Isothermal Vapor+ Liquid Equilibrium and Thermophysical Properties

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Isothermal Vapor+Liquid Equilibrium and Thermophysical Properties for 1-Butyl-3-methylimidazolium Bromide + 1-Butanol Binary System Mariana Teodorescu Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/ie502247d • Publication Date (Web): 08 Aug 2014 Downloaded from http://pubs.acs.org on August 17, 2014

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Industrial & Engineering Chemistry Research

Isothermal Vapor+Liquid Equilibrium and Thermophysical Properties for 1-Butyl-3methylimidazolium Bromide + 1-Butanol Binary System Mariana Teodorescu * ”Ilie Murgulescu” Institute of Physical Chemistry, Romanian Academy, Splaiul Independentei 202, 060021 Bucharest, Romania

ABSTRACT: Experimental isothermal vapor+liquid equilibrium (VLE) data are reported for the binary mixture containing 1-butyl-3-methylimidazolium bromide ([bmim]Br) + 1-butanol at three temperatures: (353.15, 363.15, and 373.15) K, in the range of 0-0.321 liquid mole fraction of [bmim]Br. Additionally, refractive index measurements have been performed at two temperatures: (298.15 and 308.15) K in the whole composition range. Densities, excess molar volumes, surface tensions and surface tension deviations of the binary mixture were predicted by Lorenz-Lorentz (nD-ρ) mixing rule. Dielectric permittivities and their deviations were evaluated by known equations. Vapor+liquid equilibrium data were correlated with Wilson thermodynamic model while refractive index data by the 3-parameters Redlich-Kister equation. The studied mixture presents positive abatement from the Raoult’s law. Similarly, positive refractive index deviations are obtained. The VLE data may be used in separation processes design, and the thermophysical properties as key parameters in specific applications.

INTRODUCTION Excellent properties (e.g. almost null volatility, generally zero flammability, high thermal stability1, low melting points, reasonably viscosity, high solubility and conductivity) and new applications of the ionic liquids (ILs) and their mixtures have made them very

*

Corresponding detailes: Tel: +40 213167912; Fax: +40 213121147;

E-mail address: [email protected]

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attractive materials in the last decade. ILs have been used, for example, as green entrainers with salting-out effect in extractive distillation, as highly efficient solvents in organic synthesis, and as catalysts for phase transfer. In the recent years, many authors have presented various aspects of these materials.1-8 A literature survey shows an increased trend in studies on the pure ILs and their mixtures (ion mobilities, glass temperatures and fragilities,3 different thermodynamic properties, 4 osmotic coefficients,7-10 and phase equilibria2,5,6). These studies help engineers to design, understand, and control the industrial processes including ionic liquids,10 and researchers to tailor their chemical structure so that to better match a given target.11 However, thermodynamic properties of the solutions of ILs in various solvents such as vapor pressure of the solutions and activity coefficients of the components have been more rarely studied.10 Thermodynamic data for IL-containing systems are essential for a better understanding of thermodynamic behaviors of such systems, for separation design purpose and for the development of thermodynamic models. The refractive index dependence on composition is generally used as calibration curves for phase composition determinations at pressure-temperature equilibrium data sets, by ebulliometry, when VLE data are measured. Since the refractive index of a liquid, n D, is a property easy to measure with good accuracy, it was related with other thermophysical or electrical properties, such as density, surface tension, and dielectric permittivity, by numerous empirical and theoretical equations as referenced in literature.12 Experimental measurements and prediction for these physical properties of IL mixtures are essential for application to many processes.13 The present work is a continuation of a project dealing with phase equilibria and thermophysical properties of ILs with classical solvent mixtures. In a previous paper,14 the system of a prototype IL (1-butyl-3methylimidazolium chloride) with 1-butanol has been investigated. For now, another prototype IL, 1-butyl-3methylimidazolium bromide, with the same alcohol binary system is under focus. No vapor+liquid equilibrium (VLE) data or thermophysical properties are available for it.15 Experimental isothermal VLE data are reported at three temperatures: (353.15, 363.15, and 373.15) K, in the range of 0-0.321 liquid mole fraction of [bmim]Br. The composition range was limited due to the reduced quantity of IL which it was at disposal. Additionally, refractive index measurements have been performed at two temperatures: (298.15 and 308.15) K in the whole composition range. Densities, excess molar volumes, surface tensions and surface tension deviations of the binary mixture were predicted by Lorenz-Lorentz (nD-ρ) mixing rule.16 Dielectric permittivities and their deviations were evaluated by known equations.17-18 The phase equilibria, thermodynamic and


