0.033

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T = 333.15 K

0.1563

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0.1362

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T = 303.15 K

2.9383

0.888 ± 0.030

0.928 ± 0.031

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Expanded uncertainties u at 95% confidence interval are u(T) = 0.05 K, u(p) = 0.0003 MPa.

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a

2.8140

17

ACCEPTED MANUSCRIPT Table 4. Experimental data for H2S + [C2mim][OTf] measured in this work: T, equilibrium temperature; p, equilibrium pressure; mG, molality of H2S in liquid phase.a p/

mG /

p/

mG /

p/

mG /

MPa

mol H2S·kg-1IL

MPa

mol H2S·kg-1IL

MPa

mol H2S·kg-1IL

T = 313.15 K

T = 323.15 K

0.182 ± 0.009

0.0737

0.166 ± 0.008

0.1384

0.396 ± 0.018

0.1586

0.361 ± 0.017

0.3460

1.028 ± 0.034

0.3991

0.941 ± 0.032

0.6671

2.073 ± 0.047

0.7755

1.896 ± 0.045

1.0146

3.497 ± 0.051

1.1894

3.193 ± 0.051

1.3206

5.036 ± 0.050

1.5607

4.591 ± 0.051

T = 343.15 K

0.159 ± 0.007

0.1780

0.329 ± 0.016

0.4513

0.862 ± 0.031

0.8832

1.733 ± 0.043

1.3634

2.911 ± 0.050 4.170 ± 0.051

1.8027

T = 353.15 K

0.0919

0.139 ± 0.007

0.1004

0.127 ± 0.006

0.1087

0.115 ± 0.006

0.1971

0.301 ± 0.015

0.2152

0.276 ± 0.014

0.2327

0.250 ± 0.013

0.5020

0.790 ± 0.029

0.5511

0.727 ± 0.027

0.5981

0.664 ± 0.026

0.9885

1.585 ± 0.042

1.0890

1.457 ± 0.040

1.1863

1.342 ± 0.038

1.5364

2.650 ± 0.049

1.7024

2.422 ± 0.047

1.8613

2.224 ± 0.046

2.0402

3.788 ± 0.050

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T = 333.15 K

0.0830

SC

0.0643

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T = 303.15 K

2.4553

3.238 ± 0.049

3.468 ± 0.050

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Expanded uncertainties u at 95% confidence interval are u(T) = 0.05 K, u(p) = 0.0003 MPa.

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a

2.2626

18

ACCEPTED MANUSCRIPT Table 5. Thermodynamic properties for the solubility of CO2 and H2S in [C2mim][OTf]: T, temperature; k H( 0,m) , Henry’s constant of gas G at zero pressure on the molality scale; ∆ sol Gm∞ , changes of the molar Gibbs free energy; ∆ sol H m∞ , changes of the molar enthalpy; ∆ sol S m∞ , changes of the molar entropy of gas G.

MPa

∆solGm∞ /

∆ sol H m∞ /

kJ ⋅ mol−1

kJ ⋅ mol−1

J ⋅ mol−1 ⋅ K −1

-73.44

CO2

∆ sol S m∞ /

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(k H( 0,m) ± ∆k H( 0,m) ) /

T/ K

303.15

1.669 ± 0.059

7.09

-15.17

313.15

2.009 ± 0.078

7.81

-13.92

323.15

2.350 ± 0.128

8.48

-12.64

-65.36

333.15

2.687± 0.148

9.12

-11.31

-61.32

343.15

2.995 ± 0.122

9.70

-9.95

-57.25

353.15

3.292 ± 0.109

10.26

-8.54

-53.23

3.18

-17.75

-69.03

M AN U

SC

-69.41

H2 S

0.353 ± 0.004

313.15

0.443 ± 0.004

3.88

-17.56

-68.46

323.15

0.545 ± 0.006

4.56

-17.37

-67.86

333.15

0.660 ± 0.008

5.23

-17.18

-67.24

343.15

0.786 ± 0.010

5.88

-16.97

-66.61

6.56

-16.76

-66.06

EP

TE D

303.15

0.935 ± 0.014

AC C

353.15

Table 6. Interaction parameters for the Pitzer model: T, temperature; X and AX and BX, parameters of Equation (8). System

Parameter X β

CO2 + [C2mim][OTf] H2S + [C2mim][OTf]

AX -0.09335 -0.11931

Deviations µ

BX 3.76815 25.92138

AX 0.02108 0.00168

BX -0.02148 -0.72631

19

ARD% MRD% 0.79 2.7 0.89 2.9

(ARD%)exp 4.1 3.0

(MRD%)exp 5.0 4.8

ACCEPTED MANUSCRIPT Figure 1. A schematic of the apparatus used for measuring solubility of single gases CO2 and H2S in the ionic liquid [C2mim][OTf].

T1

P1

P1: pressure transmitter, 0 – 4 MPa P2: pressure transmitter, 0 – 3 MPa

CO2

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P3: pressure transmitter, 0 – 4 MPa T1, T2, T3: pt-100 sensors, 0 – 100 oC T2

P2

H 2S

P3

T3

Vacuum

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Water out

Equilibrium Cell

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Water in Read-out

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Magnetic stirrer

20

Magnetic Bar

Water Bath

ACCEPTED MANUSCRIPT Figure 2. Comparison of experimental p versus m data of this work with Soriano et al. [7] and Shin and Lee [8] for solubility of CO2 in [C2mim][OTf] at a temperature of about 40 oC; continuous lines are added to help guiding eyes.

