Activity coefficients at infinite dilution of organic solutes, using novel N-(2′, 3′-epoxypropyl)-N-methyl-2-oxopyrrolidinium chloride ionic liquid by GLC

Journal Pre-proof Activity coefficients at infinite dilution of organic solutes, using novel N-(2′, 3′epoxypropyl)-N-methyl-2-oxopyrrolidinium chloride ionic liquid by GLC Vasanthakumar Arumugam, Bakusele Kabane, Kandasamy G. Moodley, Yanan Gao, Gan G. Redhi PII:

S0378-3812(19)30423-6

DOI:

https://doi.org/10.1016/j.fluid.2019.112362

Reference:

FLUID 112362

To appear in:

Fluid Phase Equilibria

Received Date: 12 May 2019 Revised Date:

4 October 2019

Accepted Date: 13 October 2019

Please cite this article as: V. Arumugam, B. Kabane, K.G. Moodley, Y. Gao, G.G. Redhi, Activity coefficients at infinite dilution of organic solutes, using novel N-(2′, 3′-epoxypropyl)-N-methyl-2oxopyrrolidinium chloride ionic liquid by GLC, Fluid Phase Equilibria (2019), doi: https://doi.org/10.1016/ j.fluid.2019.112362. This is a PDF file of an article that has undergone enhancements after acceptance, such as the addition of a cover page and metadata, and formatting for readability, but it is not yet the definitive version of record. This version will undergo additional copyediting, typesetting and review before it is published in its final form, but we are providing this version to give early visibility of the article. 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. © 2019 Published by Elsevier B.V.

Activity coefficients at infinite dilution of organic solutes, using novel N-(2′, 3′epoxypropyl)-N-methyl-2-oxopyrrolidinium chloride ionic liquid by GLC. Vasanthakumar Arumugam,a,b* Bakusele Kabane,a Kandasamy. G. Moodley,a Yanan Gao,b Gan G. Redhia* a

b

Faculty of Applied sciences, Department of Chemistry, Durban University of Technology, South Africa.

Key Laboratory of Ministry of Education for Advanced Materials in Tropical Island Resources, Hainan University, Haikou 570228, Hainan, China

Corresponding Author: [email protected]; [email protected]

1. Abstract The activity coefficients at infinite dilution (

) for 31 solutes, from the following classes of

compounds: alkanes, cycloalkanes, alkenes, cycloalkenes, aromatic hydrocarbons, alcohols, ketones, acetonitrile, thiophene and water were determined using an ionic liquid (IL) modified column utilizing gas liquid chromatography. In this work, N-(2′, 3′-epoxypropyl)-Nmethyl-2-oxopyrrolidinium chloride IL was used to determine the activity coefficients of organic solutes in the range of temperatures from T = 323.15 K to 343.15 K, using 10 K data, the following

intervals, under atmospheric pressure. From these experimental

thermodynamic properties at infinite dilution, including partial molar excess Gibbs free energies

∆

,

, enthalpies

Furthermore, selectivity (

∞

∆

,

and entropy term (Tref∆

) and capacity (

∞

,

) were computed.

) also calculated from experimental value of

for separation of mixtures such as (heptane/benzene, heptane/thiophene, hexane/benzene, water/butanol and hexane/ethanol) at T = 323.15 K - 343.15 K. Generally, the results of this work indicate that the investigated IL has good potential for separating aliphatic solutes from aromatic ones. Specifically, the chosen IL reduces the retention time of most of the solutes used, compared to other reported ILs. Furthermore, the findings of this study suggest that 1

thisIL [N(2′,3′-epoxypropyl)-N-methyl-2-oxopyrrolidinium chloride] may be useful for future applications such as separation and purification in petrochemical industries. Keywords Activity coefficient; ionic liquids; entropy; Gibb’s free energy; enthalpy; separation

