Investigation of the Thermophysical Properties of AMPS-Based Aprotic

Article Cite This: J. Chem. Eng. Data XXXX, XXX, XXX-XXX

pubs.acs.org/jced

Investigation of the Thermophysical Properties of AMPS-Based Aprotic Ionic Liquids for Potential Application in CO2 Sorption Processes Sabahat Sardar,*,‡,† Cecilia Devi Wilfred,‡,† Asad Mumtaz,† and Jean-Marc Leveque§ †

Department of Fundamental and Applied Sciences, and ‡Center of Research in Ionic Liquids (CORIL), Universiti Teknologi PETRONAS, 32610 Seri Iskandar, Perak Malaysia § LCME/SCeM, Université de Savoie Mont-Blanc, 73000 Chambéry, France S Supporting Information *

ABSTRACT: The thermophysical properties such as density, refractive index, and viscosity of five aprotic ILs bearing imidazolium (1-ethyl-3-methylimidazolium [emim], 1-butyl-3-methylimidazolium [bmim], 1-benzyl-3-methylimidazolium [bnmim]), pyrrolidinium (1-butyl-1methylpyrrolidinium [bmpyr]), and pyridinium (N-butylpyridinium [bpyn]) cations with 2-acryloamido-2-methylpropanesulfonate [AMPS] anion were studied, and their effect on CO2 solubility was explored. The density and viscosity were determined within the temperature range of (293.15 to 363.15) K at atmospheric pressure, while refractive indices were measured within the temperature range of (288.15 to 333.15) K. Among imidazolium cations, increasing the side chain length resulted in increased refractive index and viscosity with a corresponding decrease in density. In the presented ILs, [emim][AMPS] showed the highest density and least viscosity over the entire temperature range and an enhanced CO2 dissolution of 0.40 mole fraction at 1 MPa was observed at 298.15 K. Moreover, the Henry’s constant of [emim][AMPS] was determined to be 1.957 MPa which was 49.5, 65.5, 21, and 53% less than [bmim][AMPS], [bnmim][AMPS], [bmpyr][AMPS], and [bpyn][AMPS], respectively. The present study provides a better understanding of the structure−activity relationship between CO2 sorption and physicochemical properties of studied ILs. degradation sensitive amine solvents.1,14 For instance with the use of ILs, the required enthalpy to release physically adsorbed CO2 in a regeneration step is about 20 kJ·mol−1, which is only a quarter of the energy consumed by amine-base solvents.15 In principle, ILs are a group of low melting point organic salts in which the ions are poorly coordinated, making them liquid at room temperature (room temperature ionic liquids RTILs) or at least below 373 K. These diversified liquids are of focused interest because of their striking characteristics, that is, negligible vapor pressure, remarkable chemical and thermal stability, potential recoverability, solubility, and recyclability.16−19 Other tunable properties include good solvating capability, wider liquid phase range, nonvolatility, and noninflammability, and also act as catalysts.20−22 These unique properties place ILs among the best possible alternatives for the replacement of organic solvents for CO2 solubility.23 Because of a large available variety of cations and anions, ILs have been conferred with unique terms such as designer, protic, aprotic, task specific, and green reaction media.24−26 The suitable combination of cations and anions can alter physicochemical properties resulting in controlled change in reactivity, selectivity, and catalytic recyclability making

1. INTRODUCTION The Industrial Revolution has led to a foremost and thoughtful problem of CO2 emission which has resulted in an increase of CO2 amount in the atmosphere from about 280 ppm before the start of industrialization to 400 ppm in recent years.1,2 Under the Baseline Scenario, this increase in CO2 emission is estimated to most likely double by 2050.1 There have been various CO2 capture technologies such as chemical absorption,3,4 physical absorption,5,6 adsorption,7 and membrane separation8 that have been studied extensively. Among these technologies, currently most accessible and commercially available technique for CO2 capture is the use of aqueous alkanoamine solutions such as monoethanolamine (MEA).9−11 Although efficient, this technology possesses some limitations such as high equipment corrosion rate along with the requirement of large volume of sorber. Furthermore, amine regeneration also requires high energy consumption as capture of CO2 with amines comprises a chemical reaction with a large reaction enthalpy.12 As a result a large amount of heat is required to get rid of captured CO2 in the regeneration step.1 Ideally a liquid can be regarded as best CO2 sorbent which can adsorb CO2 rapidly and reversibly with desirable physical properties such as low volatility, low viscosity, low heat capacity and high thermal stability.13 Recently, a new class of adsorbing materials, known as ionic liquids (ILs), has been offered as a substitute to the corrosive, volatile, and © XXXX American Chemical Society

