Thermodynamic study of binary mixtures containing 1-butylpyridinium tetrafluoroborate and methanol, or ethanol

J. Chem. Thermodynamics 42 (2010) 1500–1505

Contents lists available at ScienceDirect

J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct

Thermodynamic study of binary mixtures containing 1-butylpyridinium tetrafluoroborate and methanol, or ethanol M. García-Mardones, V. Pérez-Gregorio, H. Guerrero, I. Bandrés, C. Lafuente * Departamento de Química Orgánica-Química Física, Facultad de Ciencias, Universidad de Zaragoza, 50009 Zaragoza, Spain

a r t i c l e

i n f o

Article history: Received 9 July 2010 Accepted 23 July 2010 Available online 1 August 2010 Keywords: Alkanol Ionic liquid Thermophysical properties (Vapour + liquid) equilibrium

a b s t r a c t Densities and speeds of sound have been determined for the binary mixture (1-butylpyridinium tetrafluoroborate + methanol, or ethanol) over the temperature range 293.15 K to 323.15 K. From experimental values, excess volume and excess isentropic compressibility have been calculated. The mixtures give negative values for the excess properties. Besides, (vapour + liquid) equilibrium in isothermal conditions has been obtained for these systems at T = 303.15 K and T = 323.15 K, which has allowed us to derive activity coefficients and excess Gibbs functions. Positive deviations from Raoult’s law have been found. A detailed analysis and interpretation of results have been carried out in structural and energetic terms using thermodynamic information of the pure compounds. Ó 2010 Elsevier Ltd. All rights reserved.

1. Introduction Ionic liquids (ILs) have been consolidating in the scientific panorama as a real alternative to molecular media. Consequently, their use has been widespread in many different applications due to their singular features, which presage a promising future for this type of compounds [1–5]. However, the lack of knowledge about the properties that characterize the ILs in particular conditions is, at present, a barrier that difficult their implementation in more areas. In this way, the number of studies is increasing that analyze these questions in terms of the ionic structures and intermolecular interactions, not only for the pure ILs, but also, for their mixtures, with the aim of correlating the structural and energetic aspects which influence their behaviour [6–15]. These are the reasons that have motivated us to pursue a wellplanned study of a series of ILs based on the pyridine ring in view of the need to describe their thermodynamic and transport properties and the relationship with their characteristics [16–20]. Now, in order to complete this useful information and to encourage their versatility, we are going to extend the investigation to binary mixtures constitute by ILs. Specifically in this paper, we report a systematic thermodynamic study of binary systems formed by an ionic liquid, 1-butylpyridinium tetrafluoroborate, [bpy][BF4] and an alcohol, methanol, or ethanol, as a function of temperature. Densities and speeds of sound have been determined at temperatures 293.15 K, 303.15 K, 313.15 K, and 323.15 K for these mixtures, except for the

* Corresponding author. Tel.: +34 976 762295; fax: +34 976 761202. E-mail address: [email protected] (C. Lafuente). 0021-9614/$ - see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jct.2010.07.014

([bpy][BF4] + ethanol) system in which miscibility problems have prevented the measurements at the lowest temperature. From these values, excess volume and excess isentropic compressibility have been calculated and correlated to get useful information about the phenomena that take place in the mixture. Furthermore, the isothermal vapour liquid equilibrium has been obtained at T = 303.15 K and T = 323.15 K. Activity coefficients and excess Gibbs function have been derived and interpreted in energetic terms.

2. Experimental To determine the thermodynamic properties of the mixtures studied, the alcohols methanol (mass fraction purity 0.998) and ethanol (mass fraction purity 0.995) were obtained from Aldrich. The purity of the chemicals was checked by comparing the measured properties at the temperature of 303.15 K with those reported in the literature. No further purification was considered necessary. Moreover, the ionic liquid 1-butylpyridinium tetrafluoroborate (0.99), was provided by IoLiTec. With the aim of decreasing the water content as much as possible, the IL was dried for 24 h under a vacuum of ca. 0.05 kPa while stirring and stored before use in a desiccator. The water content of the sample was less than 100  106 as determined by Karl–Fischer titration using an automatic titrator Crison KF 1S-2B. The properties of the pure components at the working temperatures and the comparison with literature values at T = 303.15 K [14,21–24] are gathered in table 1. Densities, q, and speed of sounds, u, of the samples were determined simultaneously with an Anton Paar DSA 5000 vibrating tube densimeter and sound analyser, automatically thermostatted within ±0.001 K. By measuring the damping of the oscillation of

