Synthesis, thermophysical properties and CO2 sorption of imidazolium, thiazolium, iminium and morpholinium-based protic ionic liquids paired with 2-acrylamido-2-methyl-1-propanesulfonate anion

MOLLIQ-111843; No of Pages 8 Journal of Molecular Liquids xxx (xxxx) xxx

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Synthesis, thermophysical properties and CO2 sorption of imidazolium, thiazolium, iminium and morpholinium-based protic ionic liquids paired with 2-acrylamido-2-methyl-1-propanesulfonate anion Sabahat Sardar a, ∗, Asad Mumtaz b, Masoom Yasinzai a, Cecilia Devi Wilfred c a b c

Sulaiman Bin Abdullah Aba Al-Khail - Centre for Interdisciplinary Research in Basic Sciences (SA-CIRBS), International Islamic University, H‐10 Sector, Islamabad, Pakistan School of Natural Sciences (SNS), National University of Sciences and Technology (NUST), 44000, H-12, Islamabad, Pakistan Center of Research in Ionic Liquids (CORIL), Universiti Teknologi PETRONAS, 32610, Seri Iskandar, Perak, Malaysia

a r t i c l e

i n f o

Article history: Received 24 April 2019 Received in revised form 14 September 2019 Accepted 28 September 2019 Available online xxxx Keywords: Ionic liquids AMPS Physicochemical properties CO2 sorption Structure-activity relationship

a b s t r a c t The present study aims to investigate the temperature dependence of the physicochemical properties of four cations (namely 2-chloro-1,3-dimethylimidazolium [diMeim], 3-benzyl-5-(2-hydroxyethyl)-4-thiazolium [bnthia], (chloromethylene)dimethyl iminium [diMeimi] and morpholinium [mrph]) paired with common 2-acrylamido2methyl-1-propanesulfonate [AMPS] anion. The density and viscosity of the resulting ionic liquids (ILs) were measured within temperature range of (293.15–363.15) K while refractive indexes were determined within temperature range of (288.15–333.15) K. Among the investigated novel ILs, [bnthia][AMPS] showed the highest viscosity and density over the entire temperature range and an enhanced CO2 dissolution of 0.65 mol fraction was observed at 1 MPa and 298.15 K. Furthermore, the Henry's constant of [bnthia][AMPS] was found to be 1.199 MPa which was 23.7, 32.6, and 56.0% less than [diMeim][AMPS], [diMeimi][AMPS] and [mrph][AMPS], respectively. The present study provides a detailed insight to understand structure-activity relation between physicochemical properties and CO2 solubility of studied ILs. © 2019 Elsevier B.V. All rights reserved.

1. Introduction Environmental problems, including greenhouse effect, emission reduction and energy savings have become great concerns for all living entities on this life planet. The emission of CO2 from combustion of fossil fuels and electrical power generation has been considered as dominant contributors to climate change as well as global warming [1,2]. The Inter-governmental Panel on Climate Change (IPCC) has perceived that by the year of 2100 there may be a rise of approximately 2 °C in global temperature along with other anticipated distresses [3,4]. This has made CO2 capture and sequestration into an extensively investigated topic since last decade. Many ways to reduce CO2 emissions have been projected, such as physisorption/chemisorption [5,6] membrane separation [7] or molecular sieves [8], mineral carbonation [9], amine dry scrubbing [10], etc. Though these remedies have their own acute issues related to large energy consumption, high corrosion, low capacity and high cost. Hence other more cost-effective and efficient technologies need to be explored to overcome discrepancies arising ∗ Corresponding author., E-mail addresses: [email protected], [email protected] (S. Sardar).

from conventional methods. Ideally, best CO2 sorbent can be that liquid which can sorb CO2 rapidly and reversibly with desirable physical characteristics such as low viscosity, low volatility, low heat capacity and high thermal stability [11]. The flourishing research on room temperature ionic liquids (RTILs) as environmentally benign gas processing agents is being conducted because of their tunable properties. Ionic liquids have been considered as greener liquids due to noncorrosive, nonvolatile and nonflammable behavior while recyclability and high thermal stability are considered as energy saver properties. Numerous families of ILs, such as imidazolium, pyrrolidinium, pyridinium etc. have been tried to search for the most potential one [12,13]. Improved CO2 sorption results have been reported on branching/elongation of alkyl chain, ether linkage or presence of hydroxyl/nitrile group on core structure of cation, fluorination on cation and/or anion core of respective ILs [14]. CO2 sorption studies on molecular level suggest that anionic counter-ion plays a key role than the cation while extension of alkyl chain length supports free volume or molar volume [15]. Conversely, IR spectroscopy favors the theory of Lewis acid-base interaction between CO2 and fluorinated anion of IL [16]. Interestingly, viscosity and other physicochemical properties of ILs can be altered by careful choice of structural entities (cation, anion or alkyl chain etc.) to get

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Please cite this article as: S. Sardar, A. Mumtaz, M. Yasinzai, et al., Synthesis, thermophysical properties and CO2 sorption of imidazolium, thiazolium, iminium and morpho..., Journal of Molecular Liquids, https://doi.org/10.1016/j.molliq.2019.111843