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thermophysical data reported here will bring new information required for the design of the separation processes of 1-butanol from its mixtures using an organic salt as [bmim]Br and, at the same time, will allow an insight into specific inter and intra molecular interactions or structural arrangements existing in the binary IL + 1-butanol system.

EXPERIMENTAL Materials. The used chemicals of high purity were purchased from trustful commercial sources. The providers and characteristics of the used compounds are summarized in Table 1. After purification, both chemicals were stored in closed system to dried atmosphere, on calcium chloride. A good comparison with literature values has been obtained for the refractive index and density of the pure compounds. Also a good agreement was found for the vapor pressure and ultrasonic speed of sound of 1-butanol. This is shown in Table 2. For the present measurements, the supercooled liquid [bmim]Br compound with a estimative water content smaller than 0.001 mole fraction was used. Apparatuses and procedures. The vapor pressures measurements of pure 1-butanol and of the binary mixtures were carried out by dynamic method using a modified Swietoslawski ebulliometer.23 The apparatus, is described in details previously together with the experimental procedure24 commonly used.23,25,26 The temperatures at the thermodynamic equilibrium in the apparatus were measured with an uncertainty of ±0.1 K, by means of mercury thermometers calibrated in advance at National Institute of Metrology, Bucharest. The vapor pressure measurements were performed by using a mercury manometer. The pressure readings were made with a cathetometer with uncertainty of ±0.1 mm Hg, and its reproducibility was estimated to be better than 50 Pa. The uncertainty of the pressure measurements is estimated to be 0.1% of measured value. The composition of the liquid phase in equilibrium with the vapor phase was analyzed by the refractometric method making use of a calibration curve obtained by measurements of the refractive index of weighed mixture samples (uncertainty ±0.2 mg by GH-252 A&D Japan balance) at 298.15 K and data correlation with Redlich-Kister polynomials27 with three parameters in the form:







n DE  n D  xn D ,1  1  x n D ,2  x1  x  a 0  a1 1  2 x   a 2 1  2 x 

2



(1)

The Redlich-Kister parameters a 0, a1, a 2 of eq. (1) were obtained by maximum likelihood optimization method using the following objective function:



 x  x 2 n E  nE OF    i ,exp 2 i ,calc  D ,i ,exp 2 D ,i ,calc x  nE i 1  D  N

  2




(2)


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The refractive index was measured by a digital refractometer Abbemat RXA 170 from Anton-Paar (Austria) at the wavelength of the D line of sodium, 589.3 nm, with uncertainty ±0.0001. The temperature of the Safire prism was controlled by a Peltier element to within ±0.01 K and the calibration of the apparatus was carried out with bidistilled and deionized water and by determining refractive indexes and its deviations at 298.15 K for the binary cyclooctane + toluene system at several compositions. The difference resulted from the comparison of our refractive index deviation values, nDE , and those correlated on experimental data from literature28 was 2·10 -4 in mean absolute average, which it means a good comparison. The nDE , which is defined as n DE  n D  xn D ,1  1  x n D ,2