Received: June 16, 2017 Accepted: October 25, 2017

A

DOI: 10.1021/acs.jced.7b00552 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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them designer compounds.27 Besides gas solubility and catalysis, these versatile liquids have an extensive range of utility in fields of industrial chemistry, material science, extraction processes, nuclear science, biotechnology, and electrochemistry as electrolytes in fuel cells, solar cells, batteries, biosensors, and electrochemical cells, etc.28−30 Among ILs, aprotic ILs have been extensively studied in electrochemistry, nanotechnology, catalysis, and as reaction media. This subgroup of ILs has been known for more than two decades but unfortunately there is limited literature available on physicochemical properties and characterization of these ILs. Therefore a large data bank is needed, not only for ease of design/processes of target products at the industrial level, but also for the necessary development of their structure−property correlations. Depending upon an endless variety of cations/anions combinations, ILs show a broad range of physicochemical properties. To design inherently green task specific ILs for CO2 capture, physicochemical and thermodynamic properties should be well-known to provide a better structure−activity correlation. Such properties of ILs have been extensively reported, but accurate property predicting tools are still lacking. Because there are so many possible combinations to afford ILs, it is almost impossible to measure the complete set of physicochemical properties, since it would require a vast investment of time and resources. Furthermore, possible intercorrelation between the physicochemical properties for CO2 sorption has not been yet fully explored leading to a drastic lacks of understanding of the dependence of CO2 sorption on physicochemical properties of ILs. In the present work, 2-acryloamido-2-methylpropanesulfonate (AMPS)-based ILs with imidazolium, pyrrolidinium, and pyridinium cations were synthesized and characterized, and their thermophysical properties were measured. The solubility of CO2 in these ILs over various pressures was also reported. The effects of the structural variations of ILs on their thermophysical properties and CO2 solubility in respective ILs were explored to investigate structure−activity relationship.

2.3.1. Viscosity and Density. Viscosity and density measurements were taken using an Anton Paar viscometer (SVM 3000). The instrument was calibrated using ultrapure Millipore-grade water, for which data were established. The uncertainties of viscosity and density measurements are u(T) = ±0.01 K, u(η) = ±1.5% and ur(ρ) = ±0.004 g·cm−3, respectively. The density measurement of [bnmim][AMPS] was taken using an Anton Paar densitometer (DMA 5000) over a temperature range of (293.15−363.15) K and at atmospheric pressure. Standard deviation of density and uncertainty of temperature are u(ρ) = ±0.00019 g·cm−3 and u(T) = ±0.01 K, respectively. 2.3.2. Refractive Index. An ATAGO digital refractometer (RX-5000α) was used to measure the refractive indices of ILs. The instrument was calibrated with pure organic solvents (acetonitrile and methanol) with known refractive index values and validated with ILs of known properties. The temperature range for refractive index measurement was 288.15 to 333.15 K with an accuracy of ±0.05 K and uncertainty of 6.35 × 10−5. Triplicate measurements were recorded and values were reported as an average. 2.3.3. Thermal Decomposition. The measurements of thermal decomposition temperatures were conducted using a thermogravimetric analyzer (PerkinElmer, pyris V-3.81). IL samples were heated from 323.15 to 823.15 K in a crucible under nitrogen atmosphere at a heating rate of 10 K min−1 with ±1 K temperature accuracy. 2.3.4. CO2 Solubility Measurements. The CO2 sorption measurement was determined by using a gas sorption cell. The system used is a volumetric type which uses a pressure drop method (Figure S1). The calculation to determine the mole of CO2 captured is shown in following equation: n=

Pini(Vtot − Vsample) PiniVtot − Z iniRTini ZeqRTeq

(1)

where n is mol CO2 captured, Pini is initial pressure, Vtot is total volume of system, Zini is compressibility factor (PiniTini), Zeq is compressibility factor (PeqTeq), Peq is equilibrium pressure, Vsample is sample volume, R is 8.314 J·K−1·mol−1, Tini is initial temperature and Teq is equilibrium temperature. The compressibility factor was taken by using the Soave−Redlich−Kwong (SRK) equation of state. All the measurements were taken three times, and an average value was reported with 0.5% standard deviation.