1501

M. García-Mardones et al. / J. Chem. Thermodynamics 42 (2010) 1500–1505

TABLE 1 Thermophysical properties of pure components at different working temperatures and comparison of density, speed of sound, and vapour pressure with literature values at T = 303.15 K.

q/(g  cm3)

T/K

Exp.

u/(m  s1) Lit.

Exp.

p/kPa Lit.

Exp.

Cp,m/(J  mol1  K1)

a/kK1

392.0 398.0 403.0 408.0

0.5626 0.5657 0.5688 0.5722

B/(cm3  mol1)

Lit.

Butylpyridinium tetrafluoroborate 293.15 303.15 313.15 323.15

1.217107 1.210289 1.203723 1.196678

293.15 303.15 313.15 323.15

0.791243 0.781813 0.772287 0.762628

1.2100 [14]

1611.47 1588.11 1565.57 1543.45 Methanol

303.15 313.15 323.15

0.781302 0.772569 0.763647

0.781808 [21]

0.78115 [23]

1118.83 1085.99 1053.55 1021.53 1127.20 1093.64 1060.44

1086.37 [21]

21.930

1128 [23]

xi M i

i

1

q



1

qi

114.28 118.55 123.21

1.1081 1.1206 1.1337

1

jS ¼

ð1Þ

;

1167 2730 1662

ð2Þ

;

qu2 jES ¼ jS  jidS ; where

jidS ¼

ð3Þ

jidS was obtained according to Benson and Kiyohara [25]:

X

 /i

jS;i þ

i

! P 2  X / i ai TV i a2i ; T xi V i P i C p;mi i xi C p;mi i

ð4Þ

where /i is the volume fraction of component i in the mixture referred to the unmixed state, xi is the corresponding mole fraction, T is the absolute temperature and jS,i, Vi, ai, and Cp,i are the isentropic compressibility, the molar volume, the thermal expansion, and the molar heat capacity of the pure component i, respectively. Thermal expansion coefficients were derived from densities determined in the laboratory, whereas molar heat capacities have been taken from the literature [16,26,27]. Values for both properties have been collected in table 1. Excess molar volume and excess isentropic compressibility have been correlated with the following equation [28]:

P



1828

where Mi is the molar mass of the components and xi is the mole fraction. Isentropic compressibility, jS, and excess isentropic compressibility, jES , were estimated from density, q, and speed of sound, u, using the relations:

Y ¼ x1 x2

Excess molar volumes, VE, were calculated from experimental densities of the mixture, q, and pure compounds, qi, using the following equation:

X

10.475 [24]

29.565

3. Results and discussion

VE ¼

1.1968 1.2113 1.2262 1.2417

55.785 Ethanol 10.475

the U-tube caused by the viscosity of the filled-in sample, the DSA 5000 automatically corrects viscosity related errors in the density. The calibration was carried out with ultra pure water supplied by SH Calibration service GmbH, and dry air. The final uncertainty of density and speed of sound are estimated at ±10 g  cm3 and ±0.01 m  s1, respectively. The experimental vapour liquid equilibrium was obtained using an all-glass dynamic re-circulating type still that was equipped with a Cottrell pump. It is a commercial unit (Labodest model) built by Fischer. The equilibrium temperature were measured to an accuracy of ±0.01 K by means of a thermometer (model F25 with a PT100 probe) from Automatic Systems Laboratories, and the pressure in the still was measured with a Digiquartz 735215A-102 pressure transducer from Paroscientific equipped with a Digiquartz 735 display unit. The accuracy of the pressure measurements is ±0.01% of reading. Composition of the liquid-phase was determined by density measurements. The error in the determination of mole fractions is estimated to be ±0.0002.