(3)

has an estimated uncertainty of 2·10-4. In eqs. (1) and (3), nD ,1 and nD ,2 represent the refractive indexes of the pure components 1 and 2, nD the refractive index for the mixture 1+2, and x denotes the mole fraction of component 1 in the binary mixture. Each measurement of refractive index consisted in 4 readings for the same sample. Their averaged value is reported here as one experimental point. Special attention was devoted to avoid extra moisture absorption from the atmosphere in pure chemicals and in their binary mixtures. The uncertainty in the determination of the composition was 0.001 mole fraction. The VLE apparatus and experimental procedure were successfully checked and used for investigation of different other mixtures as mentioned before.29 The density and ultrasonic speed of sound measurements of the pure compounds at the two temperatures (298.15, and 308.15) K were carried out by using a density and speed of sound meter Anton Paar DSA-5000 M with precision of ±0.005 kg m-3 and ±0.01 m s-1. Dried air and distilled deionized ultra pure water at atmospheric pressure were used as calibration fluids for the cell. The probes thermostating was maintained constant at ±0.01 K. The experimental measurement uncertainty for density was 0.1 kg m-3 and for the speed of sound 0.1 m s-1.

RESULTS AND DISCUSSIONS The vapor pressures of pure 1-butanol compound measured in the ebulliometer at the three temperatures are given in Table 2 together with literature values. Good agreement can be observed. The experimental isothermal (P, x) data measured for the binary system at (353.15, 363.15, and 373.15) K are shown in Figure 1.


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The data were correlated with Wilson thermodynamic model30 with the following basic equations:

GE   x1 lnx1   12 x2   x2 ln x 2   21 x1  RT

(4)

  12  21 ln  1   ln  x1   12 x2   x2    x1   12 x2  21 x1  x 2

  

(5)

  12  21 ln  2   ln x2   21 x1   x1    x1   12 x 2  21 x1  x 2

  

(6)

 12 

V2    11  exp  12  V1 RT  

(7)

 21 

V1     22  exp  12  V2 RT  

(8)

where xi is the liquid mole fraction of component i (i = 1, 2), V is the molar volume, γ the activity coefficient, GE excess Gibbs molar energy in J mol-1, (λij - λjj) are the model parameters expressed in J mol-1, and  12 ,  21 are the dimensionless Wilson parameters. The regression for the obtaining of the model parameters was performed by means of Barker method31 employing a program described by Hala et al..32 The objective function is defined as it follows:  P  Pi ,e S=   i ,c Pi ,e i 1  N

   

2

(9)

where N is the number of experimental points, Pi,e, Pi,c are the experimental and calculated vapor pressures for the experimental point i, respectively. All standard deviations of correlations were calculated using the expression:

Z 

 ( Z

i ,e



 Z i ,c ) 2 /( N  m )

1/ 2

, where Z is the value of the property P, n DE , and x, N is

the number of experimental points and m = 2 in the case of Wilson model or m = 3 in the case of Redlich-Kister equation. In the case of VLE data, the real behavior of the vapor phase was described with the virial equation of state. The second virial coefficients for both components and for binary mixture were evaluated by means of the Hayden and O’Connell method33, while the molar volumes were calculated by using a generalized Watson relation.34 For the calculations, the experimentally determined vapor pressures of pure 1-butanol (as given in Table 2) and those estimated for [bmim]Br by Ambrose-Walton method35 were used. The required normal



