- Email: [email protected]
Experimental and theoretical study on infinite dilution activity coefficients of various solutes in piperidinium ionic liquids
J. Chem. Thermodynamics 60 (2013) 169–178
Contents lists available at SciVerse ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Experimental and theoretical study on infinite dilution activity coefficients of various solutes in piperidinium ionic liquids Kamil Paduszyn´ski ⇑, Urszula Doman´ska Department of Physical Chemistry, Faculty of Chemistry, Warsaw University of Technology, Noakowskiego 3, 00-664 Warsaw, Poland
a r t i c l e
i n f o
Article history: Received 18 November 2012 Received in revised form 2 January 2013 Accepted 4 January 2013 Available online 26 January 2013 Keywords: Piperidinium ionic liquids Activity coefficients Extraction LSER models Regular solution theory Solubility parameters
a b s t r a c t The aim of this work is to summarize our systematic studies on activity coefficients at infinite dilution (c1 12 ) of various organic compounds and water in 1-n-alkyl-1-methylpiperidinium cation-based ionic liquids with bis[(trifluoromethyl)sulfonyl]imide anion, [CnC1Pip][NTf2] (n = 3–6). New sets of experimental data measured by gas-liquid chromatography are reported for 43 different solutes in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] in temperature range T ¼ ð308:15 to358:15Þ K. Moreover, ambient pressure liquid densities of those ionic liquids are presented at different temperatures. The results are discussed in terms of an influence of structure of both solute and ionic liquid on c1 12 . Capabilities of the studied ionic liquids in liquid-liquid extraction of aromatics or sulphur compounds from aliphatic hydrocarbons are demonstrated in terms of infinite dilution selectivity and capacity. Finally, linear solvation energy relationship (LSER) theory and regular solution theory (RST) were used to analyze the obtained experimental data. In particular, correlative and predictive power of the most current version of the LSER is demonstrated and the Hildebrand’s solubility parameters are calculated as a function of temperature by using the RST. Ó 2013 Elsevier Ltd. All rights reserved.
1. Introduction Ionic liquids (ILs) are organic salts forming a melt below 100 °C [1]. Peculiarly low solid–liquid phase transition temperature observed for this family of chemicals is closely related to overall asymmetry in the size, shape and chemical structure of ions composing a salt. The cations constituting common ILs are based mainly on functionalized and quaternarized organic bases (e.g., imidazoles, pyridines, amines), whilst the anions are usually smaller, more symmetric and inorganic or organic (e.g., halides, tetrafluoroborate, hexafluorophosphate, alkylsulphates, dialkylphosphates). Those disparities weaken the electrostatic interactions between the cations and anions and thus, substantially lower the value of lattice energy of an ionic liquid compared to ‘‘traditional’’ inorganic salts like NaCl. Since over twenty years ILs have been recognized as modern and clean (‘‘green’’) replacements for volatile organic compounds (VOCs) in many areas of pure and applied chemical sciences, e.g., synthesis and catalysis [2]. In particular, incorporating ILs as environmentally friendly entrainers for sustainable units and processes of chemical and petrochemical technologies seems to be one of their the most promising practical utilities [3]. This is mainly due to extrelemely low volatility, good thermal and chemical stability, enhanced capability of dissolving other substances (ranging from ⇑ Corresponding author. Tel.: +48 222345640. E-mail address: [email protected] (K. Paduszyn´ski). 0021-9614/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2013.01.005
traditional inorganic salts [4] to chemically simple and more complex components of biomaterials, e.g., carbohydrates [5]) and tailorable properties of ILs. Liquid–liquid extraction and azeotropic distillation of close boiling mixtures are the most important and the most widely studied examples of the possible chemical engineering applications of ILs [6–8]. In particular, several different separation problems have been investigated in literature. The most common and representative ones are: aromatic from aliphatic hydrocarbons [9–11], sulphur and nitrogen compounds (e.g., thiophene and pyridine) from hydrocarbons (model mixtures for design of novel processes of extractive desulphurisation/denitrification of gasolines and oils) [12–17] and short-chain alcohols from water (e.g., removing butanol from the aqueous fermentation media [18,19], breaking ethanol/water azeotrope [20,21]). This paper is a continuation of our systematic investigations of physico-chemical and thermodynamic properties of 1-n-alkyl-1methylpiperidinium bis[(trifluoromethyl)sulfonyl]imides, abbreviated by [CnC1Pip][NTf2] (where n = 2–6; 2 corresponds to ethyl, 3 to n-propyl, and so on) [22–26]. Limiting activity coefficients of 43 different organic solutes and water in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] were measured by using gas-liquid chromatography (GLC) technique. The data are presented in temperature range T = (308.15 to 358.15) K and compared to similar measurements published by us previously for [C 3 C 1 Pip][NTf 2 ] [22] and [C4C1Pip][NTf2] [26]. Based on the obtained results, the capabilities of the studied ILs in liquid extraction processes involving some
170
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
industrially important mixtures are discussed in terms of effects of structure of the cation on selectivity and capacity at infinite dilution. In particular, piperidinium [NTf2]-based homologous series of ILs is compared in that regard to those based on different cations, namely 1-n-alkyl-1-methylpyrrolidinium, [CnC1Pyr][NTf2], and 1-n-alkyl-3-methylimidazolium, [CnC1Im][NTf2] (where n = 1–8). Finally, applications of some theoretical approaches to represent the measured data are demonstrated, including the linear solvation energy relationship (LSER) and the regular solution theory (RST). In order to apply those approaches, ambient pressure densities of pure ILs were measured at different temperatures. On the basis of the RST analysis, the Hildebrand’s solubility parameters of piperidinium [NTf2]-based ILs were calculated as a function of temperature. 2. Experimental methods 2.1. Chemicals and materials The chemical structures of ILs considered in this work are presented in figure 1. The ionic liquid [C5C1Pip][NTf2] was synthesized in our laboratory. In the synthesis we used the accepted technique of first preparing the halide salt of the appropriate cation (in this case [C5C1Pip][Br]), followed by anion metathesis. Synthesis of [C5C1Pip][Br]. For the synthesis a 500 cm3 flask, equipped with a magnetic stirrer and condenser was used. To a solution of 19.8 g (0.200 mol) 1-methylpiperidine (SAFC, P 98%, used as received) in 100 cm3 of acetonitrile (P.O.Ch., HPLC grade) 31.7 g (0.210 mol) of 1-bromopentane (Aldrich, P 99%) were added. The mixture was stirred at 353 K for 24 h and afterwards the solution was allowed to cool down and the product crystallized. The product was filtered and washed with ethyl acetate and then it was recrystallized from acetonitrile/ethyl acetate mixture (6/1). Crystals were dried in vacuum at 353 K for 24 h yielding 48.5 g (0.193 mol) (97% of theoretical value). The following information was determined for our sample: 1 H NMR (400 MHz, D2O): dðppmÞ ¼ 0:954 (t, 3H, 3 J HH ¼ 6:8 Hz), 1:417 (m, 4H), 1:719 (o, 2H, 3 J HH ¼ 6:0 Hz), 1:810 (m, 2H), 1:924 (m, 4H), 3:071 (s, 3H), 3:380 (m, 6H). 13 C NMR (100 MHz, D2O): dðppmÞ ¼ 13:781; 20:219, 21:288; 21:448; 22:221; 28:478; 48:415; 61:739; 64:090. Synthesis of [C5C1Pip][NTf2]. For the synthesis a 250 cm3 flask, equipped with a magnetic stirrer was used. To a solution of 27.9 g (0.111 mol) of [C5C1Pip][Br] (synthesized) in 100 cm3 distilled water, 33.0 g (0.115 mol) of lithium bis[(trifluoromethyl)sulfonyl]imide (Aldrich, P99%) and 25 cm3 of dichloromethane (P.O.Ch., 99.8%) were added. Mixture was stirred at room temperature for 6 h. Two phases were separated. The water phase was extracted with dichloromethane. Afterwards, organic phases were combined and extracted with water until the water phase has given a negative response for AgNO3 aqueous solution. The product
FIGURE 1. The chemical structure of piperidinium ionic liquids investigated in this work.
was dried in vacuum at 353 K for 24 h yielding 49.1 g (0.109 mol) of [C5C1Pip][NTf2] (98% of theoretical value). The following information was determined for our sample: 1 H NMR (400 MHz, CDCl3): dðppmÞ ¼ 0:785 (t, 3H, 3 J HH ¼ 6:8 Hz), 1:235 (m, 4H), 1:600 (m, 4H), 1:741 (m, 4H), 2:869 (s, 3H), 3:125 (m, 2H), 3:195 (t, 4H, 3 J HH ¼ 6:0 Hz). 13 C NMR (100 MHz, CDCl3): dðppmÞ ¼ 13:231; 19:528; 20:258; 20:980; 21:643; 27:757; 46:966; 60:849; 64:021; 119:417 (q, 1 J CF ¼ 321 Hz). Elementary microanalysis. Found: C 34.74%, H 5.36%, N 6.33%, S 14.09%; theoretical values: C 34.66%, H 5.37%, N 6.32%, S 14.17%. The ionic liquid [C6C1Pip][NTf2] was purchased from IoLiTec. The supplied sample had a purity of P 99:9%. In order to remove any volatile chemicals and water this compound was maintained at a very low pressure of about 5 103 Pa at a temperature of about 353 K for 24 h. All solutes employed in GLC measurements were used without further purification. This is due to their high purity and the fact that the gas-liquid chromatography technique separates any impurities on the chromatography column. The list of solutes including source and purities is presented in MS Word Document provided as Supplementary Material for this article. 2.2. Water content The water content was analyzed by the Karl–Fischer titration technique (method TitroLine KF). The samples of ILs (about 3 cm3) were dissolved in methanol and titrated with steps of 2:5 103 cm3 . The results obtained have shown the water content to be less than 500 ppm in the case of both [C5C1Pip][NTf2] and [C6C1Pip][NTf2]. 2.3. GLC apparatus and experimental procedure The Perkin–Elmer Clarus 500 gas chromatograph equipped with a thermal conductivity detector (TCD) was used. The column preparation and the packing method used in this work was performed according to procedures similar to those described by us previously [22,26]. Chromosorb W/AW-DCMS 100/120 mesh supplied by Sigma–Aldrich was used as the solid support. The methanol solution of the IL was used with dispersed Chromosorb inside. After coating, the solvent was evaporated using a rotary evaporator and the Chromosorb coated with IL was placed in glass columns of length 1 m and internal diameter 4 mm. The masses of the stationary phase and of the solid support were weighed with a precision ±0.0001 g, achieving an uncertainty in IL’s loading on the column on the order of 2 104 mmol. The solvent column mass percent packings were 44.97% (4.699 mmol) and 54.65% (7.012 mmol) of [C5C1Pip][NTf2] and 43.72% (4.320 mmol) and 54.65% (6.762 mmol) of [C6C1Pip][NTf2]. Such large column packings were used to prevent the residual adsorption of solute onto the solid support. The inlet pressure of carrier gas (helium) was measured by a pressure gauge installed on the gas chromatograph with an uncertainty of ±0.1 kPa and the outlet pressure was measured using an Agilent Precision Gas Flow Meter having an uncertainty of ± 0.07 kPa. The overpressure applied varied from 15 to 60 kPa, depending on the polarity of the solute and temperature. The flow rate of carrier gas was determined using the same Agilent Precision Gas Flow Meter which was placed at the outlet after the detector and had an uncertainty of ±0.1 cm3 min1. The flow rate was set for a series of runs and was allowed to stabilize for at least 15 minutes before any determinations were made. Solute injections ranged from (0:01 0:3)103 cm3 and such amounts can be considered to be at infinite dilution on the column. The temperature of the column was maintained constant to within 0:02 K. Each experiment was repeated two to three times to establish
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
171
repeatability. Retention times were generally reproducible to within ð103 —102 ) min depending upon the temperature and the individual solute. At each temperature, values of the ‘‘dead time’’, equivalent to the retention time of a completely non-retained component, were also measured. For the TCD, air was used as a nonretainable component. The estimated overall error in limiting activity coefficient was less than 3%, taking into account the possible errors in determining the column loading, the retention times, and solute vapour pressure. Moreover, the values resulting from two different columns prepared were repeatable within errors less than ±1%.
exemplary calculations can be found in supporting information to reference [26]. Limiting activity coefficients are directly linked to excess partial molar enthalpy of mixing, denoted by DH1E;1 . The following relation (called the Gibbs–Helmholtz equation) holds:
2.4. Density measurements
3.2. LSER calculations
The density of the studied ILs was measured using an Anton Paar GmbH 4500 vibrating-tube densimeter (Graz, Austria), thermostated at different temperatures, at atmospheric pressure. Two integrated Pt 100 platinum thermometers providing a good precision of the internal control of temperature (±0.01 K) were installed. The apparatus is precise to within 1 105 gcm3, and the overall uncertainty of the measurements was estimated to be better than 5 105 gcm3. Densimeter calibration was performed at atmospheric pressure using doubly distilled and degassed water, specially-purified benzene (CHEMIPAN, Poland, 99.9%) and dried air.