2. EXPERIMENTAL SECTION 2.1. Chemicals and Reagents Used. All starting materials were used as received without any further purification procedure. N-Butylpyridinium chloride was a product of Merck while all other reagents were purchased from Sigma-Aldrich. The Chemical information table is provided in the Supporting Information (Table S3). 2.2. Synthesis of Aprotic −SO3 Functionalized Ionic Liquids. A 0.10 mol concentration of each chloride-based precursor ionic liquid (1-ethyl-3-methylimidazolium, 1-butyl-3methylimidazolium, 1-benzyl-3-methylimidazolium, 1-butyl1methylpyrrolidinium, and N-butylpyridinium) was added to 30 mL of anhydrous acetone in a round-bottom flask and then the sodium salt of 2-acrylamido-2-methyl-1-propanesulfonic acid (0.10 mol) was added slowly. The mixtures were stirred under neutral atmosphere at room temperature for 2 days. Filtration was then done under vacuum, and the acetone was evaporated. The ionic liquids were then further dissolved in dichloromethane and washed several times with small quantities of cold water to remove traces of inorganic salt that formed from the metathesis reaction. After a final washing with diethyl ether, the ionic liquids were dried under low pressure in a vacuum oven at 353.15 K for 24 h before measuring physicochemical and thermal properties. 2.3. Characterization. 1H and 13C NMR spectra (Figure S6−S10) were taken of samples in a deuterated solvent (D2O) and recorded on a Bruker Avance 500 spectrometer.

3. RESULTS AND DISCUSSION The ILs were obtained in good yield (87%−93%) with high purity as indicated by NMR, elemental analysis, water content, and halide content. The physical states, water content, and halide content (Tables S1 and S2) along with the data for 1H NMR, 13C NMR, elemental analysis, and FTIR are provided in Supporting Information. The mole fraction purity and analysis methods of synthesized ILs are presented in Table 1. 3.1. Density. Densities (ρ) were measured in the temperature range of 293.15 to 363.15 K at atmospheric pressure, and results are displayed in Figure 1 and Table S4. It was observed that the density decreased in order of [emim][AMPS] > [bmim][AMPS] > [bnmim][AMPS] > [bpyn][AMPS] > [bmpyr][AMPS]. An increase of the side alkyl chain length or incorporation of additional functional group (benzyl group) on the imidazolium cation resulted in decreased density.31−33 Briefly at room temperature, as the alkyl chain was changed from ethyl to butyl on the imidazolium cation, the density dropped from B



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molecular volume (V), standard entropy (S0), crystal energy (UPOT), and isobaric thermal expansion coefficients (αp) of the studied ILs (Table 2 and Table S6). The details of each property, that is, Vm, V, S0, UPOT, and αp is provided in the supporting data. 3.2. Refractive Index. Refractive index values (nD) of the investigated ionic liquids were determined in the temperature range from 288.15 to 333.15 K under atmospheric pressure as shown in Figure 2 and Table S7. The decreasing order of

Table 1. Chemical Name, Purification Method, Mole Fraction Purity and Analysis Method of ILs purification method

chemical name 1-ethyl-3-methylimidazolium 2acryloamido-2methylpropanesulfonate 1-butyl-3-methylimidazolium 2acryloamido-2methylpropanesulfonate 1-benzyl-3-methylimidazolium 2acryloamido-2methylpropanesulfonate 1-butyl-1-methylpyrrolidinium 2acryloamido-2methylpropanesulfonate N-butylpyridinium 2-acryloamido-2methylpropanesulfonate

mole fraction purity

analysis method

extraction + diethyl ether

>0.99

1


>0.98

1


>0.98

1


>0.97

1


>0.97

1

H NMR H NMR H NMR H NMR H NMR

Figure 2. Temperature dependence of the refractive indices (nD) of selected ILs: [emim][AMPS] (a), [bmim][AMPS] (b), [bnmim][AMPS] (c), [bmpyr][AMPS] (d), and [bpyn][AMPS] (e).

refractive indices of the studied ILs is [bnmim][AMPS] > [bpyn][AMPS] > [bmim][AMPS] > [emim][AMPS] > [bmpyr][AMPS]. The refractive index was found to be expectedly and linearly decreasing with increasing temperature. In Table S7, refractive indices of the studied ILs at 298.15 K vary from 1.492 to 1.541 with the lowest and highest values corresponding to [bmpyr][AMPS] and [bnmim][AMPS], respectively. The studied ionic liquids have comparative refractive indices with high-refractive material such as quartz crystal (nD = 1.54) and liquid immersion oil (nD = 1.51),34 which infers that these ILs might be used as optical materials. In imidazolium-based cations, incorporation of more alkyl group/ bulky group or increasing length of alkyl group chain on imidazolium cation caused increase in the value of refractive index which can be attributed to increase in electron density on alkyl group as indicated in Figure 2a−c.35 The standard deviations (SDs) and fitting parameters of refractive indices for the studied ILs are given in Table S8. 3.3. Viscosity. The viscosities (η) of studied ILs were measured in the temperature range of 293.15 to 363.15 K under atmospheric pressure as shown in Figure 3 and Table S9. The viscosity of the presented ILs markedly decreased with increased temperature (Figure 3a−d). The viscosities of ILs decreased in the order [bpyn][AMPS] > [bmim][AMPS] > [bmpyr][AMPS] > [emim][AMPS]. In the case of the imidazolium cation, the

Figure 1. Density (ρ) as a function of the temperature (T) for the aprotic ILs: [emim][AMPS] (a), [bmim][AMPS] (b), [bnmim][AMPS] (c), [bmpyr][AMPS] (d), and [bpyn][AMPS] (e).