2. Introduction On the basis of their wide-ranging properties which have been well-documented, ILs have been used as novel solvents in extraction technology [1-3], as reaction media [4, 5], as phase transfer or biphasic catalyst [6, 7], as electrolyte in batteries, coating technologies [8], pharmaceutical industries [9] and green alternative solvents for volatile organic solvents [1012]. Since ILs are tuneable they have been designed and synthesised for specific applications using different cation-anion combinations [13]. Reports from the literature indicate that pyrrolidinium based ILs have multiple applications due to their unique physiochemical properties; specifically, their good selectivity in the separation of mixtures of aliphatic hydrocarbons and aromatic compounds as well as their use in the extraction of sulfur compounds from fuels [14-16]. The choice of an oxo-pyrrolidinium IL was made on the premise that the variation in structure of a pyrrolidinium IL (via inclusion of an oxo group) would enhance the selectivity of the chosen IL in separating the several classes of compounds. ) provide valuable data on the interaction

Activity coefficients at infinite dilution (

between solute and solvent (in this regard, IL). This could be very helpful in the design of chemical separation processes such as, fractional distillation, gas stripping and liquid-liquid extraction [17, 18]. Generally, high temperatures are required for industrial processes. Thus, there is a need to determine activity coefficients of solutes at various temperatures. 2

Experimental measurements of activity coefficients for several solute-water combinations are not readily available from the literature. However, activity coefficients can be obtained by several methods; but a limited number of studies of

values have been reported and very

few of them reported results over wide ranges of temperatures. [19-27]. ) of solutes in IL systems are measured

Activity coefficients at infinite dilution (

systematically and expand the knowledge and understanding on the impact of structure and properties of IL for the feasibility of separating close boiling mixtures. Basic information on the interaction between solute and IL molecules can be inferred from the values of solute for the selected IL. High

values were observed for polar solutes such as alcohols, ethers

and ketones using IL-based separation techniques [27-32]. Moreover, many of the traditional solvents which are currently used in the separation processes are toxic in nature and have the potential to pollute the eco-system. Therefore, suitable solvents which are environmentally friendly and which possess good properties are needed for industrial processes. For the foregoing reasons, ILs are highly desirable solvents for industrial processes. Nevertheless, it is also important and essential to enlarge the knowledge regarding the mode of combination of IL with numerous solutes. This basic information can be given by activity coefficients at infinite dilution. The selectivity (

=

/

) and capacity (

= #%) values can be $

computed directly from the calculated activity coefficients data for several separation problems. In this work, the IL, namely, 2′, 3′-epoxypropyl-N-methyl-2-oxopyrrolidinium chloride was used to determine

for various organic solutes at several temperatures. From the acquired

information by gas-liquid chromatography at several temperatures, the intermolecular interactions between IL and organic solutes were deduced. Furthermore, the extraction parameters such as selectivity and capacity at infinite dilution were determined. This study 3

describes the determination of activity coefficients of various organic solutes at infinite dilution using an IL- coated column, by the gas-liquid chromatographic (GLC) method. It is noted that the chosen IL gave remarkably reduced retention times for most of the chosen solutes; thus, giving rise to a fast determination of activity coefficients at infinite dilution. [15, 33]

3. Method and Materials Table S1 lists the purities and sources of organic solutes. 3.1 Synthesis of [EPMpyr][Cl] The N-(2′, 3'-epoxypropyl)-N-methyl-2-oxopyrrolidinium chloride [EPMpyr][Cl] IL, was synthesised and characterised according to a method described in our previous reports as depicted in Fig.1 [34-37]. The molecular structure of [EPMpyr][Cl] IL has been shown in Fig.2. Furthermore, the molecular formula and molecular weight of IL are C8H14ClNO2 and 191.66 respectively. The chemical structure of IL was confirmed by 1HNMR and