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boiling point, critical temperature and pressure as well as acentric factor for [bmim]Br have been estimated by Joback method.36 In Figure 1 it can be seen a good agreement between experimental and calculated bubble curves. In the range of the studied composition, the excess Gibbs energy, GE, is positive and doesn’t show an obvious temperature variation. This can be observed in the Supporting Information. At approximate 0.140 mole fraction of [Bmim]Br in the binary mixture and at average temperature of 363.15 K, the excess enthalpy, HE, calculated from the Gibbs-Helmholtz equation, is 2401 J mol-1. No experimental calorimetric value has been found in the literature for comparison. However, since the VLE data were correlated successfully with Wilson thermodynamic model without showing systematic errors, we can say that the VLE data are thermodynamic consistent. The experimental isothermal (x, n D) data for the binary system have been determined at (298.15 and 308.15) K and they are shown in Supporting Information with the calculated refractive index deviations. The correlation of these data at each temperature has been made by 3-parameters Redlich-Kister model by using maximum likelihood method, following the procedure described for 298.15 K in the section Method. The correlation results are presented in the Supporting Information and the variation of the refractive index deviation vs. composition is shown in Figure 2. Due to the small water content of the [bmim]Br ionic liquid, expected significant values for the standard deviations of the liquid compositions have been obtained. It should be mentioned that not complete purification of the organic salt is possible. Similar situation was reported in the literature.37,38 From Figure 2 it can be observed that excess refractive indexes are positive on whole composition interval and increase slightly with increasing temperature. Similar behavior was reported previously for [bmim]Cl with 1butanol system where the same type of interactions are present.14 From refractive index vs. composition data and densities of the pure compounds at (298.15 and 308.15) K, densities, excess molar volumes, surface tensions and dielectric permittivities at optical frequency and their deviations vs. composition of the binary mixture were determined by Lorenz-Lorentz mixing rule16 or by known equations.17,18 The selection of the mixing rule was made after analyzing of ten different mixing rules results in giving the refractive indices from experimental densities for twelve binary mixtures of various polarity at 298.15 K.12 The predicted variation of the excess molar volume with composition and temperature for the binary [bmim]Br (1) + 1-butanol (2) system is shown in Figure 3. As can be observed in this Figure, the VE is negative on whole composition interval, it is more negative when the


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temperature increases, and similar behavior is given by applying other good mixing rules like Gladstone-Dale, Edwards and Eykman.39 [Bmim]Br ionic liquid can act both as a hydrogenbond acceptor ([Br]-) and donor ([bmim]+). It is expected to interact with 1-butanol which has both accepting and donating sites. On the other hand, it is well known that 1-butanol is hydrogen-bonded solvent with both high enthalpies of association and association constant. Hence, it is expected to stabilize [bmim]Br with hydrogen-bonded donor sites. Similar to [bmim]Cl, the [bmim]Br is a complicated and highly interacting molecule, especially when the other compound of IL mixture is water or alcohol.40 At constant temperature, the excess molar volumes for [bmim]Br + 1-butanol system calculated by Lorenz-Lorentz mixing rule are smaller than those of [bmim]Cl with the same alcohol.14 This is due probably to the fact that Br- is less electronegantive than Cl-. For each mentioned binary system, the temperature effect is not specific to new hydrogen bonds between unlike molecules formation; it is most likely due to the packing between hydrogen acceptor sites of 1-butanol and donor sites of imidazolium ([bmim]+) cycle. These packing effects became more dominant and increase with temperature, as it was observed for other systems cited in literature.41 The surface tension of a liquid is a property of great importance in mass transfer processes such as distillation, extraction, or absorption. It is not easily measured, and considerable attention has been paid to the development and analysis of equations allowing its prediction from properties for which data are more readily available as refractive index, for example. The surface tension σ is related to the densities of the liquid ρL and of the vapor ρV phases of the substance by using the Macelod equation17:

  ct  L  V 

4

(10)

For a pure liquid compound, multiplying both sides by the molar mass M and ignoring ρ V in comparison with ρL affords the Sugden equation18:

 1 / 4 M /  L   1 / 4Vm  Parachor  Parachor   from which it results:    Vm  

(11) 4

(12)

In eqs. (11) and (12), the Parachor is assumed as mole-wise additive and molar volume, Vm, is calculated from experimental densities of pure compounds and those predicted for the mixtures by the Lorenz-Lorentz relation16 (from Redlich-Kister correlated refractive indices). So, we used eq. (12) to predict the surface tensions of binary liquid mixtures at two temperatures (298.15 and 308.15) K and after that to calculate the surface tension deviation from ideality by applying a similar equation with eq. (3). Surface tension variation with