Linear solvation energy relationship (LSER) methodology was established in 1990s in order to quantify intermolecular solute-stationary phase interactions governing different chromatographic processes [33]. This method allows correlating and predicting various thermodynamic properties related to solute gas-to-liquid transfer processes. The most recent representation of the LSER model for ILs is expressed in terms of decimal logarithm of gasto-liquid partition coefficients (log K L ), the property which can be calculated directly from c1 12 in accordance with the following formula [34]:
DHE;1 @ ln c1 1 12 : ¼ R @T 1 P
If one assumes that DHE;1 does not vary with temperature, then it 1 can be easily determined from the linear regression of ln c1 12 data vs. T 1 .
log K L ¼ log 3. Theory
As in our previous papers [22,26], the methodology proposed and developed by Everett [27] and Cruickshank [28] was adopted to calculate the limiting activity coefficients of organic solutes and water in ILs (c1 12 ). This method uses the truncated virial equation of state to describe the behaviour of the gaseous phase over the stationary IL phase. The working equation reads as follows
n2 RT P01 J 23 U o t0R
P 01 ðB11 V 1 Þ Po J 32 ð2B13 V 1 1 Þ þ ; RT RT
ð1Þ
where subscripts 1, 2, and 3 correspond to solute, solvent (in this case IL) and carrier gas, respectively. Other quantities occurring in 1 are as follows: n2 , number of moles of solvent on the column packing; R, the universal gas constant (R ¼ 8:3144621 JK1 mol1); T, the column temperature; U o , column outlet flow rate; t 0R , corrected retention time of a solute defined as the retention time minus the retention time of a completely unretained solute; P01 , the vapour pressure of pure solute; B11 and B13 , the second virial coefficients; V 1 , the molar volume of pure solute; V 1 1 , the partial molar volume of solute at infinite dilution; P o , the outlet pressure of carrier gas; J 32 , pressure-correction term (the James–Martin coefficient)
J 32 ¼
2 ðPi =Po Þ3 1 ; 3 ðPi =Po Þ2 1
RT q2 ; 0 c1 12 P 1 M 2
ð4Þ
where M2 and q2 stand for molar mass and specific density of solvent (IL). In terms of the LSER theory
3.1. Limiting activity coefficients
ln c1 12 ¼ ln
ð3Þ
log K L ¼ c þ eE þ sS þ aA þ bB þ lL;
where the capital letters correspond to solute-specific descriptors and the lower-case letters represent the respective complementary properties of the IL. The descriptors account for the general solute characteristics: A and B are measures of the solute hydrogen-bond acidity and basicity, respectively, E is the solute excess molar refraction, S is the solute dipolarity/polarizability descriptor, L is the logarithm of the gas-to-hexadecane partition coefficient at T ¼ 298:15 K. The LSER descriptors of solutes considered in this work are listed in MS Excel Spreadsheet provided as Supplementary Material to this article. More comprehensive list can be found elsewhere [33]. In turn, the coefficients c; e; s; a; b, and l are usually obtained from multiple linear regression of experimental log K L data and they somewhat reflect complementary properties of the solvent phase. In order to improve and extend the predictive capabilities of the LSER theory for ILs, several group contribution (GC) methods for the IL-specific coefficients (c; e; s; a; b and l) have been elaborated [35–40]. Such GC-LSER models are based on ionic [35–37], or functional groups contributions [38–40]. The most current and the most comprehensive version of GC-LSER is the model developed recently by Mutelet et al. [40], who proposed the following temperature-dependent GC approach:
ð2Þ
where P i denotes the inlet pressure of carrier gas and J 23 ¼ 1=J 32 . It should be emphasized that all the temperature-dependent quantities are given at the column temperature T. Critical constants, acentric factors, vapor pressures, liquid molar volumes (it was assumed that V 1 1 V 1 ) and second virial coefficients of pure solutes (B11 ) were taken from DIPPR compilations [29]. The cross second virial coefficients (B13 ) were estimated from Tsonopoulos correlation [30] by using Hudson–McCoubrey combining rules [31,32]. For the solutes considered in this all the relevant information required to use equation 1 as well as some
ð5Þ
log K L ¼ 2:84418 þ
N N N N X X X X ni ci þ ni ei E þ ni si S þ ni ai A i¼1
N N X X ni li L þ n i bi B þ i¼1
!,
i¼1
T:
i¼1
i¼1
ð6Þ
i¼1
In equation (6), the coefficients ei ; si ; ai ; bi ; li correspond to contributions of functional group i to a respective LSER coefficients. Their values were established based on about 7000 log K L data points measured in temperature range T ¼ ð293—396Þ K for numerous solutes in 42 ILs [40]. Finally, ni stands for a number of occurrences of the group i in IL’s cations and anions. The groups defined within the
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
172
TABLE 1 E;1 a Experimental activity coefficients at infinite dilution (c1 12 ) and partial molar enthalpies of mixing at infinite dilution (DH 1 ) for different solutes in ionic liquid [C5C1Pip][NTf2] Solute
n-Hexane 2,2-Dimethylbutane n-Heptane n-Octane 2,2,4-Trimethylpentane n-Nonane n-Decane Cyclopentane Cyclohexane Methylcyclohexane Cycloheptane Cyclooctane 1-Hexene Cyclohexene 1-Heptene 1-Octene 1-Decene 1-Pentyne 1-Hexyne 1-Heptyne 1-Octyne Benzene Toluene Ethylbenzene o-Xylene m-Xylene p-Xylene Methanol Ethanol 1-Propanol 1-Butanol 1-Pentanol Water Thiophene Tetrahydrofuran Methyl tert-butyl ether Methyl tert-amyl ether di-n-Propyl ether di-n-Butyl ether Acetone 2-Pentanone 2-Hexanone a
DHE;1 (kJmol1) 1
T (K) 308.15
318.15
328.15
338.15
348.15
358.15
9.90
9.50 7.94 12.4 16.9 12.5 22.4 29.5 4.50 5.95 7.49 7.49 9.58 5.31 3.63 7.12 9.65 16.5 1.59 2.12 2.81 3.74 0.703 0.919 1.31 1.14 1.24 1.24 1.32 1.58 1.82 2.16 2.37 4.19 0.665 0.627 1.85 2.47 4.13 7.79 0.426 0.611 0.761
8.90 7.70 11.8 15.9 11.8 21.1 27.9 4.24 5.64 7.16 7.13 9.06 5.13 3.53 6.88 9.28 15.9 1.62 2.14 2.84 3.76 0.712 0.955 1.35 1.16 1.28 1.28 1.21 1.46 1.68 1.98 2.20 3.70 0.673 0.641 1.86 2.50 4.05 7.50 0.435 0.629 0.779
8.45 7.34 11.3 15.1 11.4 19.9 26.2 3.98 5.42 6.87 6.82 8.60 5.00 3.39 6.67 8.94 15.5 1.64 2.18 2.87 3.79 0.721 0.975 1.37 1.19 1.31 1.32 1.12 1.35 1.55 1.82 2.03 3.18 0.690 0.653 1.89 2.53 4.02 7.29 0.441 0.643 0.792
8.19 7.04 10.8 14.4 10.9 19.0 25.1 3.82 5.13 6.64 6.54 8.24 4.92 3.30 6.48 8.77 15.3 1.69 2.22 2.91 3.81 0.744 1.01 1.42 1.24 1.36 1.37 1.07 1.27 1.46 1.71 1.94 2.78 0.708 0.667 1.94 2.57 3.98 7.17 0.456 0.666 0.822
7.92 6.86 10.4 13.7 10.6 18.0 23.8 3.72 4.99 6.42 6.26 7.80 4.80 3.21 6.31 8.47 14.6 1.73 2.25 2.94 3.84 0.753 1.05 1.43 1.27 1.39 1.40 1.00 1.18 1.36 1.57 1.79 2.44 0.722 0.682 1.98 2.60 3.96 6.94 0.466 0.682 0.840
13.6 18.2 13.4 24.2 31.9 4.72 6.40 7.80 7.97 10.1 5.52 3.73 7.42 9.90 17.1 1.54 2.09 2.78 3.71 0.696 0.910 1.26 1.06 1.15 1.16 1.45 1.79 2.05 2.30 2.50 5.05 0.646 0.617 1.81 2.44 4.17 8.04 0.398
4.24 3.62 4.74 5.13 4.26 5.34 5.29 4.59 4.55 3.61 4.36 4.71 2.51 2.82 2.95 2.90 2.74 2.04 1.37 1.03 0.618 1.50 2.66 2.34 3.11 3.30 3.35 6.75 7.41 7.35 7.06 6.16 13.2 2.01 1.84 1.58 1.17 0.993 2.66 2.66 2.62 2.37
Standard uncertainties (u) are uðc1 12 Þ ¼ 3%; uðTÞ ¼ 0:02 K.
considered GC-LSER method (in total N ¼ 21 groups) allow representation of a great diversity of cationic and anionic structures. In this work we will test the capabilities for [C5C1Pip][NTf2], [C6C1Pip][NTf2] and another piperidinium cation-based IL reported very recently by Marciniak and Wlazło, namely 1-(2-methoxyethyl)-1-methylpiperidinium bis[(trifluoromethyl)sulfonyl]imide, [C2O1C1Pip][NTf2] [41]. It is important to note that the listed ILs were not used in development of the GC-LSER model and hence, the calculations presented in further text are pure predictions.
ln c1;comb ¼1 12
The Hildebrand–Scatchard regular solution theory (RST) combined with Flory ‘‘combinatorial’’ equation is applied in this work to reduce the experimental data on c1 12 and to estimate the Hildebrand’s solubility parameters of ILs belonging to [CnC1Pip][NTf2] series (where n = 3–6). Limiting activity coefficients are assumed to be consisted of combinatorial (‘‘comb’’) and residual (‘‘res’’) contributions due to differences in size and interactions of solute and solvent, respectively:
ð7Þ
r1 r1 þ ln ; r2 r2
ð8Þ
where the size parameters of solute (r 1 ) and IL (r 2 ) are assumed to be proportional to molar volume. Thus, r 1 =r 2 ¼ V 1 =V 2 ¼ ðM 1 q2 Þ=ðM2 q1 Þ. The residual term is given by the RST:
ln c1;res ¼ v12 ¼ 12
3.3. The regular solution theory
1;comb ln c1 þ ln c1;res : 12 ¼ ln c12 12
Combinatorial term is usually calculated by using the well-known Flory equation:
V1 ðd1 d2 Þ2 ; RT
ð9Þ
where symbol v12 denotes the binary interaction parameter, whereas d1 and d2 are the Hildebrand’s solubility parameter of solute and solvent, respectively. After applying some simple algebra, equation (9) can be rearranged as follows:
Y
d21 v12 2d2 d2 ¼ d1 2 : RT V 1 RT RT
ð10Þ
This equation shows that for a given solvent at fixed temperature T, a linear relation holds between Y and the solute solubility parameter, d1 . Then, the solubility parameter of the solvent d2 , is easily determined either from the slope or intercept of this line. The agreement of both d2 values confirms the applicability of the method to
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
173
TABLE 2 E;1 a Experimental activity coefficients at infinite dilution (c1 12 ) and partial molar enthalpies of mixing at infinite dilution (DH 1 ) for different solutes in ionic liquid [C6C1Pip][NTf2]. Solute
n-Pentane n-Hexane 2,2-Dimethylbutane n-Heptane n-Octane 2,2,4-Trimethylpentane n-Nonane n-Decane Cyclopentane Cyclohexane Methylcyclohexane Cycloheptane Cyclooctane 1-Hexene Cyclohexene 1-Heptene 1-Octene 1-Decene 1-Pentyne 1-Hexyne 1-Heptyne 1-Octyne Benzene Toluene Ethylbenzene o-Xylene m-Xylene p-Xylene Methanol Ethanol 1-Propanol 1-Butanol 1-Pentanol Water Thiophene Tetrahydrofuran Methyl tert-butyl ether Methyl tert-amyl ether di-n-Propyl ether di-n-Butyl ether Acetone 2-Pentanone 2-Hexanone a
DHE;1 (kJmol1) 1
T (K) 308.15
318.15
328.15
338.15
348.15
358.15
5.95 7.67 6.67 10.2 12.9 9.75 16.9 23.5 3.72 4.99 6.24 6.23 7.57 4.53 3.07 5.89 7.72 12.3 1.42 1.85 2.39 3.10 0.669 0.864 1.22 1.04 1.14 1.14 1.51 1.81 2.05 2.40 2.44 5.31 0.614 0.563 1.59 2.11 3.50 6.38 0.421 0.573 0.694
5.73 7.40 6.44 9.81 12.7 9.59 16.4 22.2 3.59 4.80 6.05 6.01 7.36 4.45 3.00 5.81 7.60 12.1 1.45 1.88 2.43 3.14 0.687 0.894 1.24 1.08 1.16 1.17 1.37 1.63 1.83 2.13 2.23 4.46 0.627 0.579 1.63 2.14 3.47 6.25 0.429 0.581 0.706
5.42 7.07 6.12 9.38 12.1 9.20 15.7 21.1 3.43 4.60 5.83 5.77 7.15 4.34 2.94 5.69 7.49 12.0 1.48 1.92 2.46 3.19 0.695 0.909 1.26 1.10 1.19 1.20 1.26 1.49 1.67 1.93 2.05 3.81 0.643 0.594 1.66 2.18 3.44 6.08 0.436 0.596 0.722
5.30 6.85 6.00 8.99 11.8 8.95 15.3 20.0 3.34 4.43 5.59 5.52 6.87 4.28 2.88 5.59 7.37 11.8 1.50 1.93 2.49 3.21 0.710 0.934 1.29 1.13 1.23 1.24 1.16 1.37 1.53 1.76 1.92 3.26 0.655 0.612 1.70 2.20 3.41 5.96 0.444 0.613 0.740
5.20 6.67 5.84 8.68 11.3 8.59 14.6 19.1 3.24 4.28 5.40 5.32 6.75 4.15 2.81 5.45 7.27 11.7 1.53 1.96 2.52 3.24 0.722 0.957 1.31 1.16 1.27 1.28 1.08 1.27 1.43 1.64 1.80 2.80 0.664 0.627 1.72 2.22 3.38 5.81 0.455 0.632 0.764
4.95 6.37 5.65 8.20 10.9 8.43 14.1 18.3 3.12 4.06 5.20 5.08 6.53 4.06 2.75 5.35 7.20 11.5 1.54 1.98 2.55 3.26 0.730 0.969 1.32 1.18 1.29 1.30 0.995 1.18 1.33 1.52 1.71 2.45 0.675 0.635 1.76 2.25 3.37 5.74 0.460 0.646 0.784
3.24 3.33 3.01 3.93 3.18 2.84 3.35 4.61 3.18 3.70 3.39 3.74 2.72 2.02 2.01 1.80 1.31 1.19 1.53 1.24 1.17 0.928 1.60 2.12 1.53 2.29 2.42 2.51 7.56 7.81 7.87 8.31 6.54 14.2 1.75 2.29 1.82 1.16 0.728 2.02 1.67 2.30 2.27
Standard uncertainties (u) are uðc1 12 Þ ¼ 3%; uðTÞ ¼ 0:02 K.
the considered system. Binary interaction parameter v12 present in equation 10 is determined for each solute on the basis of experimental c1 12 data, see equation (7) and (8). In turn, the d1 is calculated based on molar enthalpy of vapourization (Dvl H1 ):
d1 ¼
v 1=2 Dl H1 RT : V1
ð11Þ
In this work, DIPPR correlations [29] were used to calculate the values of Dvl H1 and V 1 for each solute as a function of temperature.