13

CNMR,

which as shown in Fig. S1 and Fig. S2. The thermophysical properties of synthesised IL has been mentioned in Table S2. The synthesised IL was distilled for 20 hrs at 70 oC, for removing volatile chemicals and water. The Karl-Fischer titration technique was used to determine the amount of water in the IL and was found to be 0.003%. The HPLC method was used to determine the purity of IL. This was found to be 98%. 3.2 Experimental Procedure. The methodology used in this work was reported in our previous publications [15, 33]. The chromatography parameters, used in this work, were sourced from literature [38, 39]. A stainless-steel GC column (1 m length and 4 mm of inner diameter) was used. IL coated chromosorb was used as a packing material for the column. Dichloromethane was used as a 4

solvent medium for uniform spreading of IL onto the surface of packing material. The volume of solvent used for column packing was 50 mL. This was sufficient to prevent any potential solute residues in the column packing material; which was effectively done by the rotary evaporation method. The weight differences of the solid material were measured at (±0. 0001 g precision) before and after coating processes. The uncertainty in moles of IL which was coated on the solid material was ±3 ×10−7 mol. Before commencing the experiment, the column was conditioned and stabilized by passing carrier gas at the high flow rate (about 2 cm3 s-1) at high temperature (373 K), and was done over a period of 48 hrs. The column was prepared with 20% and 30 % concentration of IL. This was used to monitor the separation process of various organic solutes as to determine the suitable concentration of IL. A SHIMADZU gas chromatograph (GC-2014) with column injector and thermal conductivity detector (TCD) was used. The traditional, burette with soap water method was used to determine the flow rate of carrier gas with an uncertainty of ±0.1 cm3min−1. The outlet pressure &' was taken as atmospheric pressure. Pressure drops (& − &' ) were in the range 45 and 180 kPa depending on the flow rate of the carrier gas. A pressure transducer which was employed for the determination of pressure drop had an uncertainty of ±0.1 kPa. The atmospheric pressure was measured using a membrane manometer with an uncertainty of ±0.2 kPa. The volume of the injected solutes was in the range of (0.2 to 0.5) µL in the column for activity coefficients measurements at infinite dilution. The investigated temperatures ranged from T = (323.15 to 343.15) K, at 10 K intervals. Reproducibility was checked by doing the experiments in triplicate mode. It was found that there were very small variations among the repeated measurements. Therefore, an average of duplicate values was taken. The values of

5

retention times varied from 0.5 to 10 minutes depending on the nature of the solutes. Temperatures of the injector and detector were (473 and 523) K, respectively. The evaluation of the stability of stationary phase was checked after a stream of carrier gas was passed through the column for 48 hours; to check the reproducibility of the instrument by determining the retention times for hexane and benzene under experimental condition, on a daily basis. It was observed that there was no deviation in their retention time prior to injection of the test solutes. The error propagation law was used to get the uncertainty of

.

The errors were calculated using the following parameters: flow rate of carrier gas, ±0.0017 cm3 s−1; outlet pressure, ±0.2 kPa; inlet pressure, 0.1 kPa; oven temperature, ±0.1 K; the adjusted retention time )* ±0.6 s; the mass of the stationary phase, ±0.05% and corresponding standard deviations of all these parameters.

4. Theoretical basis The activity coefficients at infinite dilution (

) for organic solutes (as shown in Table 1 )

partitioning between a carrier gas helium and an IL were computed from the retention time of the solute using the following Eq. (1) [40, 41] ln

, -. & ∗ (2 − / ∗ ) &' 45 (22 5 − / ) = ln − + … … … … … … … … . (1) /0 &∗ -. -.

where , represents the number of moles of solvent on the packed column; R is the gas constant; T indicates the temperature of column; /0 shows the net retention volume of the solute; &∗ denotes the saturated vapour pressure of the solute at investigated temperature; 2 represents the second virial coefficient of the pure solute; / ∗ is the molar volume of the solute; &' is the outlet pressure; &' 45 represents the mean column pressure, 2