1.19588 to 1.15567 g·cm−3, that is, a 3.5% decrease. Similarly, a decrease of 0.65% in density was observed when the butyl group was replaced with the benzyl group (1.15567 to 1.14818 g·cm−3) on the same type of cation. Besides that for the imidazolium cations, the density of [bmpyr][AMPS] was found to be 3.62% less than [bpyn][AMPS]. From a temperature range of 293.15 to 363.15 K, the varying decrements of 4.29% to 4.91% in measured densities were observed in the studied ILs. The increased density of [bpyn][AMPS] may result from strong electrostatic interactions, π−π stacking, and hydrogen bonding in the pyridinium cation, as compared to the pyrolidinium cation, that could lead to efficient high packing. As expected, the measured densities of the ILs decreased linearly with an increase in temperature as shown in Figure 1a−e. The standard deviations (SDs) and fitting parameters of densities for the studied ILs are given in Table S5. The experimental densities were further used to calculate other important properties such as standard molar volume (Vm),

Table 2. Molar Volume (Vm), Molecular Volume (V), Standard Entropy (S0), and Crystal Energy (UPOT) of Present ILs at 298.15 K and at Atmospheric Pressure Vm (cm3·mol−1) V (nm3) S0 (J·K−1·mol−1) UPOT (kJ·mol−1)

[emim] [AMPS]

[bmim] [AMPS]

[bnmim] [AMPS]

[bmpyr] [AMPS]

[bpyn] [AMPS]

265.1771 0.440337 578.3801 412.1777

298.6752 0.495962 647.7166 400.1888

330.2531 0.548398 713.0781 390.7591

314.9131 0.522925 681.3261 395.0051

298.2497 0.495255 646.8354 400.3290

C



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[bmim][AMPS], [bnmim][AMPS], [bmpyr][AMPS], and [bpyn][AMPS] are 577.83, 565, 552.97, 567.45, and 531.84 K, respectively. Thus, thermal stabilities show the trend [emim][AMPS] > [bmpyr][AMPS] > [bmim][AMPS] > [bnmim][AMPS] > [bpyn][AMPS]. 3.5. CO2 Solubility Measurements. 3.5.1. Effect of Pressure and Temperature on CO2 Solubility. The CO2 solubilities of the 5 ILs were determined at 298.15 K with a pressure range of 1 to 4 MPa as shown in Figure 5 and Table S11.

Figure 3. Viscosity as a function of temperature for the present ILs: [emim][AMPS] (a), [bmim][AMPS] (b), [bmpyr][AMPS] (c), and [bpyn][AMPS] (d).

viscosity of [bmim][AMPS] was found to be higher than [emim][AMPS] which can be attributed to the increased molecular weight of bmim as compared to emim.36 This increase in viscosity might be explained by van der Waals interactions caused by the extension of alkyl chain.36 Furthermore, [bnmim][AMPS] being collected as an highly viscous and sticky liquid, its viscosity could not be determined due to instrument limitation. The viscosity of [bpyn][AMPS] was much higher than [bmpyr][AMPS] and might be due to the π−π stacking, increased hydrogen bonding interactions, or electrostatic interaction derived from the pyridinium ring. The fitting parameters for determined viscosities are listed in Table S10. 3.4. Thermal Gravimetric Analysis (TGA). For the present study, the results of TGA analyses are shown in Figure 4. As

Figure 5. CO2 solubilities of ILs as a function of pressure at 298.15 K: [emim][AMPS] (a), [bmim][AMPS] (b), [bnmim][AMPS] (c), [bmpyr][AMPS] (d), and [bpyn][AMPS] (e).