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9.0

8.0

SC

7.0

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p / MPa

6.0

5.0

4.0

2.0

1.0

0.0

1.0

AC C

0.0

EP

TE D

3.0

2.0

Shin & Lee [8], 314.05 K Soriano et al. [7], 313.2 K this work, 313.15 K

3.0

4.0

mCO2 / mol CO2 · kg-1 IL

21

5.0

6.0

ACCEPTED MANUSCRIPT Figure 3. Plot of the ratio f / m/m0 versus m/m0 at various temperatures for evaluation of Henry’s law constants for (a) CO2 + [C2mim][OTf] and (b) H2S + [C2mim][OTf]: ◊, T = 303.15 K; ▲, 313.15 K; □, 323.15 K; •, 333.15 K; +, 343.15 K; ∆, 353.15 K; continuous

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lines, straight line fitting.

Figure 3a

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3.5

(a)

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3

2

313.15 K 323.15 K 333.15 K 343.15 K 353.15 K

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f /m/m 0 ̸ MPa

2.5

303.15 K

1.5

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EP

1

0.5

0

0.0

0.2

0.4

0.6

m/m0

22

0.8

1.0

1.2

1.4

ACCEPTED MANUSCRIPT Figure 3b

(b)

1.0 303.15 K 313.15 K

0.9

323.15 K

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333.15 K

0.8

343.15 K 353.15 K

SC

0.6

M AN U

0.5 0.4 0.3

0.0 0.0

1.0

EP

0.1

TE D

0.2

2.0

3.0

m/m0

AC C

f /m/m 0 ̸ MPa

0.7

23

4.0

5.0

6.0

ACCEPTED MANUSCRIPT (0) Figure 4. Comparison of mole fraction scale Henry’s constants k H,x as a function of

temperature for solubility of CO2 and H2S gases in [C2mim]+– based ionic liquids with different anions; continuous lines are added to help guiding eyes.

20

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[C2mim][EtSO4]+CO2 [20] [C2mim][Tf2N]+CO2 [31] [C2mim][eFAP]+CO2 [24] [C2mim][OTf]+CO2, this work [C2mim][EtSO4]+H2S [20]

15

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10

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kH,x(0) / MPa

[C2mim][Tf2N]+H2S [21]

0 310

320

330

T/K

AC C

EP

300

TE D

5

24

340

350

360

[C2mim][eFAP]+H2S [24] [C2mim][OTf]+H2S, this work

ACCEPTED MANUSCRIPT Figure 5. Experimental results for the solubility of CO2 in [C2mim][OTf]: ◊, T = 303.15 K; ▲, 313.15 K;

□, 323.15 K; •, 333.15 K; +, 343.15 K; ∆, 353.15 K; continuous lines,

correlations from Pitzer’s model.

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3.5

3.0

M AN U

SC

2.5

0.5

0.0 0.00

EP

1.0

TE D

1.5

0.25

AC C

p / MPa

2.0

0.50

0.75

1.00

mCO2 / mol CO2 · kg-1IL

25

303.15 K 313.15 K 323.15 K 333.15 K 343.15 K 353.15 K

1.25

1.50

ACCEPTED MANUSCRIPT Figure 6. Experimental results for the solubility of H2S in [C2mim][OTf]: ◊, T = 303.15 K; ▲, 313.15 K;

□, 323.15 K; •, 333.15 K; +, 343.15 K; ∆, 353.15 K; continuous lines,

correlations from Pitzer’s model.

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2.5

M AN U

SC

2.0

p / MPa

1.5

TE D

1.0

EP

0.5

303.15 K 313.15 K 323.15 K 333.15 K 343.15 K 353.15 K

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0.0

0.0

1.0

2.0

3.0

mH2S / mol H2S · kg-1IL

26

4.0

5.0

ACCEPTED MANUSCRIPT Figure 7. Comparison of the molality of dissolved CO2 and H2S gases in [C2mim][OTf] as a function of temperature at p = 2.0 MPa.

10.0 CO2+[C2mim][OTf]

9.0

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H2S+[C2mim][OTf] CO2+[C2mim][eFAP] H2S+[C2mim][eFAP]

8.0

H2S+[C2mim][Tf2N]

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6.0

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5.0 4.0 3.0

TE D

2.0

0.0 300

310

EP

1.0

320

330

T/K

AC C

m / mol acid gas · kg IL-1

7.0

27

340

350

360

ACCEPTED MANUSCRIPT Highlights
New experimental data are reported for solubility of single gases CO2 and H2S in [C2mim][OTf] ionic liquid.
Henry’s constant and thermodynamics functions for solubility of CO2 and H2S in

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[C2mim][OTf] are presented.


The study show that the Pitzer’s model has a good correlative accuracy for CO2/H2S + [C2mim][OTf] binary mixtures.

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This research demonstrates the high absorption capacity of [C2mim][OTf] for H2S compared to previously studied high-capacity ionic liquids [C2mim][eFAP] and

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[C2mim][Tf2N].


Results of this research show that [C2mim][OTf] is a potential condidate for natural

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gas sweetening and acid-gas enrichment processes.