6

5

(carrier gas)

is the mixed second virial coefficient of the solute and carrier gas and / is the partial molar volume of the solute at infinite dilution in the solvent. Tsonopoulos equation [15] was used to calculate 2 5 . Reference [15] gives a detailed description of the Tsonopoulos equation and reference [33] elaborates on the pressure correction term, 45 . & 2 ;&' < − 1 45 = … … … … … … … … … … … … … … (2) 3 & 5 ;& < − 1 ' Eq. (2) was used to get required parameters for the calculations of

and the values of the

constants were obtained from the literature [42]. /0 represents the overall retention volume of the solute, as given by Eq. (3) /0 = 45

=

>' ()? − )@ ) … … … … … … … … … … … … (3)

)? and )@ represents the times for retention of solute and unreturned gas respectively and >' gives the flow rate of outlet of the column. The linear van’t Hoff relation for temperature dependence is given by Eq. (4), ln

=

A + B … … … … … … … … … … … … … … … (4) .

Equation (4) (linear van”t Hoff relation) was used to calculate the partial excess thermodynamic functions at infinite dilution. Partial molar excess enthalpy, ∆ ∆ In

,

,

= -A and

entropy at infinite dilution can be obtained from the slope and intercept of the plot of vs a/T.

7

,

The excess thermodynamic functions such as ∆

,∆

,

and TEFG ∆

,

(see Table 2) at

infinite dilution are related to activity coefficients at infinite dilution. ,

∆

= -. ln(

,

)=∆

− .∆

,

………(5)

The Eq. (5) can be modified to give Eq. (6) as shown below:

H, (

)=

∆

,

-.

−

∆

,

-

… … … … … … … … … (6)

where R represent the gas constant. From Eq. 4 and Eq. 6, the thermophysical limiting partial molar excess enthalpy and entropy at infinite dilution can be computed from the slope and intercept, that is, ∆

,

= B- and ∆SK , = −A-, respectively. Partial Gibb’s free energy

was calculated using Eq. 5 at reference temperature (333.15 K). These thermodynamic properties were calculated at infinite dilution for the investigated solutes in the IL. The results are presented in Table 2 for 30 % concentration of IL. These thermodynamic functions reveal the fundamental information on the extent of interactions between the solutes and IL under investigation.

5. Results & Discussion This work set out to measure activity coefficients of various organic solutes using three different concentration of N-(2′, 3′-epoxypropyl)-N-methyl-2-oxopyrrolidinium chloride ionic liquid, namely, 20, 30 and 40 %. At 40 % coating of IL on the chromosorb resulted in a surface which appeared to be a sticky semi solid. Thus, it was not suitable for preparing an IL - coated column. Therefore, 20 % and 30 % of IL were used to make two columns separately. These were used to investigate the activity coefficients for 31 organic solutes at three different temperatures, T = 323.15 K to 343.15 K, using 10 K intervals at atmospheric pressure. However, it was found that there were large deviations in data from the two coated 8

columns. In comparing the results from both 20 % and 30 %, it was found that the 30% coated column results were generally higher than that for 20% coated column. These observations suggest that at 20 % of IL, the coverage of the chromosorb by the IL is not optimal. In the light of this, the data from 20% coated was omitted from this report. Table 1 displays the results for 30 % of IL. The existence of interactions between solute and solvent was inferred from the values of activity coefficients, time of retention decreased with increasing

at infinitive dilution. The measured

and decreased with an increase in magnitude

of the interaction between solute and IL. The highest value of

(2544.23) was obtained at

T = 343.15 K, with n-decane as a solute. Lowest values of the activity coefficients were observed for ketones, namely; cyclopentanone and cyclohexanone as (