It was observed that CO2 molar uptake increased continuously with increasing pressure with the order of [bnmim][AMPS] < [bpyn][AMPS] < [bmim][AMPS] < [bmpyr][AMPS] < [emim][AMPS]. Among the three imidazolium-based ILs studied, a marginal increase in CO2 solubility was observed for [emim][AMPS], that is, 0.40 mole fraction at 1 MPa, as compared to that for [bmim][AMPS] and [bnmim][AMPS] which have shown 0.22 and 0.15 mole fraction at 298.15 K, respectively. Karadas et al.41 measured CO2 sorption of [emim][Tf2N], [bmim][Tf2N], and [bmim][PF6] at 318 and 338 K for pressures up to 20 MPa. The CO2 sorption for [emim][Tf2N] was larger than that for [bmim][Tf2N] at 318 and 338 K for pressures up to 20 MPa. It was found that at 318 K and 0.995 MPa, x(CO2) for [emim][Tf2N], [bmim][Tf2N], and [bmim][PF6] was found to be 0.16, 0.14, and 0.08, respectively. [emim][Tf2N] and [bmim][Tf2N] are moderately viscous ionic liquids (34.1 and 51 mPa·s, for [emim][Tf2N]42 and [bmim][Tf2N],43 respectively at 298.15 K and 0.1 MPa), while the viscosity for [bmim][PF6] is larger (270.9 mPa·s at 298.15 K and 0.1 MPa44). Aki et al.27,45 reported CO2 solubility using [bmim] cation with different anions. It was found that at 298.15 K, the x(CO2) for [bmim][NO3], [bmim][BF4], [bmim][NO3], [bmim][DCA], [bmim][TfO], and [bmim][methide] were 0.11 at 1.03 MPa, 0.19 at 1.21 MPa, 0.16 at 1.27 MPa, 0.19 at 1.04 MPa, and 0.37 at 1.26 MPa, respectively. The CO2 solubility data for [emim]- and [bmim]-based ILs are displayed in Table 3. Yunus et al.46 determined CO2 solubility of pyridinium-based ILs at temperatures of 298.15, 313.15, and 333.15 K and pressures up to 1 MPa. The solubility of ILs at ambient temperature was in sequence of [C4py][Tf2N] > [C4py][TfAc] > [C4py][Dca]. At 1 MPa, CO2 solubility for the present ionic liquid ([bpyn][AMPS]) with the same cation was 0.2 mole fraction, while

Figure 4. Thermogravimetric analysis of the studied ILs between 323.15 and 823.15 K at a scan rate of 10 K min−1: [emim][AMPS] (a), [bmim][AMPS] (b), [bnmim][AMPS] (c), [bmpyr][AMPS] (d), and [bpyn][AMPS] (e).

reported in the literature, the introduction of functional groups on the structure of the cation or anion results in a decrease in thermal degradation;37−40 similar findings were estimated from the present work. Furthermore, it has also been reported that even substitution of the saturated alkyl chain with the unsaturated one generally reduces the thermal stability of ILs irrespective of the nature of the anions.40 The thermal degradation temperatures calculated for [emim][AMPS], D



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Table 3. CO2 Solubility Data for [emim]- and [bmim]-Based ILs T (K)

P (MPa)

xCO2

ref

1.999 1.999 1.999 1.5 14.9 1.15 16 1.275 75.526 0.979 43.625 1.05 7.3 1.05 9.9 1.999

0.28 0.43 0.39 0.26 0.63 0.27 0.67 0.25 0.59 0.23 0.68 0.14 0.61 0.13 0.43 0.28

Yokozeki et al.47 Yokozeki et al.47 Yokozeki et al.47 Shin and Lee48

Bmim[TFA]

298.15 298.15 298.15 303.85 303.85 303.85 303.85 333.3 323.1 293.43 293.59 293.25 293.65 292.35 313.65 298.1

Bmim[Ac]

298.1

1.999

0.43

Yokozeki et al.47

Bmim[TFES]

298

1.99

0.28

Yokozeki et al.47

Bmim[PRO] Bmim[ISB] C4mim[CF3CF2 CF2CF2SO3]

298.2 298.2 293.15 293.15 343.15 343.15

1.99 2 0.122 0.376 0.169 0.424

0.39 0.40 0.28 0.54 0.22 0.43

Yokozeki et al.47 Yokozeki et al.47 Zhou et al.51

IL

acronym

1-ethyl-3-methylimidazolium trifluoroacetate 1-ethyl-3-methylimidazolium acetate 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide 1-ethyl-3-methylimidazolium trifluoromethane-sulfonate

Emim[TFA] Emim[Ac] Emim[NTf2] C2mim[TfO]

1-butyl-3-methylimidazolium trifluoromethane-sulfonate 1-butyl-3-methylimidazolium acetate

C4mim[TfO] C4mim[Ac]

1-butyl-3-methylimidazolium trifluoroacetate

C4mim[TFA]

1-butyl-3-methylimidazolium Tetrafluoroborate 1-butyl-3-methylimidazolium Thiocyanate 1-butyl-3-methylimidazolium Trifluoroacetate 1-butyl-3-methylimidazolium acetate 1-butyl-3-methylimidazolium 1,1,2,2-tetrafluoroethanesulfonate 1-butyl-3-methylimidazolium Propionate 1-butyl-3-methylimidazolium isobutyrate 1-butyl-3-methylimidazolium Nonafluorobutylsufonate