4. Results and discussion 4.1. Limiting activity coefficients Tables 1 and 2 list the obtained values of c1 12 for various organic solutes and water in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] in the temperature range T = (308.15–358.15) K and the average partial molar excess enthalpies at infinite dilution (DHE;1 1 ) determined from Gibbs–Helmholtz equation, see equation 3. The resulting temperature dependence of c1 12 for some representative solutes are plotted in Figs. 2 and 3.
For a given IL, the c1 12 increases with an increase of the solute alkyl chain in most of the group of studied compounds. The values of c1 12 for different families of hydrocarbons decrease in the following order: n-alkanes > cyclic alkanes > 1-alkenes > 1-alkynes > aromatics. For isomeric compounds it was observed that changing a structure from linear to branched or cyclic results in lowering the c1 12 . Alkynes, aromatic hydrocarbons and polar compounds (like ketones, thiophene and tetrahydrofuran) reveal negative values of DHE;1 1 . This fact indicates that solute-solvent interactions (e.g., via polarizable p-electrons in aromatic compounds, or polar carbonyl group in ketones, or oxygen group in ethers) are stronger and more favourable than those between solute molecules. This is also a typical behaviour for other reported piperidinium ILs based on [NTf2] [22,26,41], thiocyanate [SCN] [42] and tris[(pentafluoro) ethyl]phosphate [FAP] [43] anion. Having the experimental data presented in this work and the complementary data published previously [22,26], an impact of the cation alkyl side chain in piperidinium [NTf2]-based ILs can be discussed. As it is clearly seen in figure 4, the c1 12 of nonpolar compounds (represented in this figure by n-hexane and cyclohexane) in [CnC1Pip][NTf2] decreases as n increases (at constant temperature T = 308.15 K). Similar observations were made for aromatic hydrocarbons, polar compounds and alcohols,
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
174
15
30
10
20
10 5
12
∞
γ
γ
∞
12
5
1
0.5 300
1
320
340
0.5
360
3
T/K FIGURE 2. Experimental limiting activity coefficients (c1 12 ) of some hydrocarbons in [C5C1Pip][NTf2] (empty markers) and [C6C1Pip][NTf2] (filled markers) as a function of temperature (T): n-hexane (circles), 1-hexene (squares), 1-hexyne (triangles), 1 benzene (diamonds). Solid lines designated by linear regression of ln c1 , see 12 vs. T equation 3.
T (K)
4 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15 358.15
2
γ
∞
6
FIGURE 4. An influence of the length of the piperidinium cation alkyl side chain (n) on limiting activity coefficient (c1 12 ) of some representative compounds at T = 308.15 K: n-hexane (circles), cyclohexane (squares), benzene (triangles), ethanol (diamonds), water (downward-pointing triangles), thiophene (asterisks). Dashed lines serve as guides for the eye.
6
12
5
TABLE 3 Density (q) of [C5C1Pip][NTf2] and [C6C1Pip][NTf2] at atmospheric pressure (P ¼ 0:1 MPa) as a function temperature (T).a,b
10 8
1
0.5 0.3 300
4
n in [Cn C1 Pip][NTf2]
320
340
360
T/ K
a b
FIGURE 3. Experimental limiting activity coefficients (c of some polar solutes in [C5C1Pip][NTf2] (empty markers) and [C6C1Pip][NTf2] (filled markers) as a function of temperature (T): ethanol (circles), water (squares), thiophene (triangles), di-nbutyl ether (diamonds), acetone (downward-pointing triangles). Solid lines desig1 nated by linear regression of ln c1 , see equation 3. 12 vs. T 1 12 )
q (kgm3) [C5C1Pip][NTf2]
[C6C1Pip][NTf2]
1354.47 (1356)b 1350.13 1345.85 1341.70 1337.54 1333.36 1329.20 1325.05 1320.89 1316.75 1312.60 1308.46 1304.36 1300.26
1329.53 (1329)b 1325.41 1321.30 1317.20 1313.11 1309.04 1304.98 1300.93 1296.90 1292.88 1288.87 1284.87 1280.88 1276.91
Standard uncertainties (u) are uðqÞ ¼ 0:05 kg m3 , uðTÞ ¼ 0:02 K. Reference [44].
T = (293.15–358.15) K are listed in Table 3. The experimental data can be accurately described by using the following equations
ln q=kg m3 ¼ 7:20799 6:273 104 ðT=K 298:15Þ exemplified in figure 4 by benzene, thiophene and ethanol, respectively. In the case of water, however, the reverse tendency was revealed. Those results confirm the well-known fact, that more hydrophobic character of the cation promotes attractive van der Waals dispersion interactions between the cation side chains and nonpolar groups present in solute molecules. Finally, it should be stated that the observed regularity of the variation of c1 12 with n indicates a good quality of the measured data as well as their consistency. 4.2. Density measurements The measured densities (at normal pressure) of pure [C5C1Pip][NTf2] and [C6C1Pip][NTf2] in temperature range
ð12Þ
for [C5C1Pip][NTf2] and
ln q=kg m3 ¼ 7:18948 6:213 104 ðT=K 298:15Þ
ð13Þ
for [C6C1Pip][NTf2]. At T ¼ 293:15 K, the results were compared to the values reported recently by Montanino et al. [44]. As seen from Table 3, a fair agreement (within ±0.1%) with those literature data is observed. Finally, a regular variation of molar volume with the piperidinium cation alkyl side chain length was observed when compared the experimental q measured in this work to those measured previously for [C3C1Pip][NTf2] [24] and [C4C1Pip][NTf2] [25]. An increment in molar volume is 17.0 cm3 mol1 per CH2 unit, what is in excellent agreement with the value observed for other families of ILs [45].
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
30
To assess a capability of the investigated ILs as entrainers in liquid–liquid extraction, infinite dilution selectivity (S1 ij ) and infinite 1 dilution capacity (kj ) are taken into account. For separation problem ‘‘solute j from solute i’’ they are defined as follows:
20
S1 ij ¼
c1 i2 c1 j2
1
kj ¼
1
c1 j2
ð14Þ
;
Sij∞ , 10 kj∞
4.3. Selectivity and capacity
175
1
Figure 5 shows S1 ij and kj of the studied ILs (at T = 308.15 K) for a typical pair of ‘‘aromatic from aliphatic’’ separations, i.e., n-hexane (i)/benzene (j). As can be easily noticed, efficiency of the [CnC1Pip][NTf2] ILs is comparable to N-methylpyrrolidone (NMP) [46] and sulfolane [47]. Furthermore, an increase of n deteriorates selectivity and enhances capacity of the solvent. In the case of shorter cation alkyl side chains (n 6 3), the S1 ij is even much better compared to sulfolane [47]. However, one should keep in mind that melting point of such ‘‘short-chain’’ ILs can be relatively high, what can pose a significant obstacle in their applications in real extraction processes (e.g., [C3C1Pip][NTf2] is a liquid at room temperature, while [C2C1Pip][NTf2] is not [23]). In spite of the fact
30
3
4
5
6
8
10
n 1
FIGURE 6. Infinite dilution selectivity (S1 ij ; bars) and capacity (kj ; markers) of [NTf2] anion-based piperidinium and pyrrolidinium ILs for n-hexane (i)/thiophene (j) separation problem, expressed as functions of the cation alkyl side chain length (n) at T = 308.15 K: [CnC1Pip][NTf2] (light gray bars and empty circles), [CnC1Pyr][NTf2] [50] (dark bars and markers). Dashed lines serve as guides for the eye.
that the values of S1 ij and kj are quite similar irrespective of the cation type, piperidinium ILs seem to be slightly more effective solvents than pyrrolidinium [48–50], or imidazolium [48,51] ones. Moreover, an impact of the presence of heteroatoms in alkyl chain can be established based on the results published recently by Marciniak and Wlazło for [C2O1C1Pip][NTf2] [41]. Selectivity and capacity of this IL are 19.8 and 1.11, whereas for the corresponding [C4C1Pip][NTf2] they were 16.4 and 1.27, respectively. Thus, replacing carbon atom by oxygen reduces selectivity, but increases capacity. The same behaviour was observed by Doman´ska and Marciniak for imidazolium [NTf2] ILs [52]. Another very important separation problem discussed here is extraction of sulphur compounds from aliphatic hydrocarbons 1 using [CnC1Pip][NTf2]. Figure 6 presents S1 ij and kj versus n series for n-hexane (i)/thiophene (j) model binary. Comparison to pyrrolidinium ILs reported very recently by Acree et al. [50] is also demonstrated. As seen, the observed trends are exactly the same as those for n-hexane/benzene. The values of selectivity and capacity are, however, much higher and vary from 28.7 and 1.24 for [C3C1Pip][NTf2] to 12.5 and 1.62 for [C6C1Pip][NTf2]. The influence of cationic core is not significant due to evident similarities in piperidinium and pyrrolidinium cations. In fact, thiophene capacity of piperidinium ILs seem to be slighly higher, while the selectivities follow almost exactly the same trend. Furthermore, the presence of oxygen atoms in the cation alkyl chain affects the infty values of S1 in the same manner as described above for ij and kj n-hexane / benzene separations, what is clearly indicated by the results reported by Marciniak and Wlazło [41].
S∞ ij
sulfolane
NMP
1
2
3
4
5
6
7
8
9
10
n 1.8
1.4
kj∞
0
1
20
10
10
NMP
1.0
4.4. LSER correlations and GC-LSER predictions
0.6
sulfolane 0.2
1
2
3
4
5
6
7
8
9
10
n 1
FIGURE 5. Infinite dilution selectivity (S1 ij ; upper panel) and capacity (kj ; lower panel) of [NTf2] anion- based ILs for n-hexane (i)/ benzene (j) separation problem, expressed as functions of the cation alkyl side chain length (n) at T = 308.15 K: [CnC1Pip][NTf2] [22,26] (circles), [CnC1Pyr][NTf2] [49,48,50] (squares), [CnC1Im][NTf2] [51,48] (triangles). Comparison to NMP [46] and sulfolane [47] is highlighted. Dashed lines serve as guides for the eye.
The LSER analysis was carried out at temperature T ¼ 298:15 K. In the first place, the experimental data given in tables 1 and 2 were extrapolated back to T ¼ 298:15 K, assuming a linear depen1 dence of ln c1 . Then, the corresponding log K L values were 12 on T calculated by using equation 4 followed by their multiple linear regression analysis in terms of equation 5. The results are as follows:
log K L ¼ 0:477ð0:065Þ 0:186ð0:099ÞE þ 2:639ð0:123ÞS þ 2:450ð0:119ÞA þ 0:103ð0:112ÞB þ 0:761ð0:018ÞL;
ð15Þ
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
176
ðN ¼ 41;
SD ¼ 0:075;
R2 ¼ 0:9904;
F ¼ 719Þ 4
for [C5C1Pip][NTf2] and
log K L ¼ 0:404ð0:057Þ 0:245ð0:088ÞE þ 2:469ð0:109ÞS þ 2:348ð0:105ÞA þ 0:075ð0:099ÞB
ðN ¼ 42;
SD ¼ 0:066;
ð16Þ
R2 ¼ 0:9917;
F ¼ 831Þ
for [C6C1Pip][NTf2]. In equations (15) and (16), N denotes the number of experimental values, SD refers to the standard deviation, R2 is the squared determination coefficient and F stands for Fisher F-statistic. The standard errors in the calculated coefficients are given in respective parentheses immediately following their values. The very good correlative performance of the LSER model is illustrated in figure 7. The results of GC-LSER [40] predictions are shown in figure 8 as a parity plot. To keep this figure clear, the results were plotted for temperature T = 308.15 K only. However, the standard deviations given in further text include all (i.e., temperature-dependent) data. The experimental values of log K L for [C5C1Pip][NTf2] and [C6C1Pip][NTf2] can be found in MS Excel spreadsheet provided as Supplementary Material to this article. Additionally, the predictions were made for [C2O1C1Pip][NTf2] IL reported by Marciniak and Wlazło [41]. As can be noticed, the GC-LSER method substantially overestimates log K L regardless of the type of solute (nonpolar, polar), or temperature. The standard deviations between calculated (calcd) and experimental (exptl) log K L , defined as
SD ¼
" N X
log K calcd log K exptl L;i L;i
2
#,
!1=2
ðN 6Þ
ð17Þ
i¼1
are 0.471, 0.527 and 0.436 log units for [C5C1Pip][NTf2] (N ¼ 243), [C6C1Pip][NTf2] (N ¼ 252) and [C2O1C1Pip][NTf2] (N ¼ 353), respectively. Those results cannot be viewed as satisfactory, because the applied method describes a huge amount of the available experimental data to within a much lower mean absolute error (about 0.13 log units) [40]. Indeed, it is expected that significantly more experimental data will need to be measured to increase the predictive power of the GC-LSER correlation given in equation 5.
log KLcalcd
3
þ 0:775ð0:016ÞL:
2
1
0
0
1
2
log
3
4
KLexptl
FIGURE 8. Experimental vs. GC-LSER predicted gas–liquid partition coefficients (log K L ) of the studied solutes in piperidinium ILs at T ¼ 308:15 K: [C5C1Pip][NTf2] (circles), [C6C1Pip][NTf2] (squares), [C2O1C1Pip][NTf2] [41] (triangles). Solid line designated by the diagonal. Gray markers show the results at higher temperatures.