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desired tasks. So, complete knowledge of the physicochemical properties of ILs, to be used as sorbent, is necessary to build up possible structure-activity correlation. In continuation of research progress involving ILs as sorbents, the current study is an attempt to measure CO2 solubility in four new ILs bearing different cationic core structures paired with common anion at a pressure range of 1–4 MPa. The physicochemical properties, such as density, refractive index and viscosity were also determined over a broad temperature range to investigate their influence on the CO2 sorption capacity in each IL. The effects of the structural variations of cations on measured physicochemical properties were deeply explored to predict possible structure-activity relationship with CO2 solubility in respective ILs. 2. Experimental 2.1. Materials The ILs were synthesized using analytical grade chemicals and were proceeded without any purification. All the starting materials, including 2-chloro-1,3-dimethylimidazolium chloride, 3-benzyl-5-(2hydroxyethyl)-4-thiazolium chloride, (Chloromethylene)dimethyl iminium chloride, morpholinium chloride, sodium salt of 2acrylamido-2-methyl-1-propanesulfonic acid (AMPS) and anhydrous acetone were purchased from Sigma Aldrich. The Chemical information table is provided in supplementary information (Table S1). 2.2. Synthesis of ionic liquids 0.1 mol of each chloride-based precursor chlorides (2-chloro-1,3dimethylimidazolium, 3-benzyl-5-(2-hydroxyethyl)-4-thiazolium, (Chloromethylene)dimethyl iminium, and morpholinium) was added to 25 mL of anhydrous acetone in a round bottom flask. Na-salt of 2acrylamido-2-methyl-1-propanesulfonic acid (0.1 mol) was then added slowly and mechanical stirring was carried out under nitrogen at ambient temperature for 48 h. The filtration was done to separate solid salt (dissolved in acetone) from the synthesized ILs and then solvent acetone was evaporated via rotary. ILs were dissolved further in dichloromethane followed by extraction with ice cold water (to remove the traces of salts if present). Finally after rigorous washing with diethyl ether and ethyl acetate, the ILs were dried in oven under vacuum at 353.15 K for 48 h before measuring thermal and physicochemical properties. 2.3. Characterization 1 H and 13C NMR spectra were obtained using deutrated methanol and recorded on a Bruker Avance 500 spectrometer. The chemical analysis of ILs was performed using an elemental analyser: CHNS-932 (LECO instruments). ILs were also characterized using FTIR (Schimadzu) and the spectra were recorded in range of 500–4000 cm−1.

Refractive index. The refractive indices of present ILs were taken using an ATAGO programmable digital refractometer (RX-5000α). The instrument was initially calibrated with pure organic solvents (methanol and acetonitrile) of known refractive index values and also validated with ILs of known properties. The refractive index measurements were recorded over a temperature range of (288.15–333.15) K with an accuracy of ±0.05 K and uncertainty of ±4 × 10−5. At least three measurements were taken at each temperature and values were reported as an average. Thermal decomposition. The thermal decomposition temperatures of ILs were recorded using a thermogravimetric analyser (Perkin-Elmer, Pyris V-3.81). The samples of ILs (4–8 mg) were placed in an aluminium pan under N2 atmosphere at a heating rate of 10 K min−1 from 323.15 to 823.15 K with ±1 K temperature accuracy. 2.5. CO2 sorption measurements The CO2 solubility measurement was investigated using gas sorption cell. The used system is a volumetric-type emphasized on pressure drop method (Fig. S1). The calculation used to estimate mole of CO2 captured is shown in equation (1):

Z i ni:R:T i ni−

n ¼ P i ni:V t ot P i ni:ðV t ot− V s ampleÞ

ð1Þ

Z e q:R:T e q

where n is mole CO2 captured, Pini is the initial pressure, Vtot is the total volume of system, Zini is the compressibility factor (Pini.Tini), Zeq is the compressibility factor (Peq.Teq), Peq is the equilibrium pressure, Vsample is the sample volume, R is 8.314 J.K−1mol−1, Tini is initial temperature and Teq is equilibrium temperature. The compressibility factor was measured using Soave-Redlich-Kwong (SRK) equation of state. All the measurements were taken thrice and an average value was reported with standard deviation of 0.5%. 3. Results and discussion The AMPS-based ILs were obtained in excellent yields (95%–99%) with high purity as indicated by elemental analysis, NMR, water content and halide content. The detailed description of synthesized ILs (Table S2), their water and halide contents (Table S3) along with the data for 1H-NMR, 13C-NMR, FT-IR and elemental analysis are provided in supplementary information. The spectra for FTIR and NMR are also provided in Figs. S2–S4 and Figs. S5–S12, respectively. 3.1. Density Fig. 1 and Table S4 represent the temperature dependence of density (ρ) of the present ILs in the temperature range of (293.15–363.15) K and at atmospheric pressure. For studied ILs, density decreased linearly with an increase in temperature as shown in Fig. 1. The density values of present ILs increased in order of [diMeim][AMPS] b [diMeimi][AMPS] b [bnthia][AMPS]. At 298.15 K, the values of density were 1.176 g cm−3 for [diMeim][AMPS], 1.184 g cm−3 for [diMeimi][AMPS] and 1.191 g cm−3

2.4. Properties measurements The water contents of studied ILs were measured using coulometric Karl Fischer titrator (model DL-39). The chloride contents were determined using ion chromatogram (Metrohm model 761 Compact IC). Density and Viscosity. The density and viscosity measurements of synthesized ILs were measured using the Anton Paar Viscometer (model SVM3000) over a temperature range of (293.15–363.15) K and at atmospheric pressure. Before each measurement, the instrument was calibrated using ultrapure Millipore quality water and was also validated with ILs of established properties. The standard uncertainties of measurements are u(T) = ± 0.01 K, u(η) = ± 1.5% and ur(ρ) = ± 0.004 g cm−3 for temperature, viscosity and density, respectively.