Bmim[BF4] Bmim[SCN]

x(CO2) for [C4py][Tf2N], [C4py][TfAc], and [C4py][Dca] was 0.24, 0.14, and 0.13, respectively. The solubility of [emim][AMPS] was further studied at different temperatures i.e. 298.15, 313.15, and 333.15 K (Table S13). The results for [emim][AMPS] are plotted in Figure 6. An increase in temperature resulted in a decrease in CO2 sorption in [emim][AMPS]. As expected from gas solubility in liquids, solubility increased with increase in pressure and decreased with increasing temperature for the investigated ILs.27 Because of inadequate mixing, ILs with higher viscosities may lead to uncertainties in measurements carried at low pressures.52,45 Therefore, the CO2 solubility of studied ILs was determined at higher pressures.

Shin and Lee48 Carvalho et al.49 Carvalho et al.49 Revelli et al.50 Revelli et al.50 Yokozeki et al.47

Although AMPS-based ILs have very large viscosities, the anion contains several features such as large molecule with amide and sulfonyl functionality, unsaturation with conjugation, and branched alkyl chain that might lead to better CO2 solubility. 3.5.2. Determination of Henry’s Constant Using the Krichevsky−Kasarnovsky Equation. The Krichevsky−Kasarnovsky equation has been widely used to estimate the solubility of gases in liquid solvents up to high pressure.53−56 The equation is described below:57 ln

f2 (T , P) x2

s

= ln H2P1 +

V 2∞(P − P1s) RT

(2)

where f 2(T, P) represents the fugacity of gas solute 2 in the gas phase at temperature (T) and pressure (P); mole fraction of the Ps gas dissolved in liquid solvents is presented as x2; H2 1 is Henry’s constant of gas in liquid solvents at pressure P1s; V∞ 2 is partial molar volume of gas at infinite dilution of liquid solvents; Ps1 is the standard vapor pressure of liquid solvents; and R is the gas constant. As the vapor pressure of ILs is considered as negligible, the fugacity of gas, f 2(T, P), in gas−IL systems can be substituted for the pure gaseous phase. So the saturated vapor pressure of IL (Ps1) can be considered as zero. Therefore, eq 2 becomes ln

f2 (T , P) x2

= ln H2 +

V 2∞ P RT

(3)

The fugacity of pure gas, f 2(T, P), can be found by using following equation: f2 (T , P) = ⌀2(T , P)P

(4)

where ⌀2 is the fugacity coefficient at pressure P and temperature T, and can be obtained via the SRK equation of state.58

Figure 6. CO2 solubility of [emim][AMPS] at 298.15 K (a), 313.15 K (b), and 333.15 K (c). E



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On the basis of eq 3, a graph between ln( f 2/x2) versus P at 298.15 K is shown in Figure 7, from which partial molar volume

Figure 8. ln( f 2/x2) as a function of pressure at different temperatures for [emim][AMPS] at 298.15 K (a), 313.15 K (b) and 333.15 K (c). The lines were calculated via linear regression. Figure 7. ln( f 2/x2) as a function of pressure for studied ILs at 298.15 K: [emim][AMPS] (a), [bmim][AMPS] (b), [bnmim][AMPS] (c), [bmpyr][AMPS] (d), and [bpyn][AMPS] (e). The lines were calculated via linear regression.

Also the ΔH and ΔS can be calculated from the temperature dependence of Henry’s constant as mentioned below:

of CO2 and Henry’s constant can be obtained at 298.15 K from the slope and intercept of the plot, respectively. According to Henry’s law, the solubility of a gas is directly proportional to the partial pressure of the gas above the surface of the liquid.59 The Henry’s law constants for CO2 solubility in AMPS-based ILs were calculated and are shown in Table 4. Among the studied ILs,

[emim][AMPS] [bmim][AMPS] [bnmim][AMPS] [bmpyr][AMPS] [bpyn][AMPS] a

T (K)

H2 ± σ (MPa)

298.15 313.15 333.15 298.15 298.15 298.15 298.15

1.957 ± 0.005 2.510 ± 0.004 3.276 ± 0.002 3.874 ± 0.003 5.671 ± 0.002 2.477 ± 0.005 4.162 ± 0.001

(5)

⎛ ∂ ln x 2 ⎞ ⎛ ∂ ln H ⎞ ΔS = R ⎜ ⎟ = −R ⎜ ⎟ ⎝ ∂(ln T ) ⎠ P ⎝ ∂(ln T ) ⎠ P

(6)