We are sure that the new data presented by us in this work will be very helpful in further development of the model. 4.5. Solubility parameters of ILs The Hildebrand’s solubility parameters of [CnC1Pip][NTf2] ILs (d2 ) were determined by means of linear regression of the experimental c1 12 data reported in this paper (n ¼ 5; 6), or published previously (n ¼ 3; 4) [22,26]. The values of d2 were calculated from the slope and from the intercept of the line given in equation 10. An exemplary Y vs. d1 plot is given in figure 9 as an illustration of the adopted method. The d2 of [C6C1Pip][NTf2] determined at T = 308.15 K were ð22:36 0:4Þ (slope) and ð22:55 0:3Þ (intercept) MPa1=2 . Determination coefficient of the fit was R2 ¼ 0:9879. As can be noticed, the difference between d2 estimated by the two methods is relatively small, namely, about ±0.2 MPa1=2 . Thus, the average of these values was taken as a final result.
Y ≡ δ21 / RT − χ12 / V1 (mol ⋅ cm−3)
0.7
calcd log KL
4
3
2
1
1
2
3
log
4
KLexptl
FIGURE 7. Experimental vs. LSER regressed gas-liquid partition coefficients (log K L ) of the studied solutes in piperidinium ILs at T = 298.15 K: [C5C1Pip][NTf2] (empty markers), [C6C1Pip][NTf2] (filled markers). Solid line designated by the diagonal.
0.15
0.5
0.10
0.05
14
16
18
20
0.3
0.1 10
20
30
40
50
1/2
δ1 (MPa ) FIGURE 9. An example of the Hildebrand’s solubility parameter determination by means of the RST and equation 10 (see details in text) for [C6C1Pip][NTf2] at T = 308.15 K. Solid line designated by means of linear regression of Y vs. d1 .
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn TABLE 4 The Hildebrand’s solubility parameters (d2 ) of piperidinium series of [NTf2] ILs a estimated on the basis of linear regression of the c1 12 data by using equation 10. T (K)
d2 (MPa1=2 )
177
at T ¼ 308:15 K). The differences may serve as an semi-quantitative explanation of only partial miscibility of ILs with those solutes. 5. Conclusions
[C3C1Pip][NTf2]b [C4C1Pip][NTf2]c [C5C1Pip][NTf2] [C6C1Pip][NTf2] 308.15 318.15 328.15 338.15 348.15 358.15 a
b c
23.6 23.2 22.9 22.7 22.4 22.1
(0.3) (0.2) (0.2) (0.2) (0.2) (0.3)
23.1 22.8 22.5 22.3 22.0 21.8
(0.2) (0.2) (0.3) (0.3) (0.3) (0.3)
22.8 22.6 22.4 22.2 21.9 21.7
(0.2) (0.2) (0.2) (0.3) (0.3) (0.3)
22.5 22.3 22.1 21.9 21.7 21.5
(0.2) (0.2) (0.2) (0.3) (0.3) (0.3)
The values given in parentheses are the differences in d2 calculated either from the slope and intercept of the line presented in equation 10. c1 12 data taken from reference [22]. c1 12 data taken from reference [26].
24
2
δ (MPa
1/2
)
23
22
21 300
New experimental data on the activity coefficients at infinite dilution of 43 solutes in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] were presented. It was found that the investigated ILs show high selectivity and capacity for n-hexane/benzene and n-hexane/thiophene separation problems. The values of selectivity are comparable with generally used organic solvents such as NMP or sulfolane, and therefore the ILs under consideration can be thought as ‘‘green’’ replacements for these solvents in ‘‘aromatic from aliphatic’’ and ‘‘sulphur from gasolines’’ extraction processes. The results of the GC-LSER predictions of log K L of various solutes clearly indicate that further development of this model is required in order to obtain more reliable and more accurate method. We are quite sure that the new data reported in this work (in total 507 data points) will be very helpful in further development of the model. Finally, the RST was shown to be a robust tool for an estimation of the Hildebrand’s solubility parameters of ILs. The calculations performed for ILs investigated in this work as well as those published previously were combined. Clear trends regarding influence of the piperidinium cation alkyl side chain length and temperature were obtained what confirmed internal consistency of the experimental c1 12 data for [CnC1Pip][NTf2] series (where n = 3–6), which have been measured in our laboratory since 2010. Acknowledgements
320
340
360
Mr. Maciej Zawadzki is greatly acknowledged for technical assistance in synthesis of [C5C1Pip][NTf2] ionic liquid. Funding for this research was provided by the National Science Centre in years 2011–2014 (Grant No. 2011/01/B/ST5/00800).
T (K) Appendix A. Supplementary data FIGURE 10. The Hildebrand’s solubility parameters of piperidinium cation-based [NTf2] ILs (d2 ) estimated on the basis of the RST and equation 10: [C3C1Pip][NTf2] [22] (circles), [C4C1Pip][NTf2] [26] (squares), [C5C1Pip][NTf2] (triangles), [C6C1Pip][NTf2] (diamonds). Dashed lines serve as guides for the eye.
Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.jct.2013.01.005. References
The results of d2 calculations are listed in Table 4 and shown graphically in figure 10. As can be seen, d2 decreases with an increase of temperature. This is a typical behaviour observed for different families of organic ‘‘molecular’’ compounds. The length of the cation alkyl side chain also has a significant influence on d2 . In fact, for longer chains the d2 decreases. It is noteworthy that this tendency is consistent with those known for homologous series of strongly polar and/or associating compounds (e.g., ketones, 1-alcohols) as well as for other ILs [53,54]. It is well-known fact that the Hildebrand’s solubility parameter is a measure of solvation capacity of a given solvent. Generally speaking, solute 1 is expected to be completely miscible with solvent 2, when d1 d2 . Therefore, knowing the values of d1 and d2 facilitates smart solvent screening and selection. In particular, for the ILs studied in this work, d2 varies from 23:5 to 22:4 MPa1=2 for [C3C1Pip][NTf2] and [C6C1Pip][NTf2], respectively, at T ¼ 308:15 K. Those values are quite similar to d1 for alcohols, as can be easily checked by equation 11. For example, d1 ¼ 23:2 MPa1=2 for 1-butanol. Therefore, one may expect that [C3C1Pip][NTf2] is a good candidate for 1-butanol dissolution. In fact, direct solubility measurements confirmed this hypothesis [24]. On the other hand, d2 for ILs are much higher in comparison with d1 for hydrocarbons (e.g., d1 ¼ 14:6 MPa1=2 for n-hexane at T ¼ 308:15 K) and much lower than d1 for water (d1 ¼ 47:5 MPa1=2
[1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19]
M. Freemantle, An Introduction to Ionic Liquids, 1st ed., RCS Publishing, 2009. J.P. Hallet, T. Welton, Chem. Rev. 111 (2011) 3508–3576. N.V. Plechkova, K.R. Seddon, Chem. Soc. Rev. 37 (2008) 123–150. A.B. Pereiro, J.M.M. Araújo, F.S. Oliveira, J.N. Canongia Lopes, I.M. Marrucho, L.P.N. Rebelo, J. Chem. Thermodyn. 55 (2012) 29–36. M.E. Zakrzewska, E. Bogel-Łukasik, R. BogelŁukasik, Energy Fuels 24 (2010) 737–745. H.-J. Huang, S. Ramaswamy, U.W. Tschirner, B.V. Ramarao, Sep. Purif. Technol. 62 (2008) 1–21. A. Marciniak, Fluid Phase Equilib. 294 (2010) 213–233. A.B. Pereiro, J.M.M. Araújo, J.M.S.S. Esperanca, I.M. Marrucho, L.P.N. Rebelo, J. Chem. Thermodyn. 46 (2012) 2–28. U. Doman´ska, M. Królikowski, Fluid Phase Equilib. 259 (2007) 173–179. A. Arce, M.J. Earle, S.P. Katdare, H. Rodrı´guez, K.R. Seddon, Phys. Chem. Chem. Phys. 10 (2008) 2538–2542. G.W. Meindersma, A.R. Hansmeier, A.B. De Haan, Ind. Eng. Chem. Res. 49 (2010) 7530–7540. U. Doman´ska, M. Królikowski, K. S´lesin´ska, J. Chem. Thermodyn. 41 (2009) 1303–1311. A. Seeberger, A. Jess, Green Chem. 12 (2010) 602–608. M. Francisco, A. Arce, A. Soto, Fluid Phase Equilib. 294 (2010) 39–48. A. Arce, M. Francisco, A. Soto, L. Alonso, J. Chem. Eng. Data 55 (2010) 3262– 3267. K. Kedra-Królik, F. Mutelet, J.-C. Moı¨se, J.-N. Jaubert, Energ. Fuels 25 (2011) 1559–1565. S. Lago, H. Rodrı´guez, M.K. Khoshkbarchi, A. Soto, A. Arce, RCS Adv. 2 (2012) 9392–9397. U. Doman´ska, M. Królikowski, J. Chem. Thermodyn. 53 (2012) 108–113. L.Y. Garcia-Chavez, C.M. Garsia, B. Schuur, A.B. De Haan, Ind. Eng. Chem. Res. 51 (2012) 8293–8301.
178
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
[20] N. Calvar, B. González, E. Gómez, A. Dominguez, J. Chem. Thermodyn. 53 (2008) 820–825. [21] X.-M. Li, C. Shen, C.-X. Li, J. Chem. Thermodyn. 53 (2012) 167–175. [22] U. Doman´ska, K. Paduszyn´ski, J. Chem. Thermodyn. 42 (2010) 1361–1366. [23] U. Doman´ska, M. Królikowska, K. Paduszyn´ski, Fluid Phase Equilib. 303 (2011) 1–9. [24] K. Paduszyn´ski, J. Chiyen, D. Ramjugernath, T. Letcher, U. Doman´ska, Fluid Phase Equilib. 305 (2011) 43–52. [25] K. Paduszyn´ski, U. Doman´ska, J. Phys. Chem. B 115 (2011) 12537–12548. [26] K. Paduszyn´ski, U. Doman´ska, J. Phys. Chem. B 115 (2011) 8207–8215. [27] D.H. Everett, Trans. Faraday Soc. 61 (1965) 1637–1645. [28] A.J.B. Cruickshank, B.W. Gainey, C.P. Hicks, T.M. Letcher, R.W. Moody, C.L. Young, Trans. Faraday Soc. 65 (1969) 1014–1031. [29] DIPPR Project 801 – Full Version. Design Institute for Physical Property Research/AIChE, (2005; 2008–2010). Available from:
[38] A.-L. Revelli, F. Mutelet, J.-N. Jaubert, Ind. Eng. Chem. Res. 49 (2010) 3883– 3892. [39] L.M. Grubbs, S. Ye, M. Saifullah, M.C. McMillan-Wiggins, W.E. Acree Jr., M.H. Abraham, P. Twu, J.L. Anderson, Fluid Phase Equilib. 301 (2011) 257–266. [40] F. Mutelet, V. Ortega-Villa, J.-C. Moı¨se, J.-N. Jaubert, W. Acree Jr., J. Chem. Eng. Data 56 (2011) 3598–3606. [41] A. Marciniak, M. Wlazło, J. Chem. Thermodyn. 49 (2012) 137–145. [42] U. Doman´ska, M. Królikowska, J. Chem. Eng. Data 56 (2011) 124–129. [43] A. Marciniak, M. Wlazło, J. Chem. Thermodyn. 57 (2013) 197–202. [44] M. Montanino, M. Carewska, F. Alessandrini, S. Passerini, G.B. Appetecchi, Electrochim. Acta 57 (2011) 153–159. [45] K. Paduszyn´ski, U. Doman´ska, Ind. Eng. Chem. Res. 51 (2012) 591–604. [46] M. Krummen, D. Gruber, J. Gmehling, Ind. Eng. Chem. Res. 39 (2000) 2114– 2123. [47] C. Mölmann, J. Gmehling, J. Chem. Eng. Data 42 (1997) 35–40. [48] R. Kato, J. Gmehling, J. Chem. Thermodyn. 37 (2005) 603–619. [49] S. Nebig, V. Liebert, J. Gmehling, Fluid Phase Equilib. 227 (2009) 61–67. [50] W.E. Acree Jr., G.A. Baker, A.-L. Revelli, J.-C. Moı¨se, F. Mutelet, J. Chem. Eng. Data, J. Chem. Eng. Data 57 (2012) 3510–3518. [51] M. Krummen, P. Wasserscheid, J. Gmehling, J. Chem. Eng. Data 47 (2002) 1411–1417. [52] U. Doman´ska, A. Marciniak, Fluid Phase Equilib. 286 (2009) 154–161. [53] A. Marciniak, Int. J. Mol. Sci. 11 (2010) 1973–1990. [54] F. Mutelet, J.-N. Jaubert, J. Chem. Thermodyn. 39 (2007) 1144–1150.
