Thermodynamic Behavior of Binary Mixtures of 1-Butyl-1

Article pubs.acs.org/jced

Thermodynamic Behavior of Binary Mixtures of 1‑Butyl-1methylpyrrolidinium Iodide and Alcohols Sayeed Ashique Ahmed, Aninda Chatterjee, Banibrata Maity, and Debabrata Seth* Department of Chemistry, Indian Institute of Technology Patna, Patna 800013, Bihar, India ABSTRACT: In this work, the osmotic properties of the binary mixtures of 1butyl-1-methylpyrrolidinium iodide ([BmPr][I]) with isopropanol and nbutanol were studied by using vapor pressure osmometry (VPO) technique at 323.15 K at atmospheric pressure. It was found that the values of osmotic coefficient, mean molal activity coefficient, and vapor pressure decrease with increasing molality. We observed that the change in vapor pressure of ionic liquid in n-butanol is smaller compared to ionic liquid in isopropanol. The excess Gibbs free energy of ionic liquid in isopropanol has a higher value as compared to an ionic liquid + n-butanol mixture. The binary mixture of ionic liquid with alcohols shows noticeable deviation from the Debye−Huckel limiting law. We have used the pseudophase separation model on the experimental osmotic coefficient value to estimate the critical aggregation concentration of ionic liquid in alcohols. We have showed the structural effect of the solvent on the preceding thermodynamic parameters and critical aggregation concentration value of the ionic liquid in alcohols.

1. INTRODUCTION Ionic liquids (ILs) are nonvolatile organic salts with very interesting features such as very low vapor pressure, purely ionic character, low melting temperature, comparably high thermal stability at normal pressure and temperature, and being nonflammable. Ionic liquids are used as green solvents in many research and industrial processes.1−3 For these properties of ionic liquid, presently it is a very interesting research area. To use ionic liquids for any specific purpose in research and industrial processes, systematic knowledge on the different thermodynamic as well as transport properties such as osmotic coefficients, activities, activity coefficients, viscosities, surface tensions, and electrical conductivities are necessary. To observe the deviation of solution from its ideal behavior, osmotic coefficient, activity, and excess free energy of the solution are of great interest. Knowledge on activity coefficient and osmotic coefficient are very important for IL containing media to design chemical processes, such as separation processes, extractive distillation, or salt distillation. For deeper understanding of the thermodynamic properties of IL solution, knowledge of different types of interactions at the molecular level, i.e., ionic liquid−solvent interaction, cation−cation association, and cation−anion interaction, etc., is necessary. In this work, we have reported the osmotic coefficient, activity, mean molar activity coefficient, vapor pressure, excess Gibbs free energy, critical aggregation concentration (cac) of IL + alcohol (isopropanol and n-butanol) systems at T = 323.15 K, and the probable reasons for the variation of these parameters with concentration were explained. We have explained the effect of solvent medium on these thermodynamic properties on the basis of different types of interactions at the molecular level. These types of interaction can be investigated by using © XXXX American Chemical Society

headspace chromatography, membrane osmometry, direct vapor measurement, isopiestic method, and vapor pressure osmometry (VPO). Among all of these methods, we choose VPO technique due to its numerous advantages such as high accuracy, use over wide concentration ranges, wide ranges of solvents, short analysis time, and higher detection limit (a small amount of sample is enough to proceed with the experiment). Knowledge on vapor−liquid equilibrium (VLE) of the binary mixtures of IL with different solvents is necessary before we can use IL for any specific purpose. The aim of this work is to present the variation of different types of thermodynamic parameters of binary solution of 1-butyl-1-methylpyrrolidinium iodide ([BmPr][I]) with isopropanol and n-butanol with concentration as well as to observe the effect of solvent medium. Osmotic coefficient, ϕ, activity, a, activity coefficient, γ ± , depression of vapor pressure, Δp, and excess free energy, GE were determined. Knowledge on such properties is necessary to know the vapor−liquid equilibrium (VLE) as well as thermodynamic behavior of the binary mixture of ionic liquid−solvent medium. The osmotic coefficient values are correlated by using the Archer extension of the Pitzer ion interaction model.4−6 The Archer extension of the Pitzer ion interaction model parameters were obtained by data fitting. In this work we showed the solvent effect on different thermodynamic parameters. The osmotic properties of binary mixtures containing different IL have been reported in the literature,7−28 other than the binary mixtures of 1-butyl-1methylpyrrolidinium iodide + isopropanol and 1-butyl-1Received: February 27, 2015 Accepted: July 14, 2015

A

DOI: 10.1021/acs.jced.5b00184 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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methylpyrrolidinium iodide + n-butanol studied in this work. We have estimated the value of the cac and the critical aggregation numbers, n, of 1-butyl-1-methylpyrrolidinium iodide [BmPr][I] in isopropanol and n-butanol.