Fig. 1. Density (ρ) as a function of the temperature (T) for the protic ILs; [diMeim][AMPS]: (a), [bnthia][AMPS]: (b), [diMeimi][AMPS]: (c).

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S. Sardar et al. / Journal of Molecular Liquids xxx (xxxx) xxx for [bnthia][AMPS]. The experimental densities of investigated ILs did not show a relationship with molecular weight of cation instead depended on structure of cation. Similar observation was also found in other studies [17,18]. At 298.15 K the density obtained for [diMeim][AMPS] (1.176 g cm−3), was 1.81% less than the value obtained for [emim][AMPS] (1.1977 g cm−3; imidazolium-based IL with same anion i.e., AMPS) [19]. It can be observed from the densities of [emim][AMPS] and [diMeim][AMPS] for which the density values increased at least by 1.84% when chloride group was incorporated in the imidazolium ring at C2-position. This drop in value of density of [diMeim][AMPS] could be due to strong electronegative nature of chloride group that may result in increase in charge separation, polarity and dipole moment in imidazolium ring. Similarly the density of [bnthia][AMPS] at 293.15 K was found to be 1.191 g cm−3, while reported density for 5-(2-hrdroxyethyl)-4-methylthiazolium (same cation as that of [bnthia][AMPS]) with bis(trifluoromethane)sulfonimide anion is 1.504 g cm−3 [20]. This decreased density value of [bnthia][AMPS] might be due to strong counter-ionic interactions, as besides temperature the density values of materials also depend on cation-anion interaction, size and shape of ions, and upon how ions are closely packed with each other [21]. Because of lack of available density data for iminium based ILs in literature, the density values for [diMeimi][AMPS] could not be compared. The densities of studies ILs were fitted using least square method to the below mentioned linear equation:
 ρ g:cm−3 ¼ A0 þ A1 T

ð2Þ

where, ρ denotes the density of investigated ILs in g.cm−3, T is temperature in Kelvin, and A0 and A1 are correlation coefficients. The values of correlation coefficients were estimated using least square method. The correlation coefficients, fitting parameters along with standard deviations for equation (2) are provided in supporting information (Table S5). The estimated densities were further used to calculate other significant properties such as standard molar volume (Vm), molecular volume (V), standard entropy (S0), crystal energy (UPOT), and isobaric thermal expansion coefficients (αp) of the studied ILs (Table 1 and Table S6). Standard molar volume (Vm) is the volume occupied by 1 mol of a substance at standard temperature and pressure. The Vm values of the prepared ILs were calculated at room temperature and atmospheric pressure using the following equation:   1 Vm cm3 :mol ¼ M=ρ

ð3Þ

where Vm is molar volume in cm3. mol−1, M is the molecular weight in g.mol−1, and ρ is the density in g.cm−3 at 298.15 K. The values of Vm are displayed in Table 1. The Vm values of the investigated ILs follows the order of [diMeimi][AMPS] b [diMeim][AMPS] b [bnthia] [AMPS]. The molecular volumes (V) of the ILs were calculated from molar volume and Avogadro's constant (NA) using the following equation V ¼ V m =N A

ð4Þ

The calculated molecular volumes of the ILs are also displayed in Table 1. The molecular volume of ILs increased in the same order as that of Vm. Glasser et al. [22,23] has introduced an empirical equation for determining standard entropy (S0) based on material's ionic charge, molecular volume and chemical formula. As the proposed equation does not require any structural information, so it can be used to calculate S0 of amorphous solids, ILs and hypothesized materials as well [24]. From calculated values of molecular volume (Vm), the standard entropy S0 of studied ILs were measured using following equation [25]; S0



−1

J:K−1 mol




 ¼ 1246:5 V nm3 þ 29:5

ð5Þ

where S0 is standard entropy and V is molecular volume of studied ILs in nm3. The values of S0 of present ILs are reported in Table 1. It was observed that the standard entropy (S0) followed the same trend as standard molar volume (Vm). The lattice energy of a salt is associated with potency of interactions between its ions and depicts their relative stabilities. In accordance with Glasser's theory, the crystal energy (UPOT) can be calculated using the following equation   −1 U P OT kJ:mol ¼ 1981:2ðρ=MÞ1=3 þ 103:8

3

calculated UPOT values are presented in Table 1. The UPOT values of studied ILs decreased in order of [diMeimi][AMPS] N [diMeim][AMPS] N [bnthia][AMPS], which was opposite to trend of molar and molecular volume. The isobaric thermal expansion coefficients (αp) of the investigated ILs were calculated from the experimental densities using equation (7).   1 ∂ρ A3 αp ¼ − : T p ¼− ρ ∂T A2 þ A3