JCT 12-659
Contents lists available at SciVerse ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Experimental and theoretical study on infinite dilution activity coefficients of various solutes in piperidinium ionic liquids Kamil Paduszyn´ski ⇑, Urszula Doman´ska Department of Physical Chemistry, Faculty of Chemistry, Warsaw University of Technology, Noakowskiego 3, 00-664 Warsaw, Poland
a r t i c l e
i n f o
Article history: Received 18 November 2012 Received in revised form 2 January 2013 Accepted 4 January 2013 Available online 26 January 2013 Keywords: Piperidinium ionic liquids Activity coefficients Extraction LSER models Regular solution theory Solubility parameters
a b s t r a c t The aim of this work is to summarize our systematic studies on activity coefficients at infinite dilution (c1 12 ) of various organic compounds and water in 1-n-alkyl-1-methylpiperidinium cation-based ionic liquids with bis[(trifluoromethyl)sulfonyl]imide anion, [CnC1Pip][NTf2] (n = 3–6). New sets of experimental data measured by gas-liquid chromatography are reported for 43 different solutes in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] in temperature range T ¼ ð308:15 to358:15Þ K. Moreover, ambient pressure liquid densities of those ionic liquids are presented at different temperatures. The results are discussed in terms of an influence of structure of both solute and ionic liquid on c1 12 . Capabilities of the studied ionic liquids in liquid-liquid extraction of aromatics or sulphur compounds from aliphatic hydrocarbons are demonstrated in terms of infinite dilution selectivity and capacity. Finally, linear solvation energy relationship (LSER) theory and regular solution theory (RST) were used to analyze the obtained experimental data. In particular, correlative and predictive power of the most current version of the LSER is demonstrated and the Hildebrand’s solubility parameters are calculated as a function of temperature by using the RST. Ó 2013 Elsevier Ltd. All rights reserved.
1. Introduction Ionic liquids (ILs) are organic salts forming a melt below 100 °C [1]. Peculiarly low solid–liquid phase transition temperature observed for this family of chemicals is closely related to overall asymmetry in the size, shape and chemical structure of ions composing a salt. The cations constituting common ILs are based mainly on functionalized and quaternarized organic bases (e.g., imidazoles, pyridines, amines), whilst the anions are usually smaller, more symmetric and inorganic or organic (e.g., halides, tetrafluoroborate, hexafluorophosphate, alkylsulphates, dialkylphosphates). Those disparities weaken the electrostatic interactions between the cations and anions and thus, substantially lower the value of lattice energy of an ionic liquid compared to ‘‘traditional’’ inorganic salts like NaCl. Since over twenty years ILs have been recognized as modern and clean (‘‘green’’) replacements for volatile organic compounds (VOCs) in many areas of pure and applied chemical sciences, e.g., synthesis and catalysis [2]. In particular, incorporating ILs as environmentally friendly entrainers for sustainable units and processes of chemical and petrochemical technologies seems to be one of their the most promising practical utilities [3]. This is mainly due to extrelemely low volatility, good thermal and chemical stability, enhanced capability of dissolving other substances (ranging from ⇑ Corresponding author. Tel.: +48 222345640. E-mail address: [email protected] (K. Paduszyn´ski). 0021-9614/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2013.01.005
traditional inorganic salts [4] to chemically simple and more complex components of biomaterials, e.g., carbohydrates [5]) and tailorable properties of ILs. Liquid–liquid extraction and azeotropic distillation of close boiling mixtures are the most important and the most widely studied examples of the possible chemical engineering applications of ILs [6–8]. In particular, several different separation problems have been investigated in literature. The most common and representative ones are: aromatic from aliphatic hydrocarbons [9–11], sulphur and nitrogen compounds (e.g., thiophene and pyridine) from hydrocarbons (model mixtures for design of novel processes of extractive desulphurisation/denitrification of gasolines and oils) [12–17] and short-chain alcohols from water (e.g., removing butanol from the aqueous fermentation media [18,19], breaking ethanol/water azeotrope [20,21]). This paper is a continuation of our systematic investigations of physico-chemical and thermodynamic properties of 1-n-alkyl-1methylpiperidinium bis[(trifluoromethyl)sulfonyl]imides, abbreviated by [CnC1Pip][NTf2] (where n = 2–6; 2 corresponds to ethyl, 3 to n-propyl, and so on) [22–26]. Limiting activity coefficients of 43 different organic solutes and water in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] were measured by using gas-liquid chromatography (GLC) technique. The data are presented in temperature range T = (308.15 to 358.15) K and compared to similar measurements published by us previously for [C 3 C 1 Pip][NTf 2 ] [22] and [C4C1Pip][NTf2] [26]. Based on the obtained results, the capabilities of the studied ILs in liquid extraction processes involving some
170
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
industrially important mixtures are discussed in terms of effects of structure of the cation on selectivity and capacity at infinite dilution. In particular, piperidinium [NTf2]-based homologous series of ILs is compared in that regard to those based on different cations, namely 1-n-alkyl-1-methylpyrrolidinium, [CnC1Pyr][NTf2], and 1-n-alkyl-3-methylimidazolium, [CnC1Im][NTf2] (where n = 1–8). Finally, applications of some theoretical approaches to represent the measured data are demonstrated, including the linear solvation energy relationship (LSER) and the regular solution theory (RST). In order to apply those approaches, ambient pressure densities of pure ILs were measured at different temperatures. On the basis of the RST analysis, the Hildebrand’s solubility parameters of piperidinium [NTf2]-based ILs were calculated as a function of temperature. 2. Experimental methods 2.1. Chemicals and materials The chemical structures of ILs considered in this work are presented in figure 1. The ionic liquid [C5C1Pip][NTf2] was synthesized in our laboratory. In the synthesis we used the accepted technique of first preparing the halide salt of the appropriate cation (in this case [C5C1Pip][Br]), followed by anion metathesis. Synthesis of [C5C1Pip][Br]. For the synthesis a 500 cm3 flask, equipped with a magnetic stirrer and condenser was used. To a solution of 19.8 g (0.200 mol) 1-methylpiperidine (SAFC, P 98%, used as received) in 100 cm3 of acetonitrile (P.O.Ch., HPLC grade) 31.7 g (0.210 mol) of 1-bromopentane (Aldrich, P 99%) were added. The mixture was stirred at 353 K for 24 h and afterwards the solution was allowed to cool down and the product crystallized. The product was filtered and washed with ethyl acetate and then it was recrystallized from acetonitrile/ethyl acetate mixture (6/1). Crystals were dried in vacuum at 353 K for 24 h yielding 48.5 g (0.193 mol) (97% of theoretical value). The following information was determined for our sample: 1 H NMR (400 MHz, D2O): dðppmÞ ¼ 0:954 (t, 3H, 3 J HH ¼ 6:8 Hz), 1:417 (m, 4H), 1:719 (o, 2H, 3 J HH ¼ 6:0 Hz), 1:810 (m, 2H), 1:924 (m, 4H), 3:071 (s, 3H), 3:380 (m, 6H). 13 C NMR (100 MHz, D2O): dðppmÞ ¼ 13:781; 20:219, 21:288; 21:448; 22:221; 28:478; 48:415; 61:739; 64:090. Synthesis of [C5C1Pip][NTf2]. For the synthesis a 250 cm3 flask, equipped with a magnetic stirrer was used. To a solution of 27.9 g (0.111 mol) of [C5C1Pip][Br] (synthesized) in 100 cm3 distilled water, 33.0 g (0.115 mol) of lithium bis[(trifluoromethyl)sulfonyl]imide (Aldrich, P99%) and 25 cm3 of dichloromethane (P.O.Ch., 99.8%) were added. Mixture was stirred at room temperature for 6 h. Two phases were separated. The water phase was extracted with dichloromethane. Afterwards, organic phases were combined and extracted with water until the water phase has given a negative response for AgNO3 aqueous solution. The product
FIGURE 1. The chemical structure of piperidinium ionic liquids investigated in this work.
was dried in vacuum at 353 K for 24 h yielding 49.1 g (0.109 mol) of [C5C1Pip][NTf2] (98% of theoretical value). The following information was determined for our sample: 1 H NMR (400 MHz, CDCl3): dðppmÞ ¼ 0:785 (t, 3H, 3 J HH ¼ 6:8 Hz), 1:235 (m, 4H), 1:600 (m, 4H), 1:741 (m, 4H), 2:869 (s, 3H), 3:125 (m, 2H), 3:195 (t, 4H, 3 J HH ¼ 6:0 Hz). 13 C NMR (100 MHz, CDCl3): dðppmÞ ¼ 13:231; 19:528; 20:258; 20:980; 21:643; 27:757; 46:966; 60:849; 64:021; 119:417 (q, 1 J CF ¼ 321 Hz). Elementary microanalysis. Found: C 34.74%, H 5.36%, N 6.33%, S 14.09%; theoretical values: C 34.66%, H 5.37%, N 6.32%, S 14.17%. The ionic liquid [C6C1Pip][NTf2] was purchased from IoLiTec. The supplied sample had a purity of P 99:9%. In order to remove any volatile chemicals and water this compound was maintained at a very low pressure of about 5 103 Pa at a temperature of about 353 K for 24 h. All solutes employed in GLC measurements were used without further purification. This is due to their high purity and the fact that the gas-liquid chromatography technique separates any impurities on the chromatography column. The list of solutes including source and purities is presented in MS Word Document provided as Supplementary Material for this article. 2.2. Water content The water content was analyzed by the Karl–Fischer titration technique (method TitroLine KF). The samples of ILs (about 3 cm3) were dissolved in methanol and titrated with steps of 2:5 103 cm3 . The results obtained have shown the water content to be less than 500 ppm in the case of both [C5C1Pip][NTf2] and [C6C1Pip][NTf2]. 2.3. GLC apparatus and experimental procedure The Perkin–Elmer Clarus 500 gas chromatograph equipped with a thermal conductivity detector (TCD) was used. The column preparation and the packing method used in this work was performed according to procedures similar to those described by us previously [22,26]. Chromosorb W/AW-DCMS 100/120 mesh supplied by Sigma–Aldrich was used as the solid support. The methanol solution of the IL was used with dispersed Chromosorb inside. After coating, the solvent was evaporated using a rotary evaporator and the Chromosorb coated with IL was placed in glass columns of length 1 m and internal diameter 4 mm. The masses of the stationary phase and of the solid support were weighed with a precision ±0.0001 g, achieving an uncertainty in IL’s loading on the column on the order of 2 104 mmol. The solvent column mass percent packings were 44.97% (4.699 mmol) and 54.65% (7.012 mmol) of [C5C1Pip][NTf2] and 43.72% (4.320 mmol) and 54.65% (6.762 mmol) of [C6C1Pip][NTf2]. Such large column packings were used to prevent the residual adsorption of solute onto the solid support. The inlet pressure of carrier gas (helium) was measured by a pressure gauge installed on the gas chromatograph with an uncertainty of ±0.1 kPa and the outlet pressure was measured using an Agilent Precision Gas Flow Meter having an uncertainty of ± 0.07 kPa. The overpressure applied varied from 15 to 60 kPa, depending on the polarity of the solute and temperature. The flow rate of carrier gas was determined using the same Agilent Precision Gas Flow Meter which was placed at the outlet after the detector and had an uncertainty of ±0.1 cm3 min1. The flow rate was set for a series of runs and was allowed to stabilize for at least 15 minutes before any determinations were made. Solute injections ranged from (0:01 0:3)103 cm3 and such amounts can be considered to be at infinite dilution on the column. The temperature of the column was maintained constant to within 0:02 K. Each experiment was repeated two to three times to establish
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
171
repeatability. Retention times were generally reproducible to within ð103 —102 ) min depending upon the temperature and the individual solute. At each temperature, values of the ‘‘dead time’’, equivalent to the retention time of a completely non-retained component, were also measured. For the TCD, air was used as a nonretainable component. The estimated overall error in limiting activity coefficient was less than 3%, taking into account the possible errors in determining the column loading, the retention times, and solute vapour pressure. Moreover, the values resulting from two different columns prepared were repeatable within errors less than ±1%.
exemplary calculations can be found in supporting information to reference [26]. Limiting activity coefficients are directly linked to excess partial molar enthalpy of mixing, denoted by DH1E;1 . The following relation (called the Gibbs–Helmholtz equation) holds:
2.4. Density measurements
3.2. LSER calculations
The density of the studied ILs was measured using an Anton Paar GmbH 4500 vibrating-tube densimeter (Graz, Austria), thermostated at different temperatures, at atmospheric pressure. Two integrated Pt 100 platinum thermometers providing a good precision of the internal control of temperature (±0.01 K) were installed. The apparatus is precise to within 1 105 gcm3, and the overall uncertainty of the measurements was estimated to be better than 5 105 gcm3. Densimeter calibration was performed at atmospheric pressure using doubly distilled and degassed water, specially-purified benzene (CHEMIPAN, Poland, 99.9%) and dried air.