()

Table 4. Henry’s Constant (H2) for CO2 in Studied ILs at Temperature T and Zero Pressurea IL

⎛ ⎞ ⎛ ∂ ln H ⎞ ⎜ ∂ ln x 2 ⎟ ΔH = −R ⎜ 1 ⎟ = −R ⎜ ⎟ ⎝ ∂(1/T ) ⎠ P ⎝ ∂ T ⎠P

ΔH = H l − H g ,

ΔS = S l − S g

where Hg and Sg are enthalpy and entropy of a pure gas at given temperature and pressure, x2 represents the mole fraction of the gaseous solute at saturation. The calculated enthalpy (ΔH) and entropy (ΔS) of dissolution of CO2 for [emim][AMPS] were −12.138 kJ·mol−1 and 38.476 J·mol−1·K−1, respectively. Physical sorption with enthalpies ranging from −9.7 to −12.8 kJ·mol−1 was observed in conventional organic solvents such as ethanol and heptane,60 suggesting the nature of CO2 solubility in [emim][AMPS] as physical sorption. 3.6. Structure−Activity Relationship of ILs with Their Physicochemical Properties and CO2 Sorption. A large degree of control over physicochemical properties and solubility of CO2 can be easily achieved via the proper tuning of the chemical nature of cations and anions. Although there is no direct correlation reported yet between physicochemical properties and CO2 diffusion in the literature, other than the fact that viscosity has influence on the CO2 solubility in ionic liquids.61−65 However, in the present study, the structure−activity relationship of ILs with physicochemical properties and CO2 solubility has been investigated. In the case of experimental density data observed at room temperature for the three imidazolium-based ILs, the density decreased with an increase of alkyl group length or incorporation of a bulky group following the order of [emim][AMPS] > [bmim][AMPS] > [bnmim][AMPS]. A similar trend was observed in the case of CO2 solubility in these three ILs which might predict that with a decrease in density of imidazolium-based ILs CO2 dissolution had also been decreased. When the densities of [bmpyr][AMPS] and [bpyn][AMPS] are compared, the increased density of [bpyn][AMPS] might result from electrostatic interactions along with π−π

σ is the standard deviation.

the Henry’s constant of [emim][AMPS] was determined to be 1.9571 MPa which was 49.5, 65.5, 21, and 53% less than [bmim][AMPS], [bnmim][AMPS], [bmpyr][AMPS], and [bpyn][AMPS] at 298.15 K, respectively (Figure 7). A low value for Henry’s law constant suggests high gas solubility and vice versa. Moreover, in the case of [emim][AMPS], the Henry constant was increased by 22 and 40.3% by increasing the temperature to 313.15 and 333.15 K, respectively (Figure 8). This illustrates that by increasing the temperature, CO2 sorption was decreased with a subsequent increment in Henry’s constant. The partial molar enthalpy, ΔH, or partial molar entropy, ΔS, of the solution can provide an inclusive insight of the effect of temperature on the CO2 solubility.59 The thermodynamic properties for CO2 sorption in [emim][AMPS] were calculated using Henry’s constant data. In the case of nonvolatile solvent where solubility is sufficiently small and the activity coefficient of the solute is independent of mole fraction, the following thermodynamics expression can be applied to determine partial molar enthalpy and entropy at specific pressure (eq 5 and 6). F