ϕ=

νrmr ϕ νm r

(1)

where ϕ, ν, and m are the osmotic coefficient of the IL solution, the stoichiometric number of the IL electrolyte, and the molality of the IL solution, respectively. ϕr, νr, and mr are the osmotic coefficient and stoichiometric number of the reference electrolyte and the molality showing same instrumental reading as the IL solution, respectively. We have used LiBr as the reference electrolyte.

2. EXPERIMENTAL SECTION 2.1. Materials. 1-Butyl-1-methyl pyrrolidinium iodide (high purity grade) was supplied from Sigma-Aldrich (Figure 1). 1-

3. RESULTS AND DISCUSSION 3.1. Osmotic Coefficient and Thermodynamic Modeling. The experimental osmotic coefficient values of [BmPr][I] + isopropanol and [BmPr][I] + n-butanol systems were calculated at T = 323.15 K by using mathematical expression 1. The variation osmotic coefficient values, ϕ, with molality of the solution are shown in Table 2 and Figure 2. ϕ gradually

Figure 1. (a) 1-Butyl-1-methylpyrrolidinium iodide [BmPr][I−], (b) isopropanol, and (c) n-butanol.

Table 2. Variation of Osmotic Coefficients, ϕ, with Molality, m, of the Binary Solution of Isopropanol + [BmPr][I] and nButanol + [BmPr][I] Systems at Temperature, T,a and Atmospheric Pressure

Butyl-1-methylpyrrolidinium iodide was dried under vacuum for 10 h before use. Spectroscopy graded isopropanol was received from Spectrochem, Mumbai, India, and n-butanol was received from Acros Organic. These solvents were used throughout the experiment for the preparation of solutions. All of the solutions were kept tightly sealed, and measurements were performed quickly after solution preparation to avoid the evaporation of solvent or absorption of moisture by the IL. 2.2. Apparatus and Procedure. The water content in the IL was measured by using Karl Fischer titration technique (model Metrohm 831 KF Coulometer). The water content in the IL was 0.3% (Table 1). Weighing and preparation of the solution were carried out by using a glovebox (nitrogen atmosphere) to minimize the absorption of atmospheric moisture as the IL is moisture sensitive. Each solution was prepared in a sealed vial, and each vial was weighed by an analytical balance from Mettler Toledo with an uncertainty of 10−4 g (model ML-204). The vapor pressure osmometry measurements were carried out by using a Knauer K-7000 vapor pressure osmometer (VPO) instrument. The measurements were carried out at a temperature 323.15 K and at atmospheric pressure. The principle of this technique is that the solution and the solvent droplets are generated directly on two thermistors present in the solvent chamber. Two thermistors are present in a Wheatstone bridge circuit. This bridge circuit thermistors measure resistance change, ΔR, due to the variation in temperature, ΔT. It indicates that ΔT measures in terms of ΔR. We performed the measurements at least six times for each solution, and the mean value is reported. The standard uncertainty in the output of the instrument was 1 Ω. The procedure has been described in detail previously.8 The osmotic coefficient, ϕ, value of the studied systems of IL having molality m can be expressed in the following form:6

[BmPr][I] + isopropanol

[BmPr][I] + n-butanol

T

mb

mc

K

(mol·kg−1)

ϕ

(mol·kg−1

ϕ

323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

0.0529 0.1022 0.1978 0.3344 0.5053 0.7375 1.004 1.2316 1.5529 1.839 2.2383

0.728 0.641 0.584 0.543 0.491 0.452 0.427 0.402 0.353 0.341 0.340

0.0483 0.1040 0.2081 0.2954 0.5164 0.7700 0.9771 1.2427 1.5288 1.8715 2.2662

0.714 0.640 0.560 0.511 0.458 0.402 0.384 0.336 0.330 0.327 0.314

Standard uncertainties u are u(m) = 0.001 mol·kg−1, u(T) = 0.01 K, and the combined expanded uncertainty Uϕ is = 0.01 (0.95 level of confidence). bSolvent = isopropanol, cSolvent = n-butanol.

a

decreased with the increasing molality of the IL in solution at T = 323.15 K. The experimental ϕ of the [BmPr][I] + isopropanol and [BmPr][I] + n-butanol systems were correlated by using the Archer extension of the Pitzer model, using nonlinear least-squares method.4−6 The model has the following form for a binary 1:1 electrolyte solution: ϕ − 1 = f ϕ + mBϕ + m2C ϕ

(2)

where

Table 1. Specifications of Chemical Samples mass fraction purity as per the supplier

a

water content (mass fraction)

sample

source

%

purification method

%

analysis method

[BmPr][I] isopropanol n-butanol

Sigma-Aldrich Spectrochem ACROS Organic

98.3 99.8 99 +

dried in vacuum none none

0.3 0.06 0.06

KF titrationa KF titrationa KF titrationa

Karl Fischer titration method. B



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Figure 2. . Experimental osmotic coefficients, ϕ, plotted against molality, m, at T = 323.15 K and atmospheric pressure: (a) isopropanol}; (b) ○, {[BmPr][I] + n-butanol}.