80.19 82.14 84.31 86.73

21.904 [22]

1þ

i¼0 Ai ðx1

P

 x2 Þ i

j¼1 Bj ðx1

ð5Þ

 x2 Þ j

where Y is VE or jES , Ai and Bj are adjustable parameters and x1 and x2 are the mole fractions. Values of the fitting parameter together with standard deviations, r(Q), are gathered in table 2. Excess volume and excess isentropic compressibility are plotted in figures 1 to 4.

TABLE 2 Parameters, Ai and Bi, and standard deviations, r(Q), for equation (5). Function

T/K

A0

VE/(cm3  mol1)

293.15 303.15 313.15 323.15

2.7221 2.8810 3.0464 3.5100

jES =TPa1

293.15 303.15 313.15 323.15

287.89 325.60 369.27 421.53

VE/(cm3  mol1)

303.15 313.15 323.15

2.4495 2.6387 3.0051

jES =TPa1

303.15 313.15 323.15

313.25 360.54 407.77

A1

A2

Butylpyridinium tetrafluoroborate + methanol 0.0175 1.2090 0.0711 1.5424 0.4227 1.2930 0.4860 1.4038 158.00 172.84 207.36 273.00

31.79 60.48 72.75 66.11

Butylpyridinium tetrafluoroborate + ethanol 1.4588 1.3436 1.6506 1.6360 2.1176 2.2535 325.44 403.19 460.41

230.89 344.98 463.63

A3

B1

r(Q)

0.2306 0.3758 0.7994 0.5752

0.9288 0.9560 0.9322 0.9117

0.0071 0.0072 0.0061 0.0054

39.44 71.84 25.00 120.01

0.88 0.90 0.87 0.89

0.48 1.31 1.41 1.78

1.2946 1.4018 1.4271

0.0060 0.0061 0.0071

120.04 199.86 314.21

0.92 1.07 1.04

1502

M. García-Mardones et al. / J. Chem. Thermodynamics 42 (2010) 1500–1505

0

0

-0.1

-20

-0.2

-40

T = 293.15 K -0.3

-60 -0.4

-1

κS / TPa

E

-0.6

E

3

V / cm ·mol

-1

-80

T = 293.15 K

-0.5

-100

-120

-0.7

-140 -0.8

-160

-0.9

-180

-1.0

-200

-1.1 -1.2

-220 0

0.2

0.4

0.6

0.8

0

1

0.2

0.4

x1

0.6

0.8

1

x1

FIGURE 1. Excess volumes, VE, for butylpyridinium tetrafluoroborate + methanol as a function of mole fraction, x1: (d) experimental data; (—) equation (5).

FIGURE 3. Excess isentropic compressibilities, jES , for butylpyridinium tetrafluoroborate + methanol as a function of mole fraction, x1: (d) experimental data; (—) equation (5).

0 Besides, the vapour liquid equilibria in isothermal conditions have been studied for the considered mixtures. The vapour pressure of the ionic liquid is so low at these working conditions that it can be considered zero attending to the sensitivity of the measurements [29]. Thus, it can be assumed that the total vapour pressure is equal to the partial pressure of the alcohol. Pressure-liquid composition diagrams p–x1 are graphically represented in figures 5 and 6. The Wilson equation has been used to correlate the activity coefficients of the components in the liquid-phase. Estimation of the adjustable parameters of the equation was based on minimization of the following objective function in terms of experimental and calculated pressure values:

-0.1

-0.2

-0.3

T = 303.15 K

F¼

X pexp  pcal 2 : pexp i i

ð6Þ

The calculated pressure, pcal, is obtained taking into account both the non-ideality of the vapour phase and the variation of the Gibbs function with pressure as follows:

-0.5

E

3

V / cm ·mol

-1

-0.4

-0.6

0 p

-0.7

-0.8

-0.9

-1 0

0.2

0.4

0.6

0.8

1

x1 FIGURE 2. Excess volumes, VE, for butylpyridinium tetrafluoroborate + ethanol as a function of mole fraction, x1: (d) experimental data; (—) equation (5).