Linear solvation energy relationship (LSER) methodology was established in 1990s in order to quantify intermolecular solute-stationary phase interactions governing different chromatographic processes [33]. This method allows correlating and predicting various thermodynamic properties related to solute gas-to-liquid transfer processes. The most recent representation of the LSER model for ILs is expressed in terms of decimal logarithm of gasto-liquid partition coefficients (log K L ), the property which can be calculated directly from c1 12 in accordance with the following formula [34]:
DHE;1 @ ln c1 1 12 : ¼ R @T 1 P
If one assumes that DHE;1 does not vary with temperature, then it 1 can be easily determined from the linear regression of ln c1 12 data vs. T 1 .
log K L ¼ log 3. Theory
As in our previous papers [22,26], the methodology proposed and developed by Everett [27] and Cruickshank [28] was adopted to calculate the limiting activity coefficients of organic solutes and water in ILs (c1 12 ). This method uses the truncated virial equation of state to describe the behaviour of the gaseous phase over the stationary IL phase. The working equation reads as follows
n2 RT P01 J 23 U o t0R
P 01 ðB11 V 1 Þ Po J 32 ð2B13 V 1 1 Þ þ ; RT RT
ð1Þ
where subscripts 1, 2, and 3 correspond to solute, solvent (in this case IL) and carrier gas, respectively. Other quantities occurring in 1 are as follows: n2 , number of moles of solvent on the column packing; R, the universal gas constant (R ¼ 8:3144621 JK1 mol1); T, the column temperature; U o , column outlet flow rate; t 0R , corrected retention time of a solute defined as the retention time minus the retention time of a completely unretained solute; P01 , the vapour pressure of pure solute; B11 and B13 , the second virial coefficients; V 1 , the molar volume of pure solute; V 1 1 , the partial molar volume of solute at infinite dilution; P o , the outlet pressure of carrier gas; J 32 , pressure-correction term (the James–Martin coefficient)
J 32 ¼
2 ðPi =Po Þ3 1 ; 3 ðPi =Po Þ2 1
RT q2 ; 0 c1 12 P 1 M 2
ð4Þ
where M2 and q2 stand for molar mass and specific density of solvent (IL). In terms of the LSER theory
3.1. Limiting activity coefficients
ln c1 12 ¼ ln
ð3Þ
log K L ¼ c þ eE þ sS þ aA þ bB þ lL;
where the capital letters correspond to solute-specific descriptors and the lower-case letters represent the respective complementary properties of the IL. The descriptors account for the general solute characteristics: A and B are measures of the solute hydrogen-bond acidity and basicity, respectively, E is the solute excess molar refraction, S is the solute dipolarity/polarizability descriptor, L is the logarithm of the gas-to-hexadecane partition coefficient at T ¼ 298:15 K. The LSER descriptors of solutes considered in this work are listed in MS Excel Spreadsheet provided as Supplementary Material to this article. More comprehensive list can be found elsewhere [33]. In turn, the coefficients c; e; s; a; b, and l are usually obtained from multiple linear regression of experimental log K L data and they somewhat reflect complementary properties of the solvent phase. In order to improve and extend the predictive capabilities of the LSER theory for ILs, several group contribution (GC) methods for the IL-specific coefficients (c; e; s; a; b and l) have been elaborated [35–40]. Such GC-LSER models are based on ionic [35–37], or functional groups contributions [38–40]. The most current and the most comprehensive version of GC-LSER is the model developed recently by Mutelet et al. [40], who proposed the following temperature-dependent GC approach:
ð2Þ
where P i denotes the inlet pressure of carrier gas and J 23 ¼ 1=J 32 . It should be emphasized that all the temperature-dependent quantities are given at the column temperature T. Critical constants, acentric factors, vapor pressures, liquid molar volumes (it was assumed that V 1 1 V 1 ) and second virial coefficients of pure solutes (B11 ) were taken from DIPPR compilations [29]. The cross second virial coefficients (B13 ) were estimated from Tsonopoulos correlation [30] by using Hudson–McCoubrey combining rules [31,32]. For the solutes considered in this all the relevant information required to use equation 1 as well as some
ð5Þ
log K L ¼ 2:84418 þ
N N N N X X X X ni ci þ ni ei E þ ni si S þ ni ai A i¼1
N N X X ni li L þ n i bi B þ i¼1
!,
i¼1
T:
i¼1
i¼1
ð6Þ
i¼1
In equation (6), the coefficients ei ; si ; ai ; bi ; li correspond to contributions of functional group i to a respective LSER coefficients. Their values were established based on about 7000 log K L data points measured in temperature range T ¼ ð293—396Þ K for numerous solutes in 42 ILs [40]. Finally, ni stands for a number of occurrences of the group i in IL’s cations and anions. The groups defined within the
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
172
TABLE 1 E;1 a Experimental activity coefficients at infinite dilution (c1 12 ) and partial molar enthalpies of mixing at infinite dilution (DH 1 ) for different solutes in ionic liquid [C5C1Pip][NTf2] Solute
n-Hexane 2,2-Dimethylbutane n-Heptane n-Octane 2,2,4-Trimethylpentane n-Nonane n-Decane Cyclopentane Cyclohexane Methylcyclohexane Cycloheptane Cyclooctane 1-Hexene Cyclohexene 1-Heptene 1-Octene 1-Decene 1-Pentyne 1-Hexyne 1-Heptyne 1-Octyne Benzene Toluene Ethylbenzene o-Xylene m-Xylene p-Xylene Methanol Ethanol 1-Propanol 1-Butanol 1-Pentanol Water Thiophene Tetrahydrofuran Methyl tert-butyl ether Methyl tert-amyl ether di-n-Propyl ether di-n-Butyl ether Acetone 2-Pentanone 2-Hexanone a
DHE;1 (kJmol1) 1
T (K) 308.15
318.15
328.15
338.15
348.15
358.15
9.90
9.50 7.94 12.4 16.9 12.5 22.4 29.5 4.50 5.95 7.49 7.49 9.58 5.31 3.63 7.12 9.65 16.5 1.59 2.12 2.81 3.74 0.703 0.919 1.31 1.14 1.24 1.24 1.32 1.58 1.82 2.16 2.37 4.19 0.665 0.627 1.85 2.47 4.13 7.79 0.426 0.611 0.761
8.90 7.70 11.8 15.9 11.8 21.1 27.9 4.24 5.64 7.16 7.13 9.06 5.13 3.53 6.88 9.28 15.9 1.62 2.14 2.84 3.76 0.712 0.955 1.35 1.16 1.28 1.28 1.21 1.46 1.68 1.98 2.20 3.70 0.673 0.641 1.86 2.50 4.05 7.50 0.435 0.629 0.779
8.45 7.34 11.3 15.1 11.4 19.9 26.2 3.98 5.42 6.87 6.82 8.60 5.00 3.39 6.67 8.94 15.5 1.64 2.18 2.87 3.79 0.721 0.975 1.37 1.19 1.31 1.32 1.12 1.35 1.55 1.82 2.03 3.18 0.690 0.653 1.89 2.53 4.02 7.29 0.441 0.643 0.792
8.19 7.04 10.8 14.4 10.9 19.0 25.1 3.82 5.13 6.64 6.54 8.24 4.92 3.30 6.48 8.77 15.3 1.69 2.22 2.91 3.81 0.744 1.01 1.42 1.24 1.36 1.37 1.07 1.27 1.46 1.71 1.94 2.78 0.708 0.667 1.94 2.57 3.98 7.17 0.456 0.666 0.822
7.92 6.86 10.4 13.7 10.6 18.0 23.8 3.72 4.99 6.42 6.26 7.80 4.80 3.21 6.31 8.47 14.6 1.73 2.25 2.94 3.84 0.753 1.05 1.43 1.27 1.39 1.40 1.00 1.18 1.36 1.57 1.79 2.44 0.722 0.682 1.98 2.60 3.96 6.94 0.466 0.682 0.840
13.6 18.2 13.4 24.2 31.9 4.72 6.40 7.80 7.97 10.1 5.52 3.73 7.42 9.90 17.1 1.54 2.09 2.78 3.71 0.696 0.910 1.26 1.06 1.15 1.16 1.45 1.79 2.05 2.30 2.50 5.05 0.646 0.617 1.81 2.44 4.17 8.04 0.398
4.24 3.62 4.74 5.13 4.26 5.34 5.29 4.59 4.55 3.61 4.36 4.71 2.51 2.82 2.95 2.90 2.74 2.04 1.37 1.03 0.618 1.50 2.66 2.34 3.11 3.30 3.35 6.75 7.41 7.35 7.06 6.16 13.2 2.01 1.84 1.58 1.17 0.993 2.66 2.66 2.62 2.37
Standard uncertainties (u) are uðc1 12 Þ ¼ 3%; uðTÞ ¼ 0:02 K.
considered GC-LSER method (in total N ¼ 21 groups) allow representation of a great diversity of cationic and anionic structures. In this work we will test the capabilities for [C5C1Pip][NTf2], [C6C1Pip][NTf2] and another piperidinium cation-based IL reported very recently by Marciniak and Wlazło, namely 1-(2-methoxyethyl)-1-methylpiperidinium bis[(trifluoromethyl)sulfonyl]imide, [C2O1C1Pip][NTf2] [41]. It is important to note that the listed ILs were not used in development of the GC-LSER model and hence, the calculations presented in further text are pure predictions.
ln c1;comb ¼1 12
The Hildebrand–Scatchard regular solution theory (RST) combined with Flory ‘‘combinatorial’’ equation is applied in this work to reduce the experimental data on c1 12 and to estimate the Hildebrand’s solubility parameters of ILs belonging to [CnC1Pip][NTf2] series (where n = 3–6). Limiting activity coefficients are assumed to be consisted of combinatorial (‘‘comb’’) and residual (‘‘res’’) contributions due to differences in size and interactions of solute and solvent, respectively:
ð7Þ
r1 r1 þ ln ; r2 r2
ð8Þ
where the size parameters of solute (r 1 ) and IL (r 2 ) are assumed to be proportional to molar volume. Thus, r 1 =r 2 ¼ V 1 =V 2 ¼ ðM 1 q2 Þ=ðM2 q1 Þ. The residual term is given by the RST:
ln c1;res ¼ v12 ¼ 12
3.3. The regular solution theory
1;comb ln c1 þ ln c1;res : 12 ¼ ln c12 12
Combinatorial term is usually calculated by using the well-known Flory equation:
V1 ðd1 d2 Þ2 ; RT
ð9Þ
where symbol v12 denotes the binary interaction parameter, whereas d1 and d2 are the Hildebrand’s solubility parameter of solute and solvent, respectively. After applying some simple algebra, equation (9) can be rearranged as follows:
Y
d21 v12 2d2 d2 ¼ d1 2 : RT V 1 RT RT
ð10Þ
This equation shows that for a given solvent at fixed temperature T, a linear relation holds between Y and the solute solubility parameter, d1 . Then, the solubility parameter of the solvent d2 , is easily determined either from the slope or intercept of this line. The agreement of both d2 values confirms the applicability of the method to
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
173
TABLE 2 E;1 a Experimental activity coefficients at infinite dilution (c1 12 ) and partial molar enthalpies of mixing at infinite dilution (DH 1 ) for different solutes in ionic liquid [C6C1Pip][NTf2]. Solute
n-Pentane n-Hexane 2,2-Dimethylbutane n-Heptane n-Octane 2,2,4-Trimethylpentane n-Nonane n-Decane Cyclopentane Cyclohexane Methylcyclohexane Cycloheptane Cyclooctane 1-Hexene Cyclohexene 1-Heptene 1-Octene 1-Decene 1-Pentyne 1-Hexyne 1-Heptyne 1-Octyne Benzene Toluene Ethylbenzene o-Xylene m-Xylene p-Xylene Methanol Ethanol 1-Propanol 1-Butanol 1-Pentanol Water Thiophene Tetrahydrofuran Methyl tert-butyl ether Methyl tert-amyl ether di-n-Propyl ether di-n-Butyl ether Acetone 2-Pentanone 2-Hexanone a
DHE;1 (kJmol1) 1
T (K) 308.15
318.15
328.15
338.15
348.15
358.15
5.95 7.67 6.67 10.2 12.9 9.75 16.9 23.5 3.72 4.99 6.24 6.23 7.57 4.53 3.07 5.89 7.72 12.3 1.42 1.85 2.39 3.10 0.669 0.864 1.22 1.04 1.14 1.14 1.51 1.81 2.05 2.40 2.44 5.31 0.614 0.563 1.59 2.11 3.50 6.38 0.421 0.573 0.694
5.73 7.40 6.44 9.81 12.7 9.59 16.4 22.2 3.59 4.80 6.05 6.01 7.36 4.45 3.00 5.81 7.60 12.1 1.45 1.88 2.43 3.14 0.687 0.894 1.24 1.08 1.16 1.17 1.37 1.63 1.83 2.13 2.23 4.46 0.627 0.579 1.63 2.14 3.47 6.25 0.429 0.581 0.706
5.42 7.07 6.12 9.38 12.1 9.20 15.7 21.1 3.43 4.60 5.83 5.77 7.15 4.34 2.94 5.69 7.49 12.0 1.48 1.92 2.46 3.19 0.695 0.909 1.26 1.10 1.19 1.20 1.26 1.49 1.67 1.93 2.05 3.81 0.643 0.594 1.66 2.18 3.44 6.08 0.436 0.596 0.722
5.30 6.85 6.00 8.99 11.8 8.95 15.3 20.0 3.34 4.43 5.59 5.52 6.87 4.28 2.88 5.59 7.37 11.8 1.50 1.93 2.49 3.21 0.710 0.934 1.29 1.13 1.23 1.24 1.16 1.37 1.53 1.76 1.92 3.26 0.655 0.612 1.70 2.20 3.41 5.96 0.444 0.613 0.740
5.20 6.67 5.84 8.68 11.3 8.59 14.6 19.1 3.24 4.28 5.40 5.32 6.75 4.15 2.81 5.45 7.27 11.7 1.53 1.96 2.52 3.24 0.722 0.957 1.31 1.16 1.27 1.28 1.08 1.27 1.43 1.64 1.80 2.80 0.664 0.627 1.72 2.22 3.38 5.81 0.455 0.632 0.764
4.95 6.37 5.65 8.20 10.9 8.43 14.1 18.3 3.12 4.06 5.20 5.08 6.53 4.06 2.75 5.35 7.20 11.5 1.54 1.98 2.55 3.26 0.730 0.969 1.32 1.18 1.29 1.30 0.995 1.18 1.33 1.52 1.71 2.45 0.675 0.635 1.76 2.25 3.37 5.74 0.460 0.646 0.784
3.24 3.33 3.01 3.93 3.18 2.84 3.35 4.61 3.18 3.70 3.39 3.74 2.72 2.02 2.01 1.80 1.31 1.19 1.53 1.24 1.17 0.928 1.60 2.12 1.53 2.29 2.42 2.51 7.56 7.81 7.87 8.31 6.54 14.2 1.75 2.29 1.82 1.16 0.728 2.02 1.67 2.30 2.27
Standard uncertainties (u) are uðc1 12 Þ ¼ 3%; uðTÞ ¼ 0:02 K.
the considered system. Binary interaction parameter v12 present in equation 10 is determined for each solute on the basis of experimental c1 12 data, see equation (7) and (8). In turn, the d1 is calculated based on molar enthalpy of vapourization (Dvl H1 ):
d1 ¼
v 1=2 Dl H1 RT : V1
ð11Þ
In this work, DIPPR correlations [29] were used to calculate the values of Dvl H1 and V 1 for each solute as a function of temperature.
