Liquid-liquid equilibrium between water and ionic liquids

Transcription

1 Liquid-liquid equilibrium between water and ionic liquids M. G. Freire 1,*, L. M. N. B. F. Santos 2, I. M. Marrucho 1 and J. A. P. Coutinho 1 1 CICECO, Departamento de Química, Universidade de Aveiro, Aveiro, Portugal, 2 Centro de Investigação em Química, Departamento de Química, Faculdade de Ciências, Universidade do Porto, Porto *Corresponding author Tel: ; Fax: ; mmartins@dq.ua.pt

2 Abstract For ILs to be used effectively as solvents in liquid-liquid extraction, it is important to know the mutual solubilities between ILs and the second liquid phase. On this work we address the mutual solubilities between the water and ILs. This study focuses on hydrophobic ILs presenting low water solubilities. The influence of the anion and cation on the solubility was studied. The IL content in the water-rich phase was analyzed using UV-vis spectroscopy and the IL-rich phase was analyzed by Karl Fischer titration. The temperature range of the experimental analysis was between 288 and 318 K and at atmospheric pressure. The results seem to indicate that the solubility of water in ILs is less affected by the choice of the anion than the solubility of ILs in water. The hydrophobicity of the anions incresases in the order: [BF 4 ] < [PF 6 ] < [CF 3 (SO 2 ) 2 N] and the hydrophobicity of the cations increases, as expected, with the alkyl chain length. Keywords: solubility measurements, water, ionic liquids

3 1. Introduction Ionic liquids (ILs) are novel chemical compounds with a range of interesting characteristics that are driving a lot of research in various fields. Their stability, large liquidus range and good solvation properties for both polar and nonpolar compounds make them interesting as solvents for chemical reactions and separations. Their physical properties are tuneable by wise selection of cation, anion and substituents in both the cation or in the anion. The more common ILs are based on imidazolium and pyridinium salts. Their capabilities as solvents for organic synthesis are well established leading to enhanced yields and selectivities [1-3]. Since ionic liquids are nonvolatile, they cannot contribute to atmospheric pollution. This property makes them attractive replacements for organic solvents in several applications in the chemical industry, at a moment when pollution by volatile organic compounds (VOCs) is of great concern. Among the several applications foreseeable for ionic liquids in the chemical industry there has been considerable interest in the potential of ILs for separation processes where, among others, ILs have shown promising in the liquid-liquid extraction of organics from water. Huddleston et al. [4] showed that 1-butyl-3-methylimidazolium hexafluorophosphate, [bmim][pf 6 ], could be used to extract aromatic compounds from water. Fadeev and Meagher [5] have shown that two imidazolium ionic liquids with the hexafluorophosphate anion could be used for the extraction of butanol from aqueous fermentation broths. McFarlane et al. [6] in a recent work present partition coefficients for a number of organic compounds and discuss the possibilities and limitations of organic liquids for extraction from aqueous media. Other studies have shown that one can use ionic liquids for the extraction of metal ions from solution [7], aromatics from aromatic-alkane mixtures [8],

4 sulphur-containing aromatics from gasoline [9] and in the separation of isomeric organic compounds [10]. Although, one of the areas of technological application of ionic liquids is the liquid-liquid extraction of organics from water where most work addresses the solubility of organics, in particular alcohols, in ILs and very few studies aim at the mutual solubilities of water and ionic liquids [11]. Qualitative liquid-liquid-phase behaviour of some ILs with water has been reported. For instance, Dupont and co-workers [12] report that 1-butyl-3- methylimidazolium tetrafluoroborate, [bmim][bf 4 ], is miscible with water at room temperature and that [bmim][pf 6 ] is immiscible, based on whether a 50:50 (wt%) mixture forms one or two phases. More interestingly, they report that [bmim][bf 4 ] does phase split at lower temperatures, and they report the composition of both liquid phases at temperatures between (-8 and 5) ºC. They found that this system showed UCST (upper critical solution temperature) behaviour, which is typical of hydrocarbon-water mixtures. Seddon and co-workers [13] provide some general guidelines on IL-water miscibility. They indicate that halide, ethanoate, nitrate and trifluoroacetate salts are totally miscible with water, that [PF 6 ] and [(CF 3 SO 2 ) 2 N] salts are immiscible, and that [BF 4 ] and [CF 3 SO 3 ] salts can be totally miscible or immiscible depending on the substituents on the cation. Moreover, they report the solubility of water in the IL-rich phase for [C n mim][pf 6 ] and [C n mim][bf 4 ] as a function of chain length between 4 and 8 carbon number (CN) and 6 and 10 CN, respectively. However, very few quantitative data is available in the literature. Anthony et al. [11] present mutual solubilities of water in ionic liquids at ambient temperature for [bmim][pf 6 ], 1-octyl-3-methylimidazolium