cal

¼

x2 2 p02

c

exp @

V 02  B22



RT

p  p02

1 A;

ð7Þ

where R is the gas constant, x2 is the mole fractions of alcohol in the liquid-phase, p is the total pressure, and p02 , B22, and V 02 are the vapour pressure, the second virial coefficient and the molar volume of pure alcohol. Values of experimental vapour pressures of pure alcohols at T = 303.15 K and T = 323.15 K together with the corresponding second virial coefficients taken from TRC-tables [30], are collected in table 1. Parameters for the coefficient activity correlation along with average deviations in pressure, Dp, are gathered in table 3. The values of the excess Gibbs function for the studied systems are plotted in figures 7 and 8. Values of experimental density and speed of sound together with calculated values of excess molar volume and excess isentropic compressibility for the studied mixtures at different temperatures are found in the Supplementary material. Isothermal (vapour + liquid) equilibria data, that is, vapour pressure, liquid-phase composition, activity coefficients and the corresponding excess Gibbs functions calculated using the Wilson equation, also are given as Supplementary material.

1503

M. García-Mardones et al. / J. Chem. Thermodynamics 42 (2010) 1500–1505

60

0

55

-20 50

T = 323.15 K

45

-40

T = 303.15 K 40

-60

p / kPa

-80

E

κS / TPa

-1

35 30 25

T = 303.15 K

-100 20 15

-120

10

-140 5 0

-160 0

0.2

0.4

0.6

0.8

0

1

0.2

0.4

FIGURE 4. Excess isentropic compressibilities, jES , for butylpyridinium tetrafluoroborate + ethanol as a function of mole fraction, x1: (d) experimental data; (—) equation (5).

0.6

0.8

1

x1

x1

FIGURE 5. p–x1 diagrams for butylpyridinium tetrafluoroborate + methanol: (d) experimental data; (—) Wilson’s equation.

32

28

24

T = 323.15 K

20

p / kPa

Excess volume and excess isentropic compressibility are negative over the whole composition range at all temperatures for both systems. Therefore, the mixtures studied here show the usual behaviour since VE and jES have the same sign. Values of the excess properties are larger in absolute value for the mixture containing methanol than for the system with ethanol. Moreover, it is noticeable that the curves are remarkably asymmetric, with their minima shifted toward the rich compositions in alcohol. The effect of temperature is similar for both mixtures; values of the excess properties decrease when temperature rises. Besides, a slight displacement of the minimum values toward the weak IL region when temperature increases takes place for both excess volume of (IL + methanol) and the excess isentropic compressibility of the mixture containing ethanol. Activity coefficients in both mixtures are greater than unity at T = 303.15 K and T = 323.15 K. These values are especially high for the ionic liquid at low concentration. This means that these mixtures show positive deviations from the ideal behaviour. As shown in figures 7 and 8, the excess Gibbs functions calculated from activity coefficients at the temperatures studied also show positive values for both systems that decrease slightly when temperature rises. Larger values of GE are found for the mixture containing ethanol than those when the methanol forms part of the mixture. Before facing a detailed analysis of thermodynamic properties for each of the considered systems, it is important to highlight some characteristics of the pure compounds to facilitate an appropriate interpretation of the phenomena that can occur between the components of the mixture. As we have mentioned before, the ionic liquid [bpy][BF4] is based on the pyridine ring, with a butyl chain attached to the nitrogen atom. The tetrafluoroborate anion shows a tetrahedral disposition in which the presence of the highly electronegative fluorine atoms contribute to the distribution of the anionic charge of borate. In this way, this anion exhibits a low surface electrical charge density compared to other common anions [6,16]. Regarding the intermolecular interactions in this type of IL, electrostatic and van der Waals forces are the most important, playing a key role in the behaviour of these compounds [31,32]. Moreover, they are characterized by other specific interactions such as p–p interactions and hydrogen bonds. Thus, the pyridinium cations show a high cohesive ability associated with interactions between the electronic clouds [33]. On the other side, the aromatic protons in the pyridinium cation show a lower acidic character compared to imidazolium-based IL due to the greater size of the ring and the presence of only one electronegative atom, so the contribution of these bonds seems to be less important [34].