ð7Þ

where αp, ρ, T and subscript p are thermal expansion coefficient, density, absolute temperature and constant pressure, respectively. Furthermore, A2 and A3 are the fitting parameters. The calculated values of αp as a function of temperature are included in the Supplementary Information (Table S6). The variations of αp values with the temperature were insignificant for the present ILs. Thermal expansion coefficients of present ILs increased slightly with increase in temperature and showed an inverse relation with corresponding densities. The calculated αp values are in the range of (6–7.31) × 10−4 K−1 which show the characteristic ILs behavior [26]. The αp values of ILs were found to be lesser than those of molecular organic compounds; i.e., in case of toluene, the values are in the range of 10.7–11.3 × 10−4 K−1 [27]. Also, thermal expansion coefficient values of synthesized ILs are observed to be higher than those of high temperature fused salts; for instance, αp values for KCl and NaCl at 1100 K are 3.62 × 10−4and 3.09 × 10−4 K−1, respectively [27]. 3.2. Refractive index Refractive index values (nD) of the present ILs were determined in the temperature range from (288.15–333.15) K under atmospheric pressure as shown in Fig. 2 and Table S7. The decreasing order of refractive indices of the investigated ILs is [bnthia] [AMPS] N [diMeim][AMPS] N [diMeimi][AMPS]. The refractive index was found to be linearly decreased with increasing temperature. The refractive indices of the studied ILs at 298.15 K vary from 1.4969 to 1.5514 with the lowest and highest values corresponding to [diMeimi][AMPS] and [bnthia][AMPS], respectively. No direct relation between the molar volume and refractive indexes of studied ILs was observed. Almeida et al. [28] reported that the refractive index slightly depends on the volume of anion, while Deetlefs et al. [29] reported that refractive index would be higher when the materials are more tightly packed to each other. Also, the refractive index of ILs depends on the nature of the functional groups attached with cationic portion of ILs [30]. The reported data of refractive index values are in good agreement with above statements. The temperature dependence of the refractive indices were further correlated using linear equation (8), nD ¼ A4 þ A5 T

ð8Þ

The standard deviations (SDs) and fitting parameters of refractive indices for the studied ILs are given in Table S8. 3.3. Viscosity The viscosity is a significant physicochemical property which results from species interactions such as van der Waals, hydrogen bonding and columbic forces. It is estimated that viscosity is affected by size and shape of IL ions. Higher viscosities could also be unfavorable for certain industrial applications such as mass transfer operations, synthetic solvents, conductivity and field-effect transistors [18,31]. On the other hand, large viscosities are desirable for other applications such as lubrication [31]. The viscosities (η) of prepared ILs were determined in the temperature range of (293.15–363.15) K under atmospheric pressure as shown in Fig. 3 and Table S9. The viscosity of the presented ILs markedly decreased with increase in temperature (Fig. 3). This may be due to increased movements of ions at higher temperatures which weakens the interionic forces prevailed between counter-ions and the liquid becomes less resistant to flow [17]. The viscosities of studied

ð6Þ

where ρ and M are the density (g.cm−3) and molecular mass (g.mol−1), respectively. The

Table 1 Molar volume (Vm), molecular volume (V), standard entropy (S0) and crystal energy (UPOT) of studied ILs at 298.15 K and at atmospheric pressure.

Vm (cm3mol−1) V (nm3) S0 (J.K−1.mol−1) UPOT (kJ.mol−1)

[diMeim][AMPS]

[bnthia][AMPS]

[diMeimi][AMPS]

288.99 0.4799 627.67 403.46

369.80 0.6141 794.93 379.82

252.25 0.4189 551.62 417.36

Fig. 2. Temperature dependence of the refractive indices (nD) of selected ILs; [diMeim] [AMPS]: (a), [bnthia][AMPS]: (b), [diMeimi][AMPS]: (c).

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Fig. 3. Viscosity as a function of temperature for the present ILs: [diMeim][AMPS]: (a), [bnthia][AMPS] (b), [diMeimi][AMPS] (c).

ILs decreased in order [bnthia][AMPS] N [diMeim][AMPS] N [diMeimi][AMPS]. At ambient temperature, IL with thiazolium-based cation [bnthia]+ showed a profoundly greater viscosity compared to other ILs (Table S9) that can be attributed to existence of benzyl and -OH group on cation that result in π-π stacking, increased hydrogen bonding interactions and/or electrostatic interaction derived from thiazolium ring. At 298.15 K, the values obtained for [diMeim][AMPS] was 16.03% higher than [diMeimi][AMPS]. However at higher temperatures, the difference in viscosities between [diMeim][AMPS] and [diMeimi] [AMPS] was less significant. The observed drop in viscosity with increase in temperature may be due to increased movements of ions that reduces resistant to flow of the ILs [32]. The viscosity of [diMeim][AMPS] was found much lower as compare to imidazolium-based ILs paired with AMPS anion [19]. At 298.15 K, viscosity of [diMeim] [AMPS] was found to be 228.51 mpa. s, while viscosity values of [emim][AMPS] and [bmim][AMPS] were found to be 668.13 and 2992.23 mpa. s, respectively. The reduced viscosity of [diMeim][AMPS] as compared with imidazolium-based ILs paired with AMPS anion may be due to substitution of chlorine group on the imidazolium cation and presence of an electronegative atom may alter interionic forces among ions of ILs. As reported by Seki et al., the viscosity increases with increase in molecular weight of anion and does not depend on molecular weight of cation [33]. The obtained values of density and viscosity also did not depend on molecular weight of cation but instead depended on cation structure and cation-anion interactions. Furthermore, [mrph][AMPS] being collected as yellow waxy solid, its viscosity could not be determined due to instrument limitation. Experimental viscosities (η) were fitted as a function of temperature using equation (9).