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composition and temperature is shown in Figure 4 and surface tension deviation variation with the same two variables appears in Figure 5. The parachors of the pure compounds have been predicted by additive group contribution methods of Tyn and Calus.42 The surface tensions of pure 1-butanol compares well with those determined experimentally in literature43 with mean absolute relative deviation of 2.8 % on the two temperatures examined in this work. The calculated surface tensions for pure [bmim]Br are about 10 % lower than those brutish extrapolated from literature experimental data.44 As can be seen in Figure 5, the surface tension deviations predicted here are positive on whole composition range being slightly higher at higher temperature. This is explained by the volume expansion resulted from the new H-bonds and packing between unlike molecules in the binary mixture of [bmim]Br + 1-butanol. It seems that packing of the molecules is dominant since excess surface tension is higher at higher temperature. Similar behavior was previously found for the [bmim]Cl + 1-butanol binary system where the same type of interactions and structural arrangements were expected.14 From refractive index data, the relative permittivity at optical frequency can be calculated by squaring the refractive index determined at the wavelength of 589.3 nm. For 1butanol, the relative permittivities agree well with those extrapolated at the two temperatures of (298.15 and 308.15) K from experimental literature data45 with mean absolute relative deviation of 0.3 %. No data have been found for [bmim]Br for comparison. The dielectric permittivity (Figure 6) at optical frequency is obtained by multiplying relative permittivity by vacuum permittivity. The excess permittivity (Figure 7) calculated by a similar equation like (3) is found to be positive on whole composition interval and slightly increase with increasing temperature with a maximum around 0.55 10 -12 F m-1 and 0.400 mole fraction of [bmim]Br. This indicates again that both unlike molecular species interact in such a way that they act as more H-bonded structure than those of pure compounds themselves.

CONCLUSIONS The binary 1-butyl-3-methylimidazolium bromide + 1-butanol system has been investigated isothermally at vapor+liquid equilibrium at three temperatures: (353.15, 363.15, and 373.15) K. By correlation of the experimental (P, T, x) data with Wilson thermodynamic model it was found that the system is zeotropic with high positive deviations from ideality. Excess molar enthalpy calculated from the excess molar Gibbs energy temperature dependence by Gibbs-Helmholtz equation is positive and high, indicating specific interactions between IL and organic solvent molecules. Both GE and HE have smaller values than for the


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binary [bmim]Cl + 1-butanol system14 due to the smaller strength of the intermolecular interactions between [bmim]Br and 1-butanol molecules at mixing. For the same system, refractive index measurements have been performed at two temperatures (298.15 and 308.15) K in whole composition interval. Excess refractive index is positive at both temperatures on whole composition range and increase slightly with increasing temperature. Using the Lorenz-Lorentz (n D-ρ) mixing rule, the densities, excess molar volumes, surface tensions and surface tension deviations have been predicted. By other known relations, dielectric permittivity and its deviations have been also calculated at the same temperatures of (298.15 and 308.15) K. Structural effects for the binary investigated mixture have been explained in terms of excess thermodynamic and thermophysical properties.

ACKNOWLEDGMENTS This contribution was carried out within the research programme “Chemical thermodynamics and kinetics. Quantum chemistry” of the “Ilie Murgulescu” Institute of Physical Chemistry, financed by the Romanian Academy. Support of the EU (ERDF) and Romanian Government, which allowed for the acquisition of the research infrastructure under POS-CCE O 2.2.1 Project INFRANANOCHEM - Nr. 19/01.03.2009, is gratefully acknowledged.

SUPPORTING INFORMATION AVAILABLE VLE data, activity coefficients and excess Gibbs energy for the [bmim]Br + 1-butanol system, Wilson parameters for the VLE data generalization, refractive index, and its deviation from ideality vs. composition for the same system together with Redlich-Kister model parameters obtained from their correlation. This information is available free of charge via the Internet at http://pubs.acs.org/.



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(13) Hanke, C. G.; Atamas, N. A.; Lynden-Bell, R. M. Solvation of Small Molecules in Imidazolium Ionic Liquids: A Simulation Study. Green Chem. 2002, 4, 107. (14) Teodorescu, M. Isothermal Vapour+Liquid Equilibrium and Thermophysical Properties for 1-Butyl-3-methylimidazolium Chloride + 1-Butanol Binary System. Rev. Chim. (Bucharest), 2014, in press. (15) DDBST (Software and Separation Technology) GmbH, The Dortmund Data Bank (DDB), Oldenburg, Germany, May 2014, www.ddbst.com (16) Lorentz, H. A. Theory of electrons; B. G. Teubmer (Ed.): Leipzig, 1909. (17) Fowler, R. H. A Tentative Statistical Theory of Macleod's Equation for Surface Tension, and the Parachor. Proc. Roy. Soc. Lond. A 1937, 159, 229. (18) Sugden, S. A Relation Between Surface Tension, Density, and Chemical Composition. J. Chem. Soc., Trans. 1924, 125, 1177. (19) Kim, K. -S.; Shin, B. -K.; Lee, H. Physical and Electrochemical Properties of 1-Butyl-3methylimidazolium