16

T = 303.15 K

12

8

4

0 0

0.2

0.4

0.6

0.8

1

x1 FIGURE 6. p–x1 diagrams for butylpyridinium tetrafluoroborate + ethanol: (d) experimental data; (—) Wilson’s equation.

1504

M. García-Mardones et al. / J. Chem. Thermodynamics 42 (2010) 1500–1505

The IL has been mixed with two alcohols, methanol and ethanol. It is worth noting that the size of both compounds is considerably smaller compared with the ions that constitute the IL. Moreover, their behaviour is strongly influenced by the hydrogen bonds which are established between them. It is suggested that these alcohols induce relatively strong dipolar interactions since the dipole moments in the liquid state at T = 293.15 K are 2.87 D and 1.66 D for methanol and ethanol, respectively [35]. Due to the features of the compounds which form part of the mixtures studied, it is possible to infer the structural and energetic effects that are involved in the mixture to explain their deviations from the ideal behaviour. Structural effects are connected to differences in size and shape of the unlike compounds, which lead to changes in the packing of the components when the mixture is established. It is expected that a great influence of these effects is due to the disparity in size between the IL and the alcohols. Besides, energetic effects are related with the weakness or breaking of the interactions in the pure compounds and the formation of new interactions between the chemicals once the mixture is formed. Thus, it is necessary to take into account, among others, the rupture of hydrogen bonds between the alcohol molecules and the reduction of electrostatic and van der Waals forces in the IL. Moreover, the formation of interactions between the hydroxyl group in the alcohol and the tetrafluoroborate anion as well as the new dipole–dipole interactions between both components, are also important. Some studies revealed that interactions between the alcohol and the pyridinium cation are less important than those with the anions [11,36]. Consequently, the resultant behaviour reflected in the excess properties is attributed to the multiple factors that take place in the mixture.

TABLE 3 Correlation parameters for the Wilson equation and average deviation in pressure, Dp. k12  k11/(J  mol1)

k21  k22/(J  mol1)

System

T/K

Methanol

Butylpyridinium tetrafluoroborate+ 303.15 2708.08 2585.94 323.15 2247.15 2602.16

0.046 0.133

303.15 323.15

0.054 0.093

Ethanol

10008.03 7998.47

3320.85 2960.19

Dp/kPa

Excess Gibbs functions calculated from activity coefficients show positive deviations from the ideal behaviour. Since this property is mainly related to energetic effects, it seems to be associated with the disruption of hydrogen bonds between the alcohol molecules and the weakness of electrostatic and van der Waals interactions in the IL. Experimental studies for similar systems have found exothermic mixtures in those cases in which the anion of the IL is sufficiently basic for promoting anion-alcohol hydrogen bonds [11]. Since the tetrafluoroborate anion is a weak Lewis base, without markedly donor solvent properties, the GE results for the ([bpy][BF4] + alcohol) mixtures are justified. Moreover, a decrease in the property with temperature occurs, which could be explained by the weakness of hydrogen bonds in the alcohols with increasing of temperature. Taking into account the absolute values of this property, it seems that the breaking of forces is more important in the mixtures containing ethanol. On the other hand, excess volumes and excess isentropic compressibility depend on both, structural and energetic effects. Results for these properties reflect a considerable decreasing of the unoccupied part of the volume in the mixtures compared to an ideal solution. According to the GE analysis, in which it is clear that energetic factors contribute to opposite effects, it appears that structural aspects are responsible for the negative values in the excess properties. The considerable contraction of the final volume is due to a better arrangement of the components in the mixture that can compensate the rupture of interactions. The large differences between the volumes of the components cause this behaviour and could justify the observed asymmetry of the curves [10]. Deviations are more pronounced for the smallest alcohol, probably owing to a greater difference in size than those with ethanol. Furthermore, as noted, energetic effects are most important in the mixture constituted by ethanol. The analysis of results can be compared with those previously obtained for similar systems. For example, excess properties have been measured such as excess volume or enthalpy, and deviations in viscosity for binary mixtures of (1-butyl-nmethylpyridinium tetrafluoroborate with n = 2, 3, or 4 + methanol, or ethanol) [10,11,36–38]. In this case, it is possible to study the effect of introducing a methyl group in the pyridinium cation on mixture properties. Although it is not possible to make a direct comparison, since properties are determined at different temperatures, it appears that excess volumes for our systems share the sign and are larger in absolute value than those for the mixtures constituted by the isomeric ILs. Moreover, excess enthalpies seem to reflect similar interpretations in energetic terms, since endothermic effects for these systems have been deduced. On the other side, the cationic structure dependence can be evaluated if results for mixtures formed by the analogous IL in the imidazolium family, 1-butyl-3-methylimidazolium tetra-