ln ηðmPa:sÞ ¼ C 0 þ

C 1 c2 þ 2 T T

ð9Þ

where η is the viscosity, T is the temperature and C0, C1, and C2 are the adjustable parameters. The fitting parameters are listed in Table S10. Seddon et al. [34] suggested that ILs display non-Arrhenius behavior hence VogelFulcher-Tammann (VFT) equation can be applied to explore the variation of transport properties with respect to temperature. The experimental values of viscosity of used ILs were correlated using VFT model through equation (10):

ηðT Þ ¼ ηo :exp½

Bη  T−T o

ð10Þ

where η is viscosity at temperature T, ηo is the high temperature limit of η, Bη presents fitting coefficient controlling the curvature and To is Vogel temperature that normally lies a few tens of degrees below Tg. At 298.15 K, the derived coefficients of VFT equation and η are displayed in Table 2. It was found that the correlated viscosities were in good agreement with the experimental data (Fig. 4A).

Table 2 Fitting parameters of VFT (Eq. (10)) and Arrhenius equations (Eq. (11)) at 298.15 K η∞/mPa. Eη/J.mol−1 s

ILs

ηo/mPa.s

Bη/K

To/K

[diMeim] [AMPS] [bnthia] [AMPS] [diMeimi] [AMPS]

0.042 ± 0.053 0.033 ± 0.485 0.055 ± 0.039

1004.52 ± 0.36 1134.15 ± 2.3 953.81 ± 0.27

193.43 42.87 203.35 307.7 193.99 87.58

52043.04 ± 0.023 69516.41 ± 0.027 49906.51 ± 0.022

Fig. 4. Plot of experimental viscosity (symbols); (A) Vogel-Fulcher-Tammann fitting (dotted lines) and (B) Arrhenius fitting (dotted lines), as a function of temperature for: [diMeim][AMPS]: (a), [bnthia][AMPS] (b), [diMeimi][AMPS] (c).

The Arrhenius equation (in logarithmic form) can be used to estimate the energy required for free movement of ions in the respective liquid.

ln η ¼ lnη∞ þ

Eη RT

ð11Þ

where η is dynamic viscosity, η∞ is the viscosity at infinity temperature, Eη is the activation energy, R is the universal gas constant (8.314 J K−1 mol−1) and T is the temperature in Kelvin, respectively. The empirical parameters i.e., η∞ and Eη, are tabulated in Table 2. The activation energy (Eη) is the minimum energy needed for free movement of ions in medium. With drop of Eη, the movement of ions in particular media becomes more feasible. As seen from Table 2, Eη of investigated ILs are close to activation energies of reported ILs [35]. In Fig. 4 (A and B), the experimental viscosity (symbols) had been fitted with VFT equation (dotted lines) and Arrhenius equation (dotted lines), respectively. 3.4. Thermal gravimetric analysis (TGA) Dealing with elevated temperatures, it is necessary to examine the behavior of the ILs at the operating conditions and determine their maximum operation temperatures for safe and green applications. The thermal stability of studied ILs was studied by thermogravimetric method (TGA) under nitrogen atmosphere. TGA thermogram is presented in Fig. 5. It is reported that thermal stability of ILs depends on the coordinating nature of anionic counterpart [36], i.e. strongly coordinating anions result in decrease in decomposition temperatures of ILs and poorly coordinating anions increase the thermal stability. The thermal degradation temperatures calculated for [diMeim][AMPS], [bnthia][AMPS], [diMeimi][AMPS] and [mrph][AMPS] were found to be 380.41, 501.33, 444.82, and 513.76 K, respectively. Thus thermal stabilities show the trend [diMeim][AMPS] b [diMeimi][AMPS] b [bnthia][AMPS] b [mrph][CS]. 3.5. CO2 solubility of ILs The CO2 solubility performance in all four ILs was measured experimentally at a temperature of 298.15 K from 1 to 4 MPa by CO2 sorption cell. The experimentally determined CO2 solubility data for [diMeim][AMPS], [bnthia][AMPS], [diMeimi][AMPS] and [mrph] [AMPS] are listed in Table S11 and the values are plotted for graphical representation as Fig. 6. The mole fraction-pressure plot reveals the CO2 sorption capacity of ILs, where

Please cite this article as: S. Sardar, A. Mumtaz, M. Yasinzai, et al., Synthesis, thermophysical properties and CO2 sorption of imidazolium, thiazolium, iminium and morpho..., Journal of Molecular Liquids, https://doi.org/10.1016/j.molliq.2019.111843

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Fig. 5. Thermogravimetric analysis of the studied ILs between 323.15 and 823.15 K at a scan rate of 10 K min−1; [diMeim][AMPS]: (a), [bnthia][AMPS]: (b), [diMeimi][AMPS]: (c), [mrph][AMPS]: (d).