= 23.57 and

66.44 ) respectively. This implies that these solutes interact strongly with the investigated IL in comparison with other solutes. These strong intermolecular interactions were observed between polar solutes with polar anion of IL. For non-polar solutes, it was found that, the values of

for alkenes are (

∞

= 24.13 to 1810.34) K and for alkanes are (

= 24.96 to

254423) at T = (323.15 K to 343.15 K), respectively. Furthermore, an opposite trend ( values decreases with an increase in temperatures) also was observed for the following solutes, namely, acetone, dichloromethane and isopropyl methyl ketone. The interaction between the solute and anion of IL is attributed to the presence of the nonaromatic ring and an epoxy group in this pyrrolidinium IL. H-bonds play an important role in the interaction of IL with alcohol. However, it was noted that acetonitrile also interacts via Hbond as well as through n-p orbital interaction of epoxy function group of the IL. Low values of

are believed to be responsible for the formation of new H-bonds between alcohol and

IL and the existing self-associated H-bonds in alcohol molecules are broken.

9

The double bonds of respective alkenes specifically interacted with polar anion of the IL, resulting in lower activity coefficient values. However, the same number of carboncontaining alkanes were found to have higher

values than corresponding alkenes. It is

noted that the n-n interactions occurred between anion of the IL and alkane or alkene. In view of this, the values of

should play an important role in the separation of aliphatic from

aromatics. The

values of aromatic compounds namely benzene, m-xylene, p-xylene and

o-xylene are (

= 202.92, 770.58, 599.01 and 562.05) at T = 343.15, respectively. A

complex is formed between the double bond of an aromatic compound and a lone pair of electrons on Cl- anion of the IL. In this complex there is a potential for n-π interactions which will result in stronger binding between IL and aromatics than for other compounds, which do not have aromatic character. The multiple lone pairs of electrons on chloride anion are suggested as the cause of the H-bond basicity of [EPMpyr][Cl] IL. Furthermore, the unsubstituted H-atoms and epoxy groups of pyrrolidinium cation could be the source of Hbond acidity of the IL. The values of

are temperature dependent as shown in Table 1. During interaction of

solute with IL, the effects of exothermic and endothermic properties were evident. Figs. 3 to 8 shows the results for the plots of ln

as a function of the inverse absolute temperature

using 30 % of IL for all the solutes which were selected for this study. The results show that the values of

for alkanes, alkenes and aromatics increase with increasing temperatures

while the opposite was observed for THF, acetone as well as isopropyl methyl ketone (see Figs. 6 and 8). The same type of dependence is shown by methanol whereas other chosen alcohols show weak dependence on temperature as shown in Fig. 4. Based on the chain length of the alkanes, ethers display both types of interactions. The behaviour of ketones is

10

similar to that of aromatic hydrocarbons. In general, the effect of temperature is not significant for most of the solutes except for alkanes. In addition to the usefulness of

limiting partial molar enthalpies ∆

described P,

above, from the experimental values of

of organic solutes in IL at investigated (see Table

2) temperatures, were also computed. The increasing order of

for organic solutes are

alkanes > cycloalkanes > alkenes > aromatic hydrocarbons. Generally, an increase in the alkyl chain length of the solutes results in an increase in the lower





of the solutes. It is postulated that

value of branched and cyclic hydrocarbons result from the structure of the

pyrrolidinium cation of the IL. The 5-membered lactam ring of cation possesses high polar nature with effective solvation properties. It can thus interact effectively with aromatic compounds. Chloride anion based ILs were previously reported for separation applications [43, 44]. Stronger interactions were observed between solute and solvent pairs as well as between solute and solute for polar compounds (excluding water and higher alcohols), namely, aromatic hydrocarbons. This is attributed to negative ∆

P,

values. These negative

values may be ascribed to relatively strong solute-solvent interactions. The substituents of epoxy group and chlorine anion exhibit strong hydrophobic properties which lead to these strong interactions of IL with various organic solutes. Gibb’s free energies and entropies of organic solutes have been shown in Table 2. It was also observed that the Gibb’s free energies of higher alkanes and aromatic hydrocarbons have higher values; strong H-bonding between IL and solutes. Relatively lower values of ∆