Bromide,

1-Butyl-3-methylimidazolium

Iodide,

and

1-Butyl-3-

methylimidazolium Tetrafluoroborate. Koreean J. Chem. Eng. 2004, 21(5), 1010. (20) Kim, K. -S.; Shin, B. -K.; Lee, H.; Ziegler, F. Refractive Index and Heat Capacity of 1Butyl-3-methylimidazolium Bromide and 1-Butyl-3-methylimidazolium Tetrafluoroborate, and Vapor Pressure of Binary Systems for 1-Butyl-3-methylimidazolium Bromide + Trifluoroethanol and 1-Butyl-3-methylimidazolium Tetrafluoroborate + Trifluoroethanol. Fluid Phase Equilib. 2004, 218, 215. (21) Singh, S.; Aznar, M.; Deenadayalu, N. Densities, Speeds of Sound, and Refractive Indices for Binary Mixtures of 1-Butyl-3-methylimidazolium Methyl Sulphate Ionic Liquid with Alcohols at T = (298.15, 303.15, 308.15, and 313.15) K. J. Chem. Thermodyn. 2013, 57, 238. (22) Boublik, T.; Fried, V.; Hala, E. The Vapour Pressures of Pure Substances; Elsevier: Amsterdam, 1973. (23) Rogalski, M.; Malanowski, S. Ebulliometers Modified for the Accurate Determination of Vapour—Liquid Equilibrium. Fluid Phase Equilib. 1980, 5, 97. (24) Barhala, A.; Teodorescu, M. Measurement of Vapour-Liquid Equilibria by Ebuliometry. Rev. Roum. Chim. 2001, 46(9), 967. (25) Malanowski, S. Experimental Methods for Vapour-Liquid Equilibria. Part I. Circulation Methods. Fluid Phase Equilib. 1982, 8, 197. (26) Olson, J. D. Measurement of Vapor-Liquid Equilibria by Ebulliometry. Fluid Phase Equilib. 1989, 52, 209.