5

pressures where there would be no available free volume occupied by CO2 molecules. In other words, further addition of CO2 gas into respective ILs would be nearly impossible unless strong cohesive structure of IL is broken [13]. The CO2 sorption of studied ILs was higher than reported AMPS-based ILs [19]. In our previous study [19], CO2 solubility of imidazolium-based ILs i.e., [emim][AMPS], [bmim][AMPS] and [bnmim][AMPS], at 10 bar and ambient temperature was found as 0.4, 0.22 and 0.15 mol fraction, respectively. While in current study, the CO2 sorption of [diMeim][AMPS] (an imidazolium-based IL), was found as 0.46 mol fraction which was 13%, 52% and 67% more than [emim][AMPS], [bmim][AMPS] and [bnmim][AMPS], respectively. Similarly, the Henry's constant reported for a thiazolium based IL bearing same cation as that of [bnthia][AMPS] with NTf2 anion was 4.56 MPa [20], while Henry's constant for [bnthia][AMPS] was found as 1.19 MPa, showing high CO2 sorption of [bnthia][AMPS] than reported one. Furthermore the comparison of CO2 solubility in said ILs, with literature, on the basis of their respective structures was difficult because of unmatched cation/anion combination of ILs with each other. However at comparable pressures and temperatures, current ILs showed higher CO2 sorption in comparison of reported ILs (Table 3). The solubility of [bnthia][AMPS] was further studied at higher temperatures i.e. 313.15 and 333.15 K. The results for CO2 solubility of [bnthia][AMPS] are plotted in Fig. 7. As expected, the increase in temperature resulted in decrease in CO2 solubility in [bnthia] [AMPS]. The Krichevsky-Kasarnovsky (K-K) equation is widely used to investigate the solubility of gases in liquid solvents up to high pressures [46]. The equation is mentioned below [47]: ln


 V ∞ P−P s1 f 2 ðT; P Þ ¼ ln H P2 1s þ 2 x2 RT

ð12Þ

where f2(T, P) is the fugacity of gas at pressure (P) and temperature (T), x2 is mole fraction of the gas dissolved in liquid solvents, H P2 1s is Henry's constant of gas at pressure Ps1 in liqs uid solvents, V∞ 2 is partial molar volume of the gas at infinite dilution of liquid solvents, P1 is the standard vapor pressure of the liquid solvents and R is the gas constant. The fugacity of gas, f2(T, P), in gas-IL systems can be substituted for pure gaseous phase, as the vapor pressure of ILs is considered as negligible,. So Ps1 of IL can be considered as zero. Hence, equation (12) becomes: ln

f 2 ðT; P Þ V∞ ¼ ln H 2 þ 2 P x2 RT

ð13Þ

The fugacity of pure gas, f2(T, P), can be obtained by using below mentioned equation: f 2 ðT; P Þ ¼ ∅2 ðT; P ÞP

Fig. 6. CO2 solubilities of ILs as a function of pressure at 298.15 K; [diMeimi][AMPS] (a), [bnthia][AMPS] (b), [diMeim][AMPS] (c) and [mrph][AMPS] (d).

[bnthia][AMPS] was found to be most promising candidate and [mrph][AMPS] the least efficient. However the overall CO2 sorption in ILs demonstrated the trend of [bnthia][AMPS] N [diMeim][AMPS] N [diMeimi][AMPS] N [mrph][AMPS] at different pressures. At 298.15 K, CO2 solubility in ILs first increased with increasing pressure but tended to level off at high

ð14Þ

where ∅2 is fugacity coefficient at pressure P and temperature T, and can be obtained via the SRK equation of state [48]. Based on equation (13), a graph between ln(f2/x2) versus pressure at 298.15 K is shown in Fig. 8. From Fig. 8, the partial molar volume (V∞ 2 ) of CO2 and Henry's constant (H2) can be obtained at 298.15 K from the slope and intercept of the plot, respectively. Henry's law states that the solubility of a gas is directly proportional to the partial pressure of the gas above the surface of the liquid [49]. The Henry's law constants for CO2 solubility from K-K equation of present ILs were calculated and are shown in Table 4. Among the studied ILs, the Henry's constant of [bnthia][AMPS] was determined to be 1.199 MPa

Table 3 Comparison of CO2 sorption by studied ILs with competitive reported ILs. ILs

Acronym

T (K)

P (MPa)

xCO2

Ref.

1-ethyl-3-methylimidazolium 2-acryloamido-2-methylpropanesulfonate

[emim][AMPS]

1

Emim[TFA] C2mim[TfO]

1-butyl-1methylpyrrolidinium 2-acryloamido-2-methylpropanesulfonate

[bmpyr][AMPS]

1-butyl-3-methylimidazolium trifluoroacetate

C4mim[TFA]

1-butyl-3-methylimidazolium Thiocyanate 1-butyl-3-methylimidazolium Nonafluorobutylsufonate

Bmim[SCN]

Tetrabutyl ammonium bromide Tetrabutyl ammonium tetrafluoroborate Urethane-imide based poly(ionic liquid) Cellulose-based poly(ionic liquid)

[TBA][Br] [TBA][BF4] PUIS-02-TBA [CelEt3N][PF6]

0.4 0.32 0.37 0.28 0.26 0.63 0.32 0.42 0.48 0.51 0.23 0.68 0.13 0.43 0.28 0.54 0.22 0.43 0.47 0.75 33.1 mg/g 168 mg/g