,

were observed for

dichloromethane, cyclopentanone, cyclohexane, cyclohexene, tetrahydrofuran, acetone and ether. All organic solutes gave high entropy values, except for 1-hexene, cyclohexene and benzene. This is taken to imply that there are strong H-bonds between organic solutes and IL. Additionally, isopropyl methyl ketone, tetrahydrofuran and acetone gave negative entropy 11

values. This is taken to imply that there are strong interactions between these solutes and IL. The .*QR ∆

partial ,

excess

molar

entropies

of all solutes are positive values except isopropyl methyl ketone, THF and

acetone; the negative entropy may suggest that this particular solute molecule arrange itself in the structure of IL. Many separation problems were overcome using the selectivity data which is derived from activity coefficients at infinite dilution. For developing several extraction processes, the separation factor at infinite dilution needs to be known [45]. Therefore, the data from interaction of IL with organic solutes is very useful for extraction of various materials using ILs as solvents. The following Eq. (7) and (8) were used to calculate the separation and capacity respectively.

=

T

/

=

1

(7)

(8)

In this work, [EPpyr][C] IL was used to determine the coefficients at infinite dilution. This information was used to derive the selectivity and capacity for the separation of mixtures such

as

the

following:

Heptane/Benzene,

Heptane/Thiophene,

Hexane/Benzene,

Water/Butanol, Hexane/Ethanol at temperature of 323.15 K. The results are listed in Table 3; which reveals that higher values of

were observed for alkanes, alkenes, ketones and

alcohols, while lower values were obtained for cycloalkanes, cycloalkenes, cycloketones, dichloromethane and water when using [EPpyr][Cl] IL. The above observation may indicate that the variations could be due to the differences of polarities of organic solutes. The selectivity’s and capacities at infinite dilution for [EPpyr][Cl] were compared to other ILs as 12

shown in Table 3. The highest selectivity values were observed for Heptane/Thiophene and Hexane/Ethanol, namely 3.79 and 3.27 respectively.

Conclusion

An inverse gas chromatography technique for a range of temperature (323.15 K - 343.15 K) was used to determine the activity coefficients, at infinite dilution, for a series of solutes in various concentrations of epoxypropyl pyrrolidinium IL. It is noted that the selected IL resulted in remarkably small retention times. The activity coefficients values have been fully described and discussed. Furthermore, ∆

,

,∆

,

, TEFG ∆

,

, selectivity and capacities

were computed from the measured data. These thermodynamic data may be taken to indicate that molecular interactions between the solute and IL occur via H-bonds, dipole interactions and other weak interactions.

Acknowledgement Dr. Vasanthakumar Arumugam is grateful to the Durban University of Technology, the National Research Foundation (NRF) South Africa for the Innovation Doctoral Grant (Grant UID: 109839) while Prof K G Moodley is indebted to Eskom Holdings, South Africa, for financial support.

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[EMIM][FAP]

Using

Gas-Liquid

Chromatography, J. Chem. Eng. Data 55 (2010) 2444-2450. 39. P-F. Yan, M. Yang, X-M. Liu, C. Wang, Z-C. Tan, U. Welz-Biermann, Activity coefficients at infinite dilution of organic solutes in the ionic liquid 1-ethyl-3methylimidazolium

tetracyanoborate

[EMIM][TCB]

using

gas-liquid

chromatography, J. Chem. Thermodyn. 42 (2010) 817-822. 40. A. Marciniak, M.