5 hexafluorophospahte, [omim][pf6], and [omim][bf 4 ]. They used Karl Fischer titration to establish the water content in the ionic liquid and UV spectroscopy for the solubility of ILs in water. Wong et al. [14] studied the phase equilibrium of water and the ionic liquids 1-ethyl-3-methylimidazolium hexafluorophosphate, [emim][pf 6 ], and [bmim][pf 6 ]. Rebelo et al. [15] reported the entire liquid-liquid phase diagram of [bmim][bf 4 ] and water. Above the UV spectroscopy and Karl-Fischer titration for the determination of ILs in water, two other methods have been reported. Shvedene et al [16] measured the solubility in water of two disubstituted imidazoliums cations at room temperature, 1-butyl-3-methylimidazolium hexafluorophosphate and 1-butyl-3- methylimidazolium bis(triflyimide) and, two trisubstituted imidazoliums, 1-butyl-2,3- dimethylimidazolium and 1-butyl-2,3-methylimidazolium bis(triflyimide), using ionselective electrodes. Alfassi et al. [17] measured the solubility of three typical hydrophobic ionic liquids in water, [bmim][pf 6 ], methyltributylammonium bis(trifluoromethylsulphonyl)imide and butylmethylpyrrolidonium bis(trifluoromethylsulphonyl)imide at room temperature, using electrospray ionization mass spectrometry. These techniques have revealed to be less precise than UV spectroscopy, but it is a good alternative for ILs that not contain aromatic rings that could absorb in the UV range. No other significant work related to LLE of ILs and water seems to be reported. In this work, the solubility of water in [bmim][pf 6 ], [bmim][(cf 3 SO 2 ) 2 N], [omim][bf 4 ] and [omim][pf 6 ] in the temperature range between (288and 318) K at pressures close to atmospheric were determined, by means of Karl-Fischer titration.

6 On the other side, the solubility of the ILs in the water rich phase was already determined for the [bmim][pf 6 ] and [bmim][(cf 3 SO 2 ) 2 N], in the same temperature and pressure conditions. Although, the solubility of the omim cations is being performed and the experimental results will be presented, evaluated and discussed during the meeting. With the [bmim][pf 6 ] and [bmim][(cf 3 SO 2 ) 2 N] and, [omim][bf 4 ] and [omim][pf 6 ] pair groups some conclusions about the anion effect on he mutual solubilites can be made. With the pair [bmim][pf 6 ] and [bmim][pf 6 ] some conclusions about the cation effect and the increase of the chain length can be drawn. 2. Experimental 2.1 Materials Liquid-liquid mutual solubilities were measured for several ionic liquids from the family of the imidazolium cation. The 1-butyl-3-methylimidazolium hexafluorophosphate, [bmim][pf 6 ], and the 1-methyl-3-octylimidazolium hexafluorophosphate, [omim][pf 6 ], were acquired at Solvent Innovation. The 1-butyl-3- methylimidazolium bis(trifluoromethylsulfonyl)imide, [bmim][cf 3 SO 2 ) 2 N], was acquired at Merck and, the 1-methyl-3-octylimidazolium tetrafluoroborate, [omim][bf 4 ], was acquired at Solchemar. The ILs were used without further purification since their purity was =99 mass%, with halides content smaller than 100 ppm. The ILs were stored under vacuum with silica. Water was distilled, passed by a reverse osmosis system and further treated with a Milli-Q plus 185 purification apparatus. This water had a resistivity of 18.2 µω.cm, a TOC smaller than 5 µg.l -1 and it is free of particles greater than 0.22 µm.