[19]

1-Ethyl-3-methylimidazolium trifluoroacetate 1-Ethyl-3-methylimidazolium trifluoromethane-sulfonate

298.15 303.15 313.15 298.15 303.85 303.85 298.15

C4mim[CF3CF2 CF2CF2SO3]

293.43 293.59 292.35 313.65 293.15 293.15 343.15 343.15 293.15 298.15 303.15 298.15

1.999 1.5 14.9 1 2 3 4 0.979 43.625 1.05 9.9 0.122 0.376 0.169 0.424 0.995 0.082 3

[37] [38] [19]

[39] [40] [41]

[42] [43] [44] [45]

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S. Sardar et al. / Journal of Molecular Liquids xxx (xxxx) xxx Table4 Henry's constant H for CO2 in prepared ILs at temperature T and zero pressurea. IL [bnthia][AMPS] [diMeimi][AMPS] [diMeim][AMPS] [mrph][AMPS] a

T (K)

H2 ± σ (MPa)

298.15 313.15 333.15 298.15 298.15 298.15

1.199 ± 0.05 1.315 ± 0.04 1.591 ± 0.02 1.779 ± 0.04 1.572 ± 0.04 2.728 ± 0.05

σ is the standard deviation.

The calculated enthalpy (ΔH) and entropy (ΔS) of CO2 solubility for [bnthia][AMPS] were −6.72 kJ mol−1 and 21.38 J mol−1 K−1, respectively, suggesting the existence of physical sorption of CO2 in [bnthia][AMPS] [50]. 3.6. Correlation between physicochemical properties and CO2 sorption: structure-activity relationship Fig. 7. CO2 solubility of [bnthia][AMPS] at 298.15 K (a), 313.15 K (b) and 333.15 K (c). Krichevsky-Kasarnovsky Equation; Graphs between ln(f2/x2) vs P. which was 23.7, 32.6, and 56% less than [diMeim][AMPS], [diMeimi][AMPS] and [mrph] [AMPS] at 298.15 K, respectively (Fig. 8). A low value for Henry's law constant shows high gas solubility and vice versa. In case of [bnthia][AMPS], the Henry constant was increased by 9.67 and 32.69% by increasing the temperature to 313.15 and 333.15 K, respectively (Fig. 9). This demonstrated that by increasing the temperature, CO2 solubility was decreased with subsequent increment in Henry's constant. [emim][AMPS] and [bmim][AMPS], imidazolium-based ILs paired with AMPS anion, showed CO2 sorption of 0.4 and 0.225 mol fraction at 1 MPa and 298.15 K, respectively [19]. While, CO2 solubility of [diMeim][AMPS] was 62.5% and 65.38% higher than [emim][AMPS] and [bmim][AMPS], respectively. Petrick et al. [20]. reported CO2 solubility of six novel thiazolium-based ILs paired with TF2N− anion and 3-(2-ethoxy-2oxoethyl)-4-methylthiazol-3-ium bis(trifluoromethyl)sulfonyl)amide was found to be most promising sorbent for CO2 gas. The respective IL showed Henry's constant of 3.55 MPa which was 66.2% higher than [bnthia][AMPS], showing [bnthia][AMPS] a better CO2 sorbent as compared to reported thiazolium-based ILs. The partial molar enthalpy (ΔH) or partial molar entropy (ΔS) of a solution can offer a broad insight of temperature effect on the CO2 solubility [49]. The thermodynamic properties for CO2 solubility in [bnthia][AMPS] were calculated using Henry's constant data. In non-volatile solvent, where solubility is small enough and the activity coefficient of the solute is independent of mole fraction, the below mentioned thermodynamic expressions can be used to measured ΔH and ΔS at specific pressure (equations (15) and (16)). 0