1200

550

500 1000

450

T = 303.15 K

T = 303.15 K 400 800

-1

300

G / J·mol

250

600

E

E

G / J·mol

-1

350

T = 323.15 K 200

T = 323.15 K 400

150

100

200

50

0

0 0

0.2

0.4

0.6

0.8

1

x1 FIGURE 7. Excess Gibbs functions, GE, for butylpyridinium tetrafluoroborate + methanol as a function of mole fraction, x1.

0

0.2

0.4

0.6

0.8

1

x1 FIGURE 8. Excess Gibbs functions, GE, for butylpyridinium tetrafluoroborate + ethanol as a function of mole fraction, x1.

M. García-Mardones et al. / J. Chem. Thermodynamics 42 (2010) 1500–1505 fluoroborate, are considered [9,39]. In this case, it is curious that the final values for excess properties such as excess volume and the excess Gibbs function are very similar to those obtained here, showing negative values for the first one and positive deviations from the Raoult law for the second one.

4. Conclusions A thermodynamic study of binary mixtures (IL + alcohol) is presented. Excess volumes, excess isentropic compressibility and excess Gibbs function at several temperatures have been exhaustively described and interpreted according to the features of pure compounds. From the experimental results, it is deduced that structural effects are very important to explain the volumetric properties for these systems owing to the differences in size between both components. Besides, the disruption and the weakness of interactions of pure chemicals once the mixture is established are the most significant aspects among energetic terms. Acknowledgements The authors thank the financial assistance support from Spanish Ministerio de Ciencia e Innovación (CTQ2009-09458) and Diputación General de Aragón. Victor Pérez-Gregorio and Isabel Bandrés thank her predoctoral fellowship from DGA. Appendix A. Supplementary data Supplementary data associated with this article can be found, in the online version, at doi:10.1016/j.jct.2010.07.014. References [1] P. Wasserscheid, T. Welton, Ionic Liquids in Synthesis, Wiley-VCH Verlag, Weinheim, 2003. [2] M. Armand, F. Endres, D.R. MacFarlane, H. Ohno, B. Scrosati, Nat. Mater. 8 (2009) 621–629. [3] T. Welton, Chem. Rev. 99 (1999) 2071–2083. [4] T.L. Greaves, C.J. Drummond, Chem. Rev. 108 (2008) 206–237. [5] P. Sun, D.W. Armstrong, Anal. Chim. Acta 661 (2010) 1–16. [6] H. Tokuda, K. Hayamizu, K. Ishii, M. Abu Bin Hasan Susan, M. Watanabe, J. Phys. Chem. B 108 (2004) 16593–16600. [7] Y. Yoshida, O. Baba, C. Larriba, G. Saito, J. Phys. Chem. B 111 (2007) 12204–12210. [8] Q.H. Zhang, Z.P. Li, J. Zhang, S.G. Zhang, L.Y. Zhu, J. Yang, X.P. Zhang, Y.Q. Deng, J. Phys. Chem. B 111 (2007) 2864–2872. [9] A.L. Revelli, F. Mutelet, J.N. Jaubert, J. Chem. Thermodyn. 42 (2010) 177–181.