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stacking, and these strong interactions could possibly lead to a drop of CO2 solubility of [bpyn][AMPS] as compared to that of [bmpyr][AMPS]. It is also observed that the calculated molar volume (Vm), molecular volume (V), standard entropy (S0) and isobaric thermal expansion coefficient (αp) of imidazolium-based ILs increased in order of [emim][AMPS] < [bnmim][AMPS] < [bnmim][AMPS] (Table 2) which shows an inverse relation with measured CO2 solubility values. However, the order for crystal energy (UPOT) of ILs and CO2 sorption at 298.15 K was found as [emim][AMPS] > [bmim][AMPS] > [bnmim][AMPS]. In short, it is observed that AMPS-based ILs with low V, Vm, S0, and αp are better candidates for efficient CO2 solubility. Conventionally, increasing the cation alkyl chain length of ILs generally decreases the density of the ionic liquids and subsequently increases free volume in ILs for sorption of CO2 gas to occur via a space filling mechanism.66,67 Although density is decreased with an increase in alkyl chain length in studied imidazolium-based ILs, a simultaneously dramatic increase in viscosity has been observed from 682.74 mPa·s ([emim][AMPS]) to 4874.89 mPa·s ([bmim][AMPS]) at room temperature, that might lead to limited diffusion of CO2. Subsequently the order of Henry’s constant is observed as [emim][AMPS] < [bmpyr][AMPS] < [bmim][AMPS] < [bpyn][AMPS] < [bnmim][AMPS]. Alternatively, although there was an increase in free volume by increasing alkyl chain length, the CO2 diffusion was restricted causing a significant decrease in CO2 sorption due to increased viscosity from [emim][AMPS] to [bmim][AMPS]. It was observed that the CO2 sorption of the studied ILs decreased in the order [emim][AMPS] > [bmpyr][AMPS] > [bmim][AMPS] > [bpyn][AMPS] > [bnmim][AMPS] (Figure 5) while the viscosities showed a reverse trend (Figure 3). High viscosities might be the sum of a sterically hindered anion, van der Waals interactions, chain tangling effect, or hydrogen bonding interactions between the counter parts of ILs,68−70 limiting the mass transfer phenomenon and resulting in slow diffusion of CO2 and ultimately decreased CO2 sorption.71,72 The CO2 diffusion through imidazolium-based ILs has been reported to be strongly dependent on viscosity and fluid temperature.12 ILs with higher viscosities have been reported to be least permeable to all gases due to mass transfer limitations in concerned reactions, thus limiting the gas diffusion and uptake rate.61−65 It was illustrated from results that ILs with lower viscosities have shown better CO2 sorption and vice versa. It also highlighted that the intermolecular and intramolecular interactions of ILs were of great importance in monitoring the CO2 sorption phenomena. 3.7. ILs after CO2 Sorption. The AMPS-based ILs were characterized by FTIR before and after CO2 sorption. Figure 9 shows FTIR spectra before and after CO2 sorption for [emim][AMPS] while the rest of spectra are displayed in the Supporting Information (Figures S2−S5). No peak was observed at 1639 cm−1 for the signature of the carbamate moiety indicating the absence of chemical adsorption on the corresponding ILs, whereas the interaction of CO2 with amine showed a change in the intensity of the N−H vibrational band in the range of 3280− 3450 cm−1 which may correspond to −NH in secondary amides, suggesting the physical interaction of CO2 with studied ILs. 3.8. Recyclability and Reuse of ILs. To determine recyclability of the studied ILs, the evaluation of sorption capacity of ILs for CO2 was performed by successfully recycling up to five times. For that purpose ILs were loaded in to the sorption chamber and evacuated at 348.15 K under vacuum for 2

Figure 9. FTIR spectrum before and after CO2 sorption on [emim][AMPS].

h to remove diffused CO2 from ILs.1,51 To confirm the structural integrity after each recycle, IR and NMR studies were also conducted showing no significant change. The ILs did not show reduced sorption capacity after regeneration up to five trails.

4. CONCLUSION Physicochemical properties of AMPS-based ILs were studied, and their CO2 solubility effect was measured in the pressure range of 1−4 MPa to develop a structure−activity relationship. The thermophysical properties such as density, refractive index, and viscosity were found to decrease with rising temperature. Furthermore, imidazolium cations with lower chain length showed increased density with decreased viscosity and consequently enhanced CO2 sorption. As expected, the solubility of CO2 in studied ILs increased with pressure while the temperature effect was negative. Particularly, the Henry’s constant of [emim][AMPS] was determined to be 1.957 MPa which was 49.5, 65.5, 21, and 53% less than that of [bmim][AMPS], [bnmim][AMPS], [bmpyr][AMPS] and [bpyn][AMPS], respectively. It was also observed that AMPSbased ILs with lower molar volume (Vm), molecular volume (V), standard entropy (S0), and isobaric thermal expansion coefficient (αp) were better candidates for improved CO2 solubility. This study provides a viable route to enhance CO2 sorption by tuning the physicochemical properties using structural variations in counterions of ILs.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b00552. Characterization results, CO2 solubility measurements, and setup, water and chloride contents of the ILs, tables of experimental densities (ρ), standard molar volume, standard entropy, crystal energy, isobaric thermal expansion coefficient, experimental refractive index data, standard deviations (SDs) for experimental refractive indices, experimental dynamic viscosities (mPa·s), correlation coefficients (R2), standard deviations (SD) for dynamic viscosity as a function of temperature, and FTIR of the ILs before and after CO2 sorption (PDF) G



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AUTHOR INFORMATION

Corresponding Author

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

Sabahat Sardar: 0000-0001-8148-661X Funding

The funding for this project is provided by URIF (0153AA-B35), Universiti Teknologi PETRONAS (UTP), Malaysia. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors would like to thank Center of Research in Ionic Liquids (CORIL) at Universiti Teknologi PETRONAS for laboratory facilities and chemicals provided.

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Investigation of the Thermophysical Properties of AMPS-Based Aprotic

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