7 2.2 Solubility measurements The two phases were initially vigorously agitated and let to reach equilibrium by the separation of both phases in 22 ml vials for at least 24 h. These vials were inside an aluminium block, immersed in an air bath, composed by a heater and a convention fan. In order to obtain temperatures below the room temperature a Julabo circulator, model F25-HD was employed. Temperature was controlled with a temperature control, PID associated with a Pt 100. The solubility measurements for both rich phases were performed in the temperature range ( ) K and at atmospheric pressure. Solubility of water in the IL rich phase was determined with a Metrohm 831 Karl-Fischer (KF) coulometer. Water molar fractions were determined gravimetrically weighing a syringe, with samples of 0.2 ml, before and after the KF injection. The experimental method was validated comparing our experimental data with literature data for the IL the [bmim][pf 6 ], since there are some literature reporting the equilibrium between this organic compound and water. Solubility of ILs in the water rich phase was determined by UV-spectroscopy with a SHIMADZU UV-1700 Pharma-Spec Spectrometer, at 211 nm. This wavelength was choosen because it was found to be the maximum absorption length for all ILs studied, corresponding to the UV absorption of the imidazolium ring cation. Samples of ( ) ml were taken of the water rich phase by means of a syringe and diluted by a factor of 1:1000 (v:v) in pure water. The molar fractions were determined gravimetrically, due to the higher precision obtained in this kind of dilutions. The dilutions of the water rich phase samples avoid the loss of the IL compound to the vapor phase when working at temperatures different from room temperature.

8 For each temperature, 5 different extractions were made in both equilibrium phases, and standard deviations were determined. Several measurements in different times were performed until no results differences were accomplished in order to guarantee that the two phases are completely separated and equilibrium was accomplished. 3. Results and discussion 3.1 Experimental results Experimental mole fractions values obtained for the solubility of water in liquid room temperature ILs between ( and ) K are presented in Table 1. A comparison between the measured and literature data for [bmim][pf 6 ] is presented in Figure 1. The precision of this method is well described by the good agreement of the values obtained and the reported literature, since no other significant work related with this side of the liquid-liquid equilibrium phase diagram of [bmim][pf 6 ] and water has been reported, being this, the imidazolium based ionic liquid most studied in the literature. There is very little data available in the open literature for these systems, and most of these data is widely scattered, as can be seen. The comparison between our results of the solubility of water in the different ILs studied is presented in Figure 2. From the experimental data, it can be seen that the [omim][bf 4 ] is the more hygroscopic IL studied and the [omim][pf 6 ] is the less hygroscopic one. This is a direct effect of the anion that defines the hygroscopicity of each salt.

9 For the three more hydrophobic ILs, the solubility of ILs in water increases in the order: [omim][pf 6 ] < [bmim][cf 3 (SO 2 ) 2 N] < [bmim][pf 6 ]. With these results it can be concluded that the [bmim] cation presents a higher affinity for water when compared with the [omim] with a higher chain length and, because of that, a more hydrophilic cation. On the other side the [PF 6 ] anion is more hydrophilic than the [CF 3 (SO 2 ) 2 N]. Experimental mole fractions values obtained for the solubility of ILs in water at several temperatures and at atmospheric pressure are presented in Table 2. A comparison between literature data and our experimental values is presented in Figure 3. Since there is no much available data in the open literature for the solubility of [bmim][pf 6 ] in the water rich phase, and most of these data are widely scattered, it can be concluded that our experimental results agree well with the reported literature, with much more precision and smaller standard deviations, confirming the viability of the technique adopted. Analyzing the experimental data obtained until the moment, its visible that the [bmim][pf 6 ] is much more soluble in water (one order of magnitude in mole fraction higher) than the [bmim][cf 3 (SO 2 ) 2 N], confirming the results obtained in the other phase of the equilibrium, where the first one is more hydrophilic with a higher affinity for water. From this preliminary results it seems to be a generalized trend, the solubility of water in ILs is less affected by the anion than the solubility of ILs in water. A more precise conclusion will be presented during the meeting. Thermodynamic functions such as Gibbs energy, enthalpy and entropy of solution and solvation in both equilibrium phases can be determined from the solubility dependence. The thermodynamic functions of solution and solvation for water in the ILs