1505

[10] J. Ortega, R. Vreekamp, E. Marrero, E. Penco, J. Chem. Eng. Data 52 (2007) 2269–2276. [11] A. Navas, J. Ortega, R. Vreekamp, E. Marrero, J. Palomar, Ind. Eng. Chem. Res. 48 (2009) 2678–2690. [12] M. Anouti, A. Vigeant, J. Jacquemin, C. Brigouleix, D. Lemordant, J. Chem. Thermodyn. 42 (2010) 834–845. [13] U. Domanska, M. Krolikowski, K. Paduszynski, J. Chem. Thermodyn. 41 (2009) 932–938. [14] B. Mokhtarani, A. Sharifi, H.R. Mortaheb, M. Mirzaei, M. Mafi, F. Sadeghian, J. Chem. Thermodyn. 41 (2009) 323–329. [15] T. Singh, A. Kumar, M. Kaur, G. Kaur, H. Kumar, J. Chem. Thermodyn. 41 (2009) 717–723. [16] I. Bandres, F.M. Royo, I. Gascon, M. Castro, C. Lafuente, J. Phys. Chem. B 114 (2010) 3601–3607. [17] I. Bandres, D.F. Montano, I. Gascon, P. Cea, C. Lafuente, Electrochim. Acta 55 (2010) 2252–2257. [18] I. Bandres, G. Pera, S. Martín, M. Castro, C. Lafuente, J. Phys. Chem. B 113 (2009) 11936–11942. [19] I. Bandres, B. Giner, I. Gascon, M. Castro, C. Lafuente, J. Phys. Chem. B 112 (2008) 12461–12467. [20] I. Bandres, B. Giner, H. Artigas, F.M. Royo, C. Lafuente, J. Phys. Chem. B 112 (2008) 3077–3084. [21] M.T. Zafarani-Moattar, H. Shekaari, J. Chem. Eng. Data 50 (2005) 1694–1699. [22] R. Garriga, F. Sanchez, P. Perez, M. Gracia, J. Chem. Thermodyn. 29 (1997) 649–659. [23] B. Gonzalez, A. Dominguez, J. Tojo, J. Chem. Thermodyn. 38 (2006) 1172–1185. [24] R. Garriga, S. Martinez, P. Perez, M. Gracia, Fluid Phase Equilib. 207 (2003) 97–109. [25] G.C. Benson, O. Kiyohara, J. Chem. Thermodyn. 11 (1979) 1061–1064. [26] H.G. Carlson, E.F. Westrum, J. Chem. Phys. 54 (1971) 1464–1471. [27] I. Klesper, Z. Phys. Chem. Frankfurt 51 (1966) 1–12. [28] A. Heintz, B. Schmittecker, D. Wagner, R.N. Lichtenthaler, J. Chem. Eng. Data 31 (1986) 487–492. [29] H. Katayanagi, K. Nishikawa, H. Shimozaki, K. Miki, P. Westh, Y. Koga, J. Phys. Chem. B 108 (2004) 19451–19457. [30] TRC Thermodynamics Table Hydrocarbons and Non-Hydrocarbons, Selected values of Properties of Chemical Compounds, Thermodynamic Research Center, Texas A&M University, College Station, TX. [31] W. Xu, L.M. Wang, R.A. Nieman, C.A. Angell, J. Phys. Chem. B 107 (2003) 11749–11756. [32] L. Santos, J.N.C. Lopes, J.A.P. Coutinho, J. Esperanca, L.R. Gomes, I.M. Marrucho, L.P.N. Rebelo, J. Am. Chem. Soc. 129 (2007) 284–285. [33] Y. Yoshida, O. Baba, G. Saito, J. Phys. Chem. B 111 (2007) 4742–4749. [34] T. Singh, A. Kumar, J. Phys. Chem. B 111 (2007) 7843–7851. [35] J.A. Riddick, W.B. Bunger, A. Weissberger, Organic Solvents; Physical Properties and Methods of Purification, Wiley-Interscience, New York, 1970. [36] J. Ortega, R. Vreekamp, E. Penco, E. Marrero, J. Chem. Thermodyn. 40 (2008) 1087–1094. [37] A. Heintz, D. Klasen, J.K. Lehmann, J. Solution Chem. 31 (2002) 467–476. [38] G. Garcia-Miaja, J. Troncoso, L. Romani, J. Chem. Eng. Data 52 (2007) 2261–2265. [39] I.M. Abdulagatov, A. Tekin, J. Safarov, A. Shahverdiyev, E. Hassel, Int. J. Thermophys. 29 (2008) 505–533.

JCT 10-236