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(27) Redlich, O.; Kister, A. T. Algebraic Representation of Thermodynamic Properties and the Classification of Solutions. Ind. Eng. Chem. 1948, 40, 345. (28) Gonzales, B.; Dominguez, I.; Gonzales, E. J.; Dominguez, A. Density, Speed of Sound, and Refractive Index of the Binary Systems Cyclohexane (1) or Methylcyclohexane (1) or Cyclooctane (1) with Benzene (2), Toluene (2), and Ethylbenzene (2) at Two Temperatures. J. Chem. Eng. Data 2010, 55, 1003. (29) Teodorescu, M.; Dragoescu, D.; Gheorghe, D. Isothermal (Vapour + Liquid) Equilibria for (Nitromethane or Nitroethane + 1,4-Dichlorobutane) Binary Systems at Temperatures Between (343.15 and 363.15) K. J. Chem. Thermodyn. 2013, 56, 32. (30) Wilson, G. M. Vapor-Liquid Equilibrium. XI. A New Expression for the Excess Free Energy of Mixing. J. Am. Chem. Soc. 1964, 86, 127. (31) Barker, J. A. Determination of Activity Coefficients From Total Pressure Measurements. Aust. J. Chem. 1953, 6, 207. (32). Hala, E.; Aim, K.; Boublik, T.; Linek, J.; Wichterle, I. Vapor-Liquid Equilibrium at Normal and Reduced Pressures, Academia: Praha, 1982, (in Czech). (33) Hayden, J. G.; O’Connell, J. P. A Generalized Method for Predicting Second Virial Coefficients. Ind. Eng. Chem., Proc. Des. Develop. 1975, 14, 209. (34) Hougen, O. A.; Watson, K. M. Chemical Process Principles, Part II; J. Wiley: New York, 1947. (35) Ambrose, D.; Walton, J. Vapour Pressures up to Their Critical Temperatures of Normal Alkanes and 1-Alcohols. Pure Appl. Chem., 1989, 61(8), 1395. (36) Joback, K. G.; Reid, R. C. Estimation of Pure-Component Properties from GroupContributions. Chem. Eng. Commun., 1987, 57, 233. (37) Zafarani-Moattar, M. T.; Shekaari, H. Apparent Molar Volume and Isentropic Compressibility of Ionic Liquid 1-Butyl-3-methylimidazolium Bromide in Water, Methanol, and Ethanol at T = (298.15 to 318.15) K. J. Chem Thermodyn. 2005, 37, 1029. (38) Sadeghi, R.; Golabiazar, R.; Shekaari, H. The Salting-out Effect and Phase Separation in aqueous Solutions of Trisodium Citrate and 1-Butyl-3-methylimidazolium Bromide. J. Chem. Thermodyn. 2010, 42, 441. (39) Baraldi, P.; Giorgini, M. G.; Manzini, D.; Marchetti, A.; Tassi, L. Density, Refractive Index, and Related Properties for 2-Butanone + n-Hexane Binary Mixtures at Various Temperatures. J. Sol. Chem. 2002, 31(11), 873. (40) Domanska, U.; Bogel-Eukasik, E.; Bogel-Eukasik, R. 1-Octanol/Water Partition Coefficients of 1-Alkyl-3-methylimidazolium Chloride. Chem. Eur. J. 2003, 9, 3033.


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(41) Valtz, A.; Teodorescu, M.; Wichterle, I.; Richon, D. Liquid Densities and Excess Molar Volumes for Water + Diethylene Glycolamine, and Water, Methanol, Ethanol, 1-Propanol + Triethylene Glycol Binary Systems at Atmospheric Pressure and Temperatures in the Range of 283.15-363.15 K. Fluid Phase Equilib. 2004, 215, 129. (42) Poling, B. E.; Prausnitz, J. M.; O’Connell, J. P. The properties of gases and liquids, 5th Edition; McGraw Hill: New York, USA, 2001. (43) Jiang, H.; Zhao, Y.; Wang, J.; Zhao, F.; Liu, R.; Hu, Y. Density and Surface Tension of Pure Ionic Liquid 1-Butyl-3-methyl-imidazolium l-Lactate and Its Binary Mixture With Alcohol and Water. J. Chem. Thermodyn. 2013, 64, 1. (44) Fu, D.; Wang, H. M.; Du, L. X. Experiments and Model for the Surface Tension of (MDEA + [Bmim][BF4]) and (MDEA + [Bmim][Br]) Aqueous Solutions. J. Chem. Thermodyn. 2014, 71, 1. (45) Rana, V. A.; Chaube, H. A. Static Permittivity, Density, Viscosity and Refractive Index of Binary Mixtures of Anisole with 1-Butanol and 1-Heptanol at Different Temperatures. J. Molec. Liq. 2012, 173, 71.



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Table 1. Commercial sources, purities and methods of purification of the used chemical compounds Compound

[bmim]Br

1-butanol

Commercial Purity / source mass fraction Fluka > 0.970

Purification method

Riedel de Häen

Dried and stored on molecular sieves 4Å

> 0.995

Dried in the oven at 70 0C and 0.1 kPa for 10 days

Table 2. Refractive indices, nD, densities, ρ, and ultrasonic speeds of sound, u, at atmospheric pressure and vapor pressures, P, of pure compounds nDa

ρb / (kg m-3) This Lit. work

Compound / Te / K [bmim]Br/ 298.15

This work

Lit.