10 solvents, which are difficult to obtain through calorimetric measurements, are presented in Table 3. Comparing all the enthalpies of solvation, it can be concluded that the [BF 4 ] anion effect derives from the higher enthalpy of solvation that the [omim][bf 4 ] compound presents. Also, the difference in the hydrophobicity of the compounds [omim][bf 4 ] and [omim][pf 6 ] arises from the differences that can be observed in the enthalpies of solvation between the two ILs. On the other side, the close enthalpy of solvation of [bmim][pf 6 ] and [bmim][cf 3 (SO 2 ) 2 N] confirm the similar hydrophobic properties between the two compounds. When analyzing the entropy of solution and solvation of all the compounds it can be detected a proximity between values for compounds with the same cation. This effect results form the orientation of each cation in the solvent phase, since they present different chain length. The thermodynamic functions for the solubility of ILs in water will be presented and discussed during the meeting, after finishing the experimental work for all the compounds above interest in the temperature range proposed. Conclusions Original data for the mutual solubilities between water and ILs in the temperature range between ( and ) K and at pressures close to atmospheric were presented.

11 The mutual solubilities are primarily defined by the anion present in the salt. However the solubility of water in ILs is less affected by the presence of the anion than the solubility of ILs in water. By the experimental data obtained some conclusions about the hydrophobichydrophilic character of the salts can be drawn. The hydrophobicity of the anions incresases in the order: [BF 4 ] < [PF 6 ] < [CF 3 (SO 2 ) 2 N] and the hydrophobicity of the cations increases, as expected, with the alkyl chain length. Acknowledgements This project was financed by Fundação para a Ciência e Tecnologia, POCTI/EQU/58152/2004. M. G. Freire acknowledge the financial support from Fundação para a Ciência e a Tecnologia through their PhD. (SFRH/BD/14134/2003) scholarship.

12 References [1] T. Welton, Chem. Rev., 99 (1999) [2] P. Wasserscheid, T. Welton, Ionic Liquids in Synthesis, Wiley-VCH, Weinheim (2003) 364. [3] P. Wasserscheid, W. Keim, Angew. Chem. Int. Ed. 39 (2000) [4] J.G Huddleston, H.D. Willauer, R.P. Swatloski, A.E. Visser, R.D. Rogers, Chem. Commun. (1998) [5] A.G. Fadeev, M.M. Meagher, Chem. Commun. (2001) 295. [6] J. McFarlane, W.B. Ridenour, H. Luo, R.D. Hunt, D.W. DePaoli, R.X. Ren, Separation Science and Technology 40 (2005) [7] M.S. Selvan, M.D. McKinley, R.H. Dubois, J.L.J. Atwood, J. L., J. Chem. Eng. Data 45 (2000) 841. [8] T.M. Letcher, N. Deenadayalu, B. Soko, D. Ramjugernath, P.K. Naicker, J. Chem. Eng. Data 48 (2003) 904. [9] S. Zhang, Z.C. Zhang, Green Chem. 4 (2002) 376. [10] L.C. Branco, J.G. Crespo, C.A.M. Afonso, Angew. Chem. Int. Ed. 41 (2002) [11] J.L. Anthony, E.J. Maginn, J.F. Brennecke, J. Phys. Chem. B 105 (2001) [12] J.E.L. Dullius, P.A.Z. Suarez, S. Einloft, R.F. de Souza, J. Dupont, J. Fischer, A. De Cian, Organometallics 17 (1998) 815. [13] K.R. Seddon, A. Stark, M.J. Torres, Pure Appl. Chem. 72 (2000) [14] D.S.H. Wong, J.P. Chen, J.M. Chang, C.H. Chou, Fluid Phase Equilib. 194 (2002) 1089.