4. Results and discussion 4.1. Limiting activity coefficients Tables 1 and 2 list the obtained values of c1 12 for various organic solutes and water in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] in the temperature range T = (308.15–358.15) K and the average partial molar excess enthalpies at infinite dilution (DHE;1 1 ) determined from Gibbs–Helmholtz equation, see equation 3. The resulting temperature dependence of c1 12 for some representative solutes are plotted in Figs. 2 and 3.
For a given IL, the c1 12 increases with an increase of the solute alkyl chain in most of the group of studied compounds. The values of c1 12 for different families of hydrocarbons decrease in the following order: n-alkanes > cyclic alkanes > 1-alkenes > 1-alkynes > aromatics. For isomeric compounds it was observed that changing a structure from linear to branched or cyclic results in lowering the c1 12 . Alkynes, aromatic hydrocarbons and polar compounds (like ketones, thiophene and tetrahydrofuran) reveal negative values of DHE;1 1 . This fact indicates that solute-solvent interactions (e.g., via polarizable p-electrons in aromatic compounds, or polar carbonyl group in ketones, or oxygen group in ethers) are stronger and more favourable than those between solute molecules. This is also a typical behaviour for other reported piperidinium ILs based on [NTf2] [22,26,41], thiocyanate [SCN] [42] and tris[(pentafluoro) ethyl]phosphate [FAP] [43] anion. Having the experimental data presented in this work and the complementary data published previously [22,26], an impact of the cation alkyl side chain in piperidinium [NTf2]-based ILs can be discussed. As it is clearly seen in figure 4, the c1 12 of nonpolar compounds (represented in this figure by n-hexane and cyclohexane) in [CnC1Pip][NTf2] decreases as n increases (at constant temperature T = 308.15 K). Similar observations were made for aromatic hydrocarbons, polar compounds and alcohols,
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
174
15
30
10
20
10 5
12
∞
γ
γ
∞
12
5
1
0.5 300
1
320
340
0.5
360
3
T/K FIGURE 2. Experimental limiting activity coefficients (c1 12 ) of some hydrocarbons in [C5C1Pip][NTf2] (empty markers) and [C6C1Pip][NTf2] (filled markers) as a function of temperature (T): n-hexane (circles), 1-hexene (squares), 1-hexyne (triangles), 1 benzene (diamonds). Solid lines designated by linear regression of ln c1 , see 12 vs. T equation 3.
T (K)
4 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15 358.15
2
γ
∞
6
FIGURE 4. An influence of the length of the piperidinium cation alkyl side chain (n) on limiting activity coefficient (c1 12 ) of some representative compounds at T = 308.15 K: n-hexane (circles), cyclohexane (squares), benzene (triangles), ethanol (diamonds), water (downward-pointing triangles), thiophene (asterisks). Dashed lines serve as guides for the eye.
6
12
5
TABLE 3 Density (q) of [C5C1Pip][NTf2] and [C6C1Pip][NTf2] at atmospheric pressure (P ¼ 0:1 MPa) as a function temperature (T).a,b
10 8
1
0.5 0.3 300
4
n in [Cn C1 Pip][NTf2]
320
340
360
T/ K
a b
FIGURE 3. Experimental limiting activity coefficients (c of some polar solutes in [C5C1Pip][NTf2] (empty markers) and [C6C1Pip][NTf2] (filled markers) as a function of temperature (T): ethanol (circles), water (squares), thiophene (triangles), di-nbutyl ether (diamonds), acetone (downward-pointing triangles). Solid lines desig1 nated by linear regression of ln c1 , see equation 3. 12 vs. T 1 12 )
q (kgm3) [C5C1Pip][NTf2]
[C6C1Pip][NTf2]
1354.47 (1356)b 1350.13 1345.85 1341.70 1337.54 1333.36 1329.20 1325.05 1320.89 1316.75 1312.60 1308.46 1304.36 1300.26
1329.53 (1329)b 1325.41 1321.30 1317.20 1313.11 1309.04 1304.98 1300.93 1296.90 1292.88 1288.87 1284.87 1280.88 1276.91
Standard uncertainties (u) are uðqÞ ¼ 0:05 kg m3 , uðTÞ ¼ 0:02 K. Reference [44].
T = (293.15–358.15) K are listed in Table 3. The experimental data can be accurately described by using the following equations
ln q=kg m3 ¼ 7:20799 6:273 104 ðT=K 298:15Þ exemplified in figure 4 by benzene, thiophene and ethanol, respectively. In the case of water, however, the reverse tendency was revealed. Those results confirm the well-known fact, that more hydrophobic character of the cation promotes attractive van der Waals dispersion interactions between the cation side chains and nonpolar groups present in solute molecules. Finally, it should be stated that the observed regularity of the variation of c1 12 with n indicates a good quality of the measured data as well as their consistency. 4.2. Density measurements The measured densities (at normal pressure) of pure [C5C1Pip][NTf2] and [C6C1Pip][NTf2] in temperature range
ð12Þ
for [C5C1Pip][NTf2] and
ln q=kg m3 ¼ 7:18948 6:213 104 ðT=K 298:15Þ
ð13Þ
for [C6C1Pip][NTf2]. At T ¼ 293:15 K, the results were compared to the values reported recently by Montanino et al. [44]. As seen from Table 3, a fair agreement (within ±0.1%) with those literature data is observed. Finally, a regular variation of molar volume with the piperidinium cation alkyl side chain length was observed when compared the experimental q measured in this work to those measured previously for [C3C1Pip][NTf2] [24] and [C4C1Pip][NTf2] [25]. An increment in molar volume is 17.0 cm3 mol1 per CH2 unit, what is in excellent agreement with the value observed for other families of ILs [45].
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
30
To assess a capability of the investigated ILs as entrainers in liquid–liquid extraction, infinite dilution selectivity (S1 ij ) and infinite 1 dilution capacity (kj ) are taken into account. For separation problem ‘‘solute j from solute i’’ they are defined as follows:
20
S1 ij ¼
c1 i2 c1 j2
1
kj ¼
1
c1 j2
ð14Þ
;
Sij∞ , 10 kj∞
4.3. Selectivity and capacity
175
1
Figure 5 shows S1 ij and kj of the studied ILs (at T = 308.15 K) for a typical pair of ‘‘aromatic from aliphatic’’ separations, i.e., n-hexane (i)/benzene (j). As can be easily noticed, efficiency of the [CnC1Pip][NTf2] ILs is comparable to N-methylpyrrolidone (NMP) [46] and sulfolane [47]. Furthermore, an increase of n deteriorates selectivity and enhances capacity of the solvent. In the case of shorter cation alkyl side chains (n 6 3), the S1 ij is even much better compared to sulfolane [47]. However, one should keep in mind that melting point of such ‘‘short-chain’’ ILs can be relatively high, what can pose a significant obstacle in their applications in real extraction processes (e.g., [C3C1Pip][NTf2] is a liquid at room temperature, while [C2C1Pip][NTf2] is not [23]). In spite of the fact
30
3
4
5
6
8
10
n 1
FIGURE 6. Infinite dilution selectivity (S1 ij ; bars) and capacity (kj ; markers) of [NTf2] anion-based piperidinium and pyrrolidinium ILs for n-hexane (i)/thiophene (j) separation problem, expressed as functions of the cation alkyl side chain length (n) at T = 308.15 K: [CnC1Pip][NTf2] (light gray bars and empty circles), [CnC1Pyr][NTf2] [50] (dark bars and markers). Dashed lines serve as guides for the eye.
that the values of S1 ij and kj are quite similar irrespective of the cation type, piperidinium ILs seem to be slightly more effective solvents than pyrrolidinium [48–50], or imidazolium [48,51] ones. Moreover, an impact of the presence of heteroatoms in alkyl chain can be established based on the results published recently by Marciniak and Wlazło for [C2O1C1Pip][NTf2] [41]. Selectivity and capacity of this IL are 19.8 and 1.11, whereas for the corresponding [C4C1Pip][NTf2] they were 16.4 and 1.27, respectively. Thus, replacing carbon atom by oxygen reduces selectivity, but increases capacity. The same behaviour was observed by Doman´ska and Marciniak for imidazolium [NTf2] ILs [52]. Another very important separation problem discussed here is extraction of sulphur compounds from aliphatic hydrocarbons 1 using [CnC1Pip][NTf2]. Figure 6 presents S1 ij and kj versus n series for n-hexane (i)/thiophene (j) model binary. Comparison to pyrrolidinium ILs reported very recently by Acree et al. [50] is also demonstrated. As seen, the observed trends are exactly the same as those for n-hexane/benzene. The values of selectivity and capacity are, however, much higher and vary from 28.7 and 1.24 for [C3C1Pip][NTf2] to 12.5 and 1.62 for [C6C1Pip][NTf2]. The influence of cationic core is not significant due to evident similarities in piperidinium and pyrrolidinium cations. In fact, thiophene capacity of piperidinium ILs seem to be slighly higher, while the selectivities follow almost exactly the same trend. Furthermore, the presence of oxygen atoms in the cation alkyl chain affects the infty values of S1 in the same manner as described above for ij and kj n-hexane / benzene separations, what is clearly indicated by the results reported by Marciniak and Wlazło [41].
S∞ ij
sulfolane
NMP
1
2
3
4
5
6
7
8
9
10
n 1.8
1.4
kj∞
0
1
20
10
10
NMP
1.0
4.4. LSER correlations and GC-LSER predictions
0.6
sulfolane 0.2
1
2
3
4
5
6
7
8
9
10
n 1
FIGURE 5. Infinite dilution selectivity (S1 ij ; upper panel) and capacity (kj ; lower panel) of [NTf2] anion- based ILs for n-hexane (i)/ benzene (j) separation problem, expressed as functions of the cation alkyl side chain length (n) at T = 308.15 K: [CnC1Pip][NTf2] [22,26] (circles), [CnC1Pyr][NTf2] [49,48,50] (squares), [CnC1Im][NTf2] [51,48] (triangles). Comparison to NMP [46] and sulfolane [47] is highlighted. Dashed lines serve as guides for the eye.
The LSER analysis was carried out at temperature T ¼ 298:15 K. In the first place, the experimental data given in tables 1 and 2 were extrapolated back to T ¼ 298:15 K, assuming a linear depen1 dence of ln c1 . Then, the corresponding log K L values were 12 on T calculated by using equation 4 followed by their multiple linear regression analysis in terms of equation 5. The results are as follows:
log K L ¼ 0:477ð0:065Þ 0:186ð0:099ÞE þ 2:639ð0:123ÞS þ 2:450ð0:119ÞA þ 0:103ð0:112ÞB þ 0:761ð0:018ÞL;
ð15Þ
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
176
ðN ¼ 41;
SD ¼ 0:075;
R2 ¼ 0:9904;
F ¼ 719Þ 4
for [C5C1Pip][NTf2] and
log K L ¼ 0:404ð0:057Þ 0:245ð0:088ÞE þ 2:469ð0:109ÞS þ 2:348ð0:105ÞA þ 0:075ð0:099ÞB
ðN ¼ 42;
SD ¼ 0:066;
ð16Þ
R2 ¼ 0:9917;
F ¼ 831Þ
for [C6C1Pip][NTf2]. In equations (15) and (16), N denotes the number of experimental values, SD refers to the standard deviation, R2 is the squared determination coefficient and F stands for Fisher F-statistic. The standard errors in the calculated coefficients are given in respective parentheses immediately following their values. The very good correlative performance of the LSER model is illustrated in figure 7. The results of GC-LSER [40] predictions are shown in figure 8 as a parity plot. To keep this figure clear, the results were plotted for temperature T = 308.15 K only. However, the standard deviations given in further text include all (i.e., temperature-dependent) data. The experimental values of log K L for [C5C1Pip][NTf2] and [C6C1Pip][NTf2] can be found in MS Excel spreadsheet provided as Supplementary Material to this article. Additionally, the predictions were made for [C2O1C1Pip][NTf2] IL reported by Marciniak and Wlazło [41]. As can be noticed, the GC-LSER method substantially overestimates log K L regardless of the type of solute (nonpolar, polar), or temperature. The standard deviations between calculated (calcd) and experimental (exptl) log K L , defined as
SD ¼
" N X
log K calcd log K exptl L;i L;i
2
#,
!1=2
ðN 6Þ
ð17Þ
i¼1
are 0.471, 0.527 and 0.436 log units for [C5C1Pip][NTf2] (N ¼ 243), [C6C1Pip][NTf2] (N ¼ 252) and [C2O1C1Pip][NTf2] (N ¼ 353), respectively. Those results cannot be viewed as satisfactory, because the applied method describes a huge amount of the available experimental data to within a much lower mean absolute error (about 0.13 log units) [40]. Indeed, it is expected that significantly more experimental data will need to be measured to increase the predictive power of the GC-LSER correlation given in equation 5.
log KLcalcd
3
þ 0:775ð0:016ÞL:
2
1
0
0
1
2
log
3
4
KLexptl
FIGURE 8. Experimental vs. GC-LSER predicted gas–liquid partition coefficients (log K L ) of the studied solutes in piperidinium ILs at T ¼ 308:15 K: [C5C1Pip][NTf2] (circles), [C6C1Pip][NTf2] (squares), [C2O1C1Pip][NTf2] [41] (triangles). Solid line designated by the diagonal. Gray markers show the results at higher temperatures.