Wlazło, Activity coefficients at infinite dilution and

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glycerol. Benzene + nitrogen and benzene + carbon dioxide at 50 C, Trans. Faraday Soc. 65 (1969) 1014-1031. 42. U. Domańska, A. Marciniak, Physicochemical properties and activity coefficients at infinite dilution for organic solutes and water in the ionic liquid 1-decyl-3methylimidazolium tetracyanoborate, J. Phys. Chem. B 114 (2010) 16542-16547. 43. M.G. Prieto, M.D. Williams-Wynn, I. Bahadur, F.A. Sánchez, A.H. Mohammadi, S. Pereda, D. Ramjugernath, Activity coefficients at infinite dilution of hydrocarbons in glycols: Experimental data and thermodynamic modelling with the GCA-EoS, J. Chem. Thermodyn. 105 (2017) 226-237. 44. A. Marciniak, M.

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Kabane,

Gan

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Application

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dicyanamide ionic liquid for various types of separations problems: Activity coefficients at infinite dilution measurements utilizing GLC method, Fluid Phase Equilib. 493 (2019) 181-187 47. F. Mutelet, A.L. Revelli, J-N. Jaubert, L.M. Sprunger, W.E. Acree Jr, G.A. Baker, Partition

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Compounds

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20

Table 1: Average activity coefficients at infinite dilution for the solutes in the ionic liquid (30 %) at three different a temperatures: for the standard state of solutes hypothetical liquid at zero pressure.

Average

values

Solutes n-pentane n-hexane n-heptane n-octane n-Decane cyclohexane hex-1-ene hept-1-ene non-1-ene Dec-1-ene Cyclohexene Methanol Ethanol Butan-1-ol pentan-2-ol Acetonitrile Dichloro methane 4-methyl pentane-2-one Cyclohexanone Cyclopentanone Isopropyl methyl ketone Methyl ethyl ketone m-xylene o-xylene p-xylene Tetrahydrofuron Thiophene Benzene Toluene acetone water a

T = 323.15 K

T = 333.15 K

T = 343.15 K

73.03 262.81 459.34 692.29 1028.42 24.96 229.42 372.09 753.18 921.06 24.13 46.28 80.34 133.58 149.16 125.60 21.08 229.9 35.82 10.48 83.39 119.21 417.44 300.41 356.11 100.98 121.11 162.78 240.28 86.06 29.76

122.10 357.18 691.02 1109.47 1849.00 45.15 274.24 471.04 1035.73 1490.36 41.17 69.19 120.39 202.90 220.24 195.77 28.75 374.55 59.59 17.90 76.56 191.21 566.94 472.23 494.5 67.14 171.96 181.71 375.69 63.97 49.22

172.82 443.16 759.97 1250.80 2544.23 64.19 307.96 597.45 1349.10 1810.34 60.66 82.07 134.47 236.88 260.76 275.16 41.26 509.51 66.44 23.57 33.70 266.92 770.58 562.05 599.01 48.38 194.32 202.92 477.43 59.61 69.94

Standard uncertainties u are u(

) = 5 %, u(T) =0.02 K.

Table 2: Limiting partial molar enthalpies ∆

,

, Gibbs energies ∆

investigated ionic liquid at the reference temperature

solute n-pentane n-hexane n-heptane n-octane n-Decane cyclohexane hex-1-ene hept-1-ene non-1-ene Dec-1-ene Cyclohexene methanol ethanol Butan-1-ol pentan-2-ol Acetonitrile Dichloro methane 4-methyl pentane-2-one Cyclohexanone Cyclopentanone Isopropyl methyl ketone Methyl ethyl ketone m-xylene o-xylene p-xylene Tetrahydrofuron Thiophene benzene toluene acetone water a

,

and entropies

∆

,

for solutes in the

= 333.15 K.