13 [15] L.P.N. Rebelo, V. Najdanovic-Visak, Z.P. Visak, M.N. da Ponte, J. Szydlowski, C.A. Cerdeirina, J. Troncoso, L. Romani, J.M.S.S. Esperança, H.J.R. Guedes, H.C. de Sousa, Green Chemistry 6 (2004) 369. [16] N.V. Shvedene, S.V. Borovskaya, V.V. Sviridov, E.R. Ismailova, I.V. Pletenev, Anal. Bioanl. Chem. 381 (2005) 427. [17] Z. B. Alfassi, R.E. Huie, B.L. Milman, P. Neta, Anal. Bioanal. Chem. 377 (2003) 159.

14 Table 1: Experimental mole fraction solubilities (x) of water in the ILs. [bmim][pf 6 ] [bmim][(cf 3 SO 2 ) 2 N] [omim][bf 4 ] [omim][pf 6 ] T / (K) 10 (x ± σ a ) 10 (x ± σ a ) 10 (x ± σ a ) 10 (x ± σ a ) ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± a standard deviation

15 Table 2: Experimental mole fraction solubilities (x) of ILs in water. [bmim][pf 6 ] [bmim][(cf 3 SO 2 ) 2 N] T / (K) 10 3 (x ± σ a ) 10 4 (x ± σ a ) ± ± ± ± ± ± ± ± ± ± a standard deviation

16 Table 3: Thermodynamic functions of solution and solvation of water in the several ILs at K Thermodynamic Function / IL [bmim][pf 6 ] [bmim][(cf 3 SO 2 ) 2 N] [omim][bf 4 ] [omim][pf 6 ] ( sol H ± σ a ) / (KJ mol -1 ) 10.6 ± ± ± ± 0.75 ( solv H ± σ a ) / (KJ mol -1 ) ± ± ± ± 0.8 ( sol G ± σ a ) / (KJ mol -1 ) 3.3 ± ± ± ± ( solv G ± σ a ) / (KJ mol -1 ) -5.3 ± ± ± ± ( sol S ± σ a ) / (J mol -1 K -1 ) ± ± ± ± 0.03 ( solv S ± σ a ) / (J mol -1 mol -1 ) ± ± ± ± 0.03 a standard deviation

17 Figure Captions Figure 1: Mole fraction solubility (x) of water in [bmim][pf 6 ] at several temperatures:, this work. Other literature data:, [11];, [13]; ο, [14]; +, [16]. Figure 2: Mole fraction solubility (x) of water at several temperatures in:, [omim][bf 4 ];, [bmim][pf 6 ]; ο, [bmim][cf 3 (SO 2 ) 2 N];, [omim][pf 6 ]. Figure 3: Weight mass fraction solubility (wt%) of [bmim][pf 6 ] in water at several temperatures:, this work. Other literature data:, [11]; ο, [14]; +, [16];, [17].

18 Figure 1-0,90-1,00-1,10-1,20 ln(x) -1,30-1,40-1,50-1,60-1,70 0, , , , , , /Temperature/ (K -1 )

19 Figure 2 0,00-0,20-0,40-0,60 ln(x ) -0,80-1,00-1,20-1,40-1,60-1,80 0, , , , , /Temperature/ (K -1 )

20 Figure 3 4,20 3,70 3,20 wt(%) 2,70 2,20 1,70 1,20 285,00 290,00 295,00 300,00 305,00 310,00 315,00 320,00 325,00 Temperature / K