We are sure that the new data presented by us in this work will be very helpful in further development of the model. 4.5. Solubility parameters of ILs The Hildebrand’s solubility parameters of [CnC1Pip][NTf2] ILs (d2 ) were determined by means of linear regression of the experimental c1 12 data reported in this paper (n ¼ 5; 6), or published previously (n ¼ 3; 4) [22,26]. The values of d2 were calculated from the slope and from the intercept of the line given in equation 10. An exemplary Y vs. d1 plot is given in figure 9 as an illustration of the adopted method. The d2 of [C6C1Pip][NTf2] determined at T = 308.15 K were ð22:36 0:4Þ (slope) and ð22:55 0:3Þ (intercept) MPa1=2 . Determination coefficient of the fit was R2 ¼ 0:9879. As can be noticed, the difference between d2 estimated by the two methods is relatively small, namely, about ±0.2 MPa1=2 . Thus, the average of these values was taken as a final result.
Y ≡ δ21 / RT − χ12 / V1 (mol ⋅ cm−3)
0.7
calcd log KL
4
3
2
1
1
2
3
log
4
KLexptl
FIGURE 7. Experimental vs. LSER regressed gas-liquid partition coefficients (log K L ) of the studied solutes in piperidinium ILs at T = 298.15 K: [C5C1Pip][NTf2] (empty markers), [C6C1Pip][NTf2] (filled markers). Solid line designated by the diagonal.
0.15
0.5
0.10
0.05
14
16
18
20
0.3
0.1 10
20
30
40
50
1/2
δ1 (MPa ) FIGURE 9. An example of the Hildebrand’s solubility parameter determination by means of the RST and equation 10 (see details in text) for [C6C1Pip][NTf2] at T = 308.15 K. Solid line designated by means of linear regression of Y vs. d1 .
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn TABLE 4 The Hildebrand’s solubility parameters (d2 ) of piperidinium series of [NTf2] ILs a estimated on the basis of linear regression of the c1 12 data by using equation 10. T (K)
d2 (MPa1=2 )
177
at T ¼ 308:15 K). The differences may serve as an semi-quantitative explanation of only partial miscibility of ILs with those solutes. 5. Conclusions
[C3C1Pip][NTf2]b [C4C1Pip][NTf2]c [C5C1Pip][NTf2] [C6C1Pip][NTf2] 308.15 318.15 328.15 338.15 348.15 358.15 a
b c
23.6 23.2 22.9 22.7 22.4 22.1
(0.3) (0.2) (0.2) (0.2) (0.2) (0.3)
23.1 22.8 22.5 22.3 22.0 21.8
(0.2) (0.2) (0.3) (0.3) (0.3) (0.3)
22.8 22.6 22.4 22.2 21.9 21.7
(0.2) (0.2) (0.2) (0.3) (0.3) (0.3)
22.5 22.3 22.1 21.9 21.7 21.5
(0.2) (0.2) (0.2) (0.3) (0.3) (0.3)
The values given in parentheses are the differences in d2 calculated either from the slope and intercept of the line presented in equation 10. c1 12 data taken from reference [22]. c1 12 data taken from reference [26].
24
2
δ (MPa
1/2
)
23
22
21 300
New experimental data on the activity coefficients at infinite dilution of 43 solutes in [C5C1Pip][NTf2] and [C6C1Pip][NTf2] were presented. It was found that the investigated ILs show high selectivity and capacity for n-hexane/benzene and n-hexane/thiophene separation problems. The values of selectivity are comparable with generally used organic solvents such as NMP or sulfolane, and therefore the ILs under consideration can be thought as ‘‘green’’ replacements for these solvents in ‘‘aromatic from aliphatic’’ and ‘‘sulphur from gasolines’’ extraction processes. The results of the GC-LSER predictions of log K L of various solutes clearly indicate that further development of this model is required in order to obtain more reliable and more accurate method. We are quite sure that the new data reported in this work (in total 507 data points) will be very helpful in further development of the model. Finally, the RST was shown to be a robust tool for an estimation of the Hildebrand’s solubility parameters of ILs. The calculations performed for ILs investigated in this work as well as those published previously were combined. Clear trends regarding influence of the piperidinium cation alkyl side chain length and temperature were obtained what confirmed internal consistency of the experimental c1 12 data for [CnC1Pip][NTf2] series (where n = 3–6), which have been measured in our laboratory since 2010. Acknowledgements
320
340
360
Mr. Maciej Zawadzki is greatly acknowledged for technical assistance in synthesis of [C5C1Pip][NTf2] ionic liquid. Funding for this research was provided by the National Science Centre in years 2011–2014 (Grant No. 2011/01/B/ST5/00800).
T (K) Appendix A. Supplementary data FIGURE 10. The Hildebrand’s solubility parameters of piperidinium cation-based [NTf2] ILs (d2 ) estimated on the basis of the RST and equation 10: [C3C1Pip][NTf2] [22] (circles), [C4C1Pip][NTf2] [26] (squares), [C5C1Pip][NTf2] (triangles), [C6C1Pip][NTf2] (diamonds). Dashed lines serve as guides for the eye.
Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.jct.2013.01.005. References
The results of d2 calculations are listed in Table 4 and shown graphically in figure 10. As can be seen, d2 decreases with an increase of temperature. This is a typical behaviour observed for different families of organic ‘‘molecular’’ compounds. The length of the cation alkyl side chain also has a significant influence on d2 . In fact, for longer chains the d2 decreases. It is noteworthy that this tendency is consistent with those known for homologous series of strongly polar and/or associating compounds (e.g., ketones, 1-alcohols) as well as for other ILs [53,54]. It is well-known fact that the Hildebrand’s solubility parameter is a measure of solvation capacity of a given solvent. Generally speaking, solute 1 is expected to be completely miscible with solvent 2, when d1 d2 . Therefore, knowing the values of d1 and d2 facilitates smart solvent screening and selection. In particular, for the ILs studied in this work, d2 varies from 23:5 to 22:4 MPa1=2 for [C3C1Pip][NTf2] and [C6C1Pip][NTf2], respectively, at T ¼ 308:15 K. Those values are quite similar to d1 for alcohols, as can be easily checked by equation 11. For example, d1 ¼ 23:2 MPa1=2 for 1-butanol. Therefore, one may expect that [C3C1Pip][NTf2] is a good candidate for 1-butanol dissolution. In fact, direct solubility measurements confirmed this hypothesis [24]. On the other hand, d2 for ILs are much higher in comparison with d1 for hydrocarbons (e.g., d1 ¼ 14:6 MPa1=2 for n-hexane at T ¼ 308:15 K) and much lower than d1 for water (d1 ¼ 47:5 MPa1=2
[1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19]
M. Freemantle, An Introduction to Ionic Liquids, 1st ed., RCS Publishing, 2009. J.P. Hallet, T. Welton, Chem. Rev. 111 (2011) 3508–3576. N.V. Plechkova, K.R. Seddon, Chem. Soc. Rev. 37 (2008) 123–150. A.B. Pereiro, J.M.M. Araújo, F.S. Oliveira, J.N. Canongia Lopes, I.M. Marrucho, L.P.N. Rebelo, J. Chem. Thermodyn. 55 (2012) 29–36. M.E. Zakrzewska, E. Bogel-Łukasik, R. BogelŁukasik, Energy Fuels 24 (2010) 737–745. H.-J. Huang, S. Ramaswamy, U.W. Tschirner, B.V. Ramarao, Sep. Purif. Technol. 62 (2008) 1–21. A. Marciniak, Fluid Phase Equilib. 294 (2010) 213–233. A.B. Pereiro, J.M.M. Araújo, J.M.S.S. Esperanca, I.M. Marrucho, L.P.N. Rebelo, J. Chem. Thermodyn. 46 (2012) 2–28. U. Doman´ska, M. Królikowski, Fluid Phase Equilib. 259 (2007) 173–179. A. Arce, M.J. Earle, S.P. Katdare, H. Rodrı´guez, K.R. Seddon, Phys. Chem. Chem. Phys. 10 (2008) 2538–2542. G.W. Meindersma, A.R. Hansmeier, A.B. De Haan, Ind. Eng. Chem. Res. 49 (2010) 7530–7540. U. Doman´ska, M. Królikowski, K. S´lesin´ska, J. Chem. Thermodyn. 41 (2009) 1303–1311. A. Seeberger, A. Jess, Green Chem. 12 (2010) 602–608. M. Francisco, A. Arce, A. Soto, Fluid Phase Equilib. 294 (2010) 39–48. A. Arce, M. Francisco, A. Soto, L. Alonso, J. Chem. Eng. Data 55 (2010) 3262– 3267. K. Kedra-Królik, F. Mutelet, J.-C. Moı¨se, J.-N. Jaubert, Energ. Fuels 25 (2011) 1559–1565. S. Lago, H. Rodrı´guez, M.K. Khoshkbarchi, A. Soto, A. Arce, RCS Adv. 2 (2012) 9392–9397. U. Doman´ska, M. Królikowski, J. Chem. Thermodyn. 53 (2012) 108–113. L.Y. Garcia-Chavez, C.M. Garsia, B. Schuur, A.B. De Haan, Ind. Eng. Chem. Res. 51 (2012) 8293–8301.
178
´ ski, U. Doman ´ ska / J. Chem. Thermodynamics 60 (2013) 169–178 K. Paduszyn
[20] N. Calvar, B. González, E. Gómez, A. Dominguez, J. Chem. Thermodyn. 53 (2008) 820–825. [21] X.-M. Li, C. Shen, C.-X. Li, J. Chem. Thermodyn. 53 (2012) 167–175. [22] U. Doman´ska, K. Paduszyn´ski, J. Chem. Thermodyn. 42 (2010) 1361–1366. [23] U. Doman´ska, M. Królikowska, K. Paduszyn´ski, Fluid Phase Equilib. 303 (2011) 1–9. [24] K. Paduszyn´ski, J. Chiyen, D. Ramjugernath, T. Letcher, U. Doman´ska, Fluid Phase Equilib. 305 (2011) 43–52. [25] K. Paduszyn´ski, U. Doman´ska, J. Phys. Chem. B 115 (2011) 12537–12548. [26] K. Paduszyn´ski, U. Doman´ska, J. Phys. Chem. B 115 (2011) 8207–8215. [27] D.H. Everett, Trans. Faraday Soc. 61 (1965) 1637–1645. [28] A.J.B. Cruickshank, B.W. Gainey, C.P. Hicks, T.M. Letcher, R.W. Moody, C.L. Young, Trans. Faraday Soc. 65 (1969) 1014–1031. [29] DIPPR Project 801 – Full Version. Design Institute for Physical Property Research/AIChE, (2005; 2008–2010). Available from:
[38] A.-L. Revelli, F. Mutelet, J.-N. Jaubert, Ind. Eng. Chem. Res. 49 (2010) 3883– 3892. [39] L.M. Grubbs, S. Ye, M. Saifullah, M.C. McMillan-Wiggins, W.E. Acree Jr., M.H. Abraham, P. Twu, J.L. Anderson, Fluid Phase Equilib. 301 (2011) 257–266. [40] F. Mutelet, V. Ortega-Villa, J.-C. Moı¨se, J.-N. Jaubert, W. Acree Jr., J. Chem. Eng. Data 56 (2011) 3598–3606. [41] A. Marciniak, M. Wlazło, J. Chem. Thermodyn. 49 (2012) 137–145. [42] U. Doman´ska, M. Królikowska, J. Chem. Eng. Data 56 (2011) 124–129. [43] A. Marciniak, M. Wlazło, J. Chem. Thermodyn. 57 (2013) 197–202. [44] M. Montanino, M. Carewska, F. Alessandrini, S. Passerini, G.B. Appetecchi, Electrochim. Acta 57 (2011) 153–159. [45] K. Paduszyn´ski, U. Doman´ska, Ind. Eng. Chem. Res. 51 (2012) 591–604. [46] M. Krummen, D. Gruber, J. Gmehling, Ind. Eng. Chem. Res. 39 (2000) 2114– 2123. [47] C. Mölmann, J. Gmehling, J. Chem. Eng. Data 42 (1997) 35–40. [48] R. Kato, J. Gmehling, J. Chem. Thermodyn. 37 (2005) 603–619. [49] S. Nebig, V. Liebert, J. Gmehling, Fluid Phase Equilib. 227 (2009) 61–67. [50] W.E. Acree Jr., G.A. Baker, A.-L. Revelli, J.-C. Moı¨se, F. Mutelet, J. Chem. Eng. Data, J. Chem. Eng. Data 57 (2012) 3510–3518. [51] M. Krummen, P. Wasserscheid, J. Gmehling, J. Chem. Eng. Data 47 (2002) 1411–1417. [52] U. Doman´ska, A. Marciniak, Fluid Phase Equilib. 286 (2009) 154–161. [53] A. Marciniak, Int. J. Mol. Sci. 11 (2010) 1973–1990. [54] F. Mutelet, J.-N. Jaubert, J. Chem. Thermodyn. 39 (2007) 1144–1150.
JCT 12-659