1   B ∂lnx2 C ∂ln H C ΔH ¼ −R B @ 1A ¼ R ∂ð1=T Þ p ∂ T p  ΔS ¼ R

   ∂lnx2 ∂ln H ¼R ∂ðln T Þ p ∂ðln T Þ p

ð15Þ

ð16Þ

The physicochemical properties and CO2 sorption in ILs can be fluctuated with the combination of anions and cations. The theory of anion effect suggests that free volume lies in interstitial space near cationic alkyl chain and hence cations with long chain are favorable for increasing free volume with increased CO2 solubility. Furthermore, quantum chemical calculations and molecular dynamic simulation [51,52] speculate that free volume lies in interionic space between cation and anion, emphasizing that weaker cationanion interactions are favorable for increased CO2 sorption. In this respect, molar volumes of the ILs can be related with CO2 solubility in ILs [13], i.e., larger the molar volumes of ILs, higher would be CO2 solubility. Hence molar volumes, calculated from experimental values of densities of present ILs, have been used to describe CO2 sorption of studied ILs. In order to develop better understanding of the influence of different cations (paired with common anion) on CO2 sorption, calculated Vm of each IL and their respective CO2 sorption results in X2 have been displayed in Fig. 10. It can be observed from Table 1 that high Vm of [bnthia][AMPS] may be responsible for high CO2 sorption of respective IL i.e., AMPS-based ILs with high Vm gave better CO2 solubility and vice versa. The results obtained from Fig. 10 supported the theory of dominancy of free volume over anioninteraction with CO2. High viscosities may be the sum of chain tangling effect, bulky anion, van der Waals interactions or hydrogen bonding interactions between the counter ions of ILs limiting the mass transfer phenomenon and consequently result in slow diffusion of CO2 [17,53]. Although [bnthia][AMPS] has higher viscosity than other investigated ILs, still it was found to be most promising candidate among the studied ILs. At 298.15 K, the viscosity values for [bnthia][AMPS], [diMeim][AMPS] and [diMeimi][AMPS] were found to be 6289.31, 695.85 and 574.5 mPa s, respectively. On the other hand, CO2 sorption of [bnthia][AMPS] at 298.15 K and 1 MPa was found to be 25.99% and 40.05% higher than [diMeim][AMPS] and [diMeimi][AMPS], respectively. Because of higher viscosity of [bnthia][AMPS], the equilibrium time found for [bnthia][AMPS] was 31 h while time for [diMeim][AMPS] and [diMeimi][AMPS] to reach equilibrium was 19 and 15 h, respectively. After attaining equilibrium, further increase in time duration did not affect the CO2 solubility in either case. Hence in present the study, the high viscosity directly affected the rate of CO2 diffusion while X2 remained unaffected.

ΔH ¼ H l −Hg ; ΔS ¼ Sl −Sg

3.7. ILs after CO2 sorption

where Hg and Sg are enthalpy and entropy of pure gas at given temperature and pressure, x2 is the mole fraction of the gaseous solute at saturation.

The IR spectra of studied ILs were taken before and after CO2 sorption. Fig. 11 shows FTIR spectra before and after CO2 sorption for [bnthia][AMPS] while rest of spectra are

Fig. 8. ln(f2/x2) as a function of pressure for studied ILs at 298.15 K; [diMeimi][AMPS] (a), [bnthia][AMPS] (b), [diMeim][AMPS] (c) and [mrph][AMPS] (e). The lines were calculated via linear regression.

Fig. 9. ln(f2/x2) as a function of pressure at different temperatures for [bnthia][AMPS] at 298.15 K (a), 313.15 K (b) and 333.15 K (c). The lines were calculated via linear regression.

Please cite this article as: S. Sardar, A. Mumtaz, M. Yasinzai, et al., Synthesis, thermophysical properties and CO2 sorption of imidazolium, thiazolium, iminium and morpho..., Journal of Molecular Liquids, https://doi.org/10.1016/j.molliq.2019.111843

S. Sardar et al. / Journal of Molecular Liquids xxx (xxxx) xxx

7

current study provides a sustainable route to improve CO2 sorption by proper tuning the thermophysical properties using structural variations in ions of ILs. Acknowledgment The authors would like to thank Higher Education Commission (HEC) of Pakistan for Start-Up Research Grant (2444) and Center of Research in Ionic Liquids (CORIL) at Universiti Teknologi PETRONAS for the financial support from the grant number URIF\/0153AA-B35. Appendix A. Supplementary data Supplementary data to this article can be found online at https://doi. org/10.1016/j.molliq.2019.111843. Fig. 10. Plot of experimentally determined CO2 solubility (X2) at pressure of 1 MPa and temperature of 298.15 K versus calculated molar volume (cm3. mol−1).

displayed in supporting information (Figs. S2–4). No distinct peak at 1545 cm−1 and 835 cm−1 was observed for the signature of carbamate moiety and bicarbonate anion, respectively, indicating the absence of chemical sorption on the corresponding ILs. The result showed that ILs anion could be chemically inert during sorption process, While, the interaction of CO2 with amine of anion showed a change in the intensity of N-H vibrational band in the range of 3296–3450 cm−1 which may correspond to –NH in secondary amides suggesting the physical interaction of CO2 with studied ILs. New distinct bands, corresponding to ammonium cation, between 2800 and 3000 cm−1 and in 200–2800 cm−1 region, were also not found [54]. Instead a new peak was observed at 2338 cm−1 that corresponded physically dissolved CO2. 3.8. Recyclability and reuse of ILs To investigate the recyclability of the present ILs, the evaluation of sorption capability of ILs for CO2 was checked by successful recycling up to 4–5 times. For that purpose, studied ILs were loaded in to the sorption chamber and evacuated at 348.15 K under vacuum for 3 h to remove diffused CO2 from ILs [1,41]. To confirm the structural integrity of investigated ILs after each recycle, IR and NMR studies were also conducted showing no significant change. The ILs did not show reduced sorption after regeneration up to 5 trails.

4. Conclusion The present work highlighted the potential investigations on physicochemical properties of AMPS-based ILs and the effect of these properties on CO2 solubility was further explored in the pressure range of 1–4 MPa to develop a structure-activity relation. Thiazolium cation with chloro group substitution exhibited increased density and viscosity along with enhanced CO2 sorption. The Henry's constant of [bnthia] [AMPS] was determined to be 1.199 MPa which was 23.7, 32.6 and 56% less than [diMeim][AMPS], [diMeimi][AMPS] and [mrph][AMPS], respectively. It was also found that AMPS-based ILs with higher molar volumes were promising candidates for enhanced CO2 solubility. The

Fig. 11. FTIR spectrum before and after CO2 sorption on [bnthia][AMPS].

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