1.5353

308.15

1.5326

1.545019 1.5420 1.542019 1.5420

uc / (m s-1) This Lit. work

1290.9

130019

1664.4

1284.3

129019

1642.1

Pd / kPa This Lit. work

1-butanol / 298.15 1.3973 1.3974721 805.6 805.9321 1240.6 1239.2821 308.15 1.3932 1.3934221 798.0 798.2121 1207.1 1206.2621 353.15 21.24 21.8022 363.15 33.85 34.1922 373.15 51.69 51.8922 a unD = 0.0001, buρ = 0.1 kg m-3, cuu = 0.1 m s-1, duP = 0.1% of measured value, and euT = 0.1 K.


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Caption to figures Figure 1. Bubble curves for [bmim]Br (1) + 1-butanol (2) system at 353.15 K (), 363.15 K (▲), and 373.15 K (■); Filled symbols: experimental bubble curves (pressure P vs. liquid composition in [bmim]Br x); () Wilson correlation of the bubble curves with real behavior assumption of the vapor phase.

Figure 2. Refractive index deviations n DE vs. composition in [bmim]Br x for [bmim]Br (1) + 1-butanol (2) system at 298.15 K () and 308.15 K (▲); Filled symbols are calculated values by Eq. (3); Lines are 3-parameters Redlich-Kister correlation for 298.15 K () and 308.15 K (---). Figure 3. Prediction of excess molar volumes VE vs. composition in [bmim]Br x for [bmim]Br (1) + 1-butanol (2) system at 298.15 K () and 308.15 K (- - -) by different mixing rules: (○) Lorenz-Lorentz, (Δ) Gladston-Dale, (□) Edwards, and (◊) Eykmann.

Figure 4. Prediction of surface tensions σ vs. composition in [bmim]Br x for [bmim]Br (1) + 1-butanol (2) system at 298.15 K (○) and 308.15 K (Δ). Figure 5. Prediction of excess surface tensions σE vs. composition in [bmim]Br x for [bmim]Br (1) + 1-butanol (2) system at 298.15 K (○)and 308.15 K (Δ).

Figure 6. Prediction of dielectric permittivities at optical frequency ε vs. composition in [bmim]Br x for [bmim]Br (1) + 1-butanol (2) system at 298.15 K (○) and 308.15 K (Δ). Figure 7. Prediction of excess dielectric permittivities at optical frequency εE vs. composition in [bmim]Br x for [bmim]Br (1) + 1-butanol (2) system at 298.15 K (○) and 308.15 K (Δ).



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16

55 50 45

P / kPa

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

40 35 30 25 20 0.00

0.10

0.20

0.30

x

Figure 1.


0.40

Page 17 of 23


17

0.025

0.020

0.015

n DE

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

0.010

0.005

0.000 0.0

0.2

0.4

0.6

0.8

x

Figure 2.


1.0


Page 18 of 23

18

0.0

0.2

0.4

0.6

0.8

0.0 -0.1 -0.2

V E 106 / (m3 mol-1)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

-0.3 -0.4 -0.5 -0.6 -0.7 -0.8 -0.9 -1.0

x

Figure 3.


1.0

Page 19 of 23


19

45

40

-1 3  10 / (N m )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

35

30

25

20 0

0.2

0.4

0.6

0.8

x

Figure 4.


1


Page 20 of 23

20

3.0

2.5

E 3 -1  10 / (N m )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

2.0

1.5

1.0

0.5

0.0 0.0

0.2

0.4

0.6

0.8

x

Figure 5.


1.0

Page 21 of 23


21

21.5 21.0 20.5 12 -1 e 10 / (F m )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

20.0 19.5 19.0 18.5 18.0 17.5 17.0 0.0

0.2

0.4

0.6

0.8

x

Figure 6.


1.0


Page 22 of 23

22

0.60

0.50 E 12 -1 e 10 / (F m )

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

0.40

0.30

0.20

0.10

0.00 0

0.2

0.4

0.6

0.8

x

Figure 7.


1

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Graphical Abstract for Table of Contents Only


Isothermal Vapor+ Liquid Equilibrium and Thermophysical Properties

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