∆ , (kJ mol−1) -39.79 -24.14 -23.26 -27.32 -33.22 -33.97 -13.60 -21.87 -17.95 -31.21 -14.04 -26.46 -23.79 -26.46 -19.51 -22.25 -31.02 -26.40 -28.54 -37.45 41.86 -18.33 -27.74 -28.94 -24.02 33.99 -18.29 -10.18 -31.72 16.96 -28.10

Standard uncertainties u are u(

∆ , ( . ) 13.31 16.28 18.11 19.42 20.84 10.55 15.55 17.05 19.23 20.24 10.30 11.74 13.27 14.72 14.94 14.62 9.30 16.41 11.32 7.99 12.02 14.55 17.56 17.06 17.18 11.65 14.26 14.41 16.42 11.52 10.79

∆

,

( . ) 53.1 40.4 41.4 46.7 54.1 44.5 29.2 38.9 37.2 51.5 24.3 38.2 37.1 41.2 34.5 36.9 40.3 42.8 39.9 45.4 -29.8 32.9 45.3 46.0 41.2 -22.3 32.5 24.6 48.1 -5.4 38.9

) =5 %, u(T) =0.02 K.

Table 3 Comparison of Selectivity and Capacity of various ionic liquids with present study

Selectivity

Capacity

at 323.15 K

IL

Refs Heptane /Benzene

Heptane/ Thiophene

Hexane/ Benzene

Water/ Butanol

Hexane/ Ethanol

Heptane/ Benzene

Heptane/ Thiophene

Hexane/ Benzene

Water/ Butanol

Hexane/ Ethanol

[EPpyr][Cl]

2.82

3.79

1.61

0.22

3.27

0.006

0.008

0.006

0.007

0.012

Present

[P666 14][DCA]

3.87

3.72

-

3.95

-

2.62

2.52

-

2.22

-

46

[EMIM][DCA]

66.2

109

-

-

-

0.39

0.64

-

-

-

47

[P666

4.44

4.97

-

4.53

-

2.18

2.44

-

2.36

-

48

[BMIM][DCA]

44.7

71.5

-

0.29

-

0.51

0.82

-

0.9

-

27

[AMIM][DCA]

64.1

106.9

-

0.19

-

0.32

0.53

-

0.56

-

49

[P1444][TOS]

12.6

19.7

-

1.74

-

0.78

1.22

-

3.14

-

50

[B4MPY][DCA]

50.6

80.5

-

0.33

-

0.72

1.15

-

1.13

-

29

14][TCM]

Figures

O

N

CH3 CN

O

Cl

O

N

Cl

80 oC O +

-

Figure 1: Synthesis of [EPMpyr] [Cl] IL

+

+

Figure 2: (a) 3D structure of the ionic liquid [EPMpyr] [Cl]−. (b) Structure of the ionic liquid [EPMpyr] [Cl]−.

7

n-pentane n-hexane

5.5

n-heptane 4

n-octane n-decane

2.5 2.9

2.95

3

3.05

3.1

1000 T/K Figure 3: Plot of ln

against

1

1/T for the alkanes.

cyclohexane

7 hex-1-ene 5.5

hep-1-ene non-1-ene

4

dec-1-ene cyclohexene

2.5 2.9

2.95

3

3.05

3.1

1000 K/T Figure 4: Plot of ln

against

1/T for the alkenes.

5

methanol ethanol butan-1-ol pentan-2-ol

3.5 2.9

2.95

3

3.05

3.1

1000 K/T Figure 5: Plot of ln

against

2

1/T for the alcohols.

6.5 4-methyl pentan-2-one

5

cyclohexanone cyclopentanone 3.5

isopropyl methyl ketone methyl ethyl ketone

2 2.9

2.95

3

3.05

3.1

acetone

1000 K/T

Figure 6: Plot of ln

against

1/T for the ketones.

benzene

6

toluene m-xylene o-xylene p-xylene 4.5 2.9

2.95

3

3.05

3.1

1000 K/T Figure 7: Plot of ln



versus 1/T for aromatic hydrocarbons.

3

5.8

acetonitrile dichloromethane

4.3

tetrahydrofuran thiophene water 2.8 2.9

2.95

3

3.05

3.1

1000 K/T

Figure 8: Plot of ln



versus 1/T for water, acetonitrile, thiophene, dichloromethane and tetrahydrofuran.

4