▲,

{[BmPr][I] +

Table 3. Fit Parameters of the Extended Archer Extension of Pitzer Model for {[BmPr][I] + Isopropanol} and {[BmPr][I] + nButanol} Systems Together with the Standard Deviation, σ, at Temperature, T, and Atmospheric Pressure system

T/K

β(0)

β(1)

β(2)

C(0)

C(1)

σ

[BmPr][I] + isopropanol [BmPr][I] + n-butanol

323.15 323.15

−0.3758 −0.2779

4.9407 5.3338

52.4860 57.1878

0.1310 0.1165

222.7235 381.2861

0.010 0.008

fϕ =

Aϕ =

In the preceding expressions 2 to 6, b, α1, α2, and α3 are the adjustable parameters. The values of b, α1, α2, and α3 are reported in the literature, the yields of which are reliable as well as good fits for some electrolyte solutions.9,16,17 To obtained the correlation parameters, we choose sets of Archer extended Pitzer ion interaction model parameters of the studied systems: b = 1.2 kg1/2·mol−1/2, α1 = 1.4 kg1/2·mol−1/2, α2 = 14 kg1/2· mol−1/2, and α3 = 10 kg1/2·mol−1/2. The correlation parameters are shown in Table 3. In our previous study, we obtained very good correlation by using the preceding adjustable parameters.16 From Figure 2, it is observed that with increasing molality of the solution, ϕ gradually decreased. It is well known behavior of the osmotic coefficient of an electrolyte in solvent medium. ϕ depends on the following factors: (i) ϕ decreases with increasing molality of the solution; (ii) ϕ increases with increasing ionic radius, as the repulsion between the ions is higher; (iii) ϕ increases with increasing solvation; and (iv) ion association decreases ϕ.28 In our study, the osmotic coefficient value decreases with increasing molality of the solution. With increasing concentration of the solution, ion association takes place in both solvents, which reduces the osmotic coefficient value.15,29 The value of the osmotic coefficient for isopropanol is greater as compared to that of n-butanol at each concentration, due to the greater solvation of IL in isopropanol as compared n-butanol. Solvation of IL in n-butanol is lower as compared to isopropanol due to presence of a short nonpolar chain within n-butanol, whereas it is absent in isopropanol; for this reason the interaction of IL with a more polar solvent, i.e., isopropanol, is greater as compared to that with the less polar nbutanol. Hence, the osmotic coefficient value for isopropanol is greater than n-butanol for the studied ionic liquid. The terms β(0) and β(1) represent the presence of interaction between like and unlike charges and the short-range interactions between unlike charged ions. The value of β(0) is found to be negative, whereas the value of β(1) is found to be positive by using the Archer extension of the Pitzer model. It is suggested that the interactions between unlike charged ions are predominant over like charged ions. The presence of Cϕ (C(0) and C(1)) shows the importance of the third virial coefficient.

−AϕI1/2 1 + bI1/2

(3)

⎛ ⎞3/2 ⎛1⎞ e2 ⎜ ⎟ 2πN d ⎜ ⎟ A 0 ⎝3⎠ ⎝ 2πε0DekT ⎠

(4)

Bϕ = β (0) + β (1) exp( −α1I1/2) + β (2) exp( −α2I1/2)

(5)

C ϕ = C(0) + C(1) exp( −α3I1/2)

(6)

In the preceding equations, β(0), β(1), β(2), C(0), and C(1) are the ion interaction parameters of the Archer extension of the Pitzer model. These ion interaction parameters depend on the temperature and pressure. In eq 4, Aϕ is the Debye−Hückel constant and NA, d0, e, ε0, De, and k are the Avogadro number, the density of the solvent, the electronic charge, the vacuum permittivity, the dielectric constant, and the Boltzmann constant, respectively. The values of Aϕ for isopropanol and n-butanol are (3.434 and 3.950) kg1/2·mol−1/2 at 323.15 K, respectively.18 Ionic strength of the solution is calculated by the following expression: I = (1/2) ∑ mizi 2

(7)

In the preceding equations I, mi, and zi are the ionic strength in molality unit, the molality of the ith ion, and the absolute value of ionic charge on the ith ion, respectively. The correlation of the experimental osmotic coefficient value was done by using the Archer extended Pitzer ion interaction model and has been shown in Figure 2. The solid lines present in Figure 2 are generated by using the Archer extended Pitzer ion interaction model. The standard deviation of the fit is calculated using eq 8, which depends on the number of data points, N, and the number of model parameters, s. ⎛ ∑ (ϕexp − ϕcal)2 ⎞1/2 ⎟⎟ σ(ϕ) = ⎜⎜ N−s ⎝ ⎠

(8) C



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for electrolytes in solvents. In Figure 3, the depression of vapor pressure (Δp = p* − p) vs m of the IL solutions is plotted at T

From these data, it is proved that multiple cation−cation equilibria are present in the solution. 3.2. Activity and Vapor Pressure. Activity, a, and vapor pressure, p, of the alcoholic binary IL solution have been calculated by using the following correlations (9 and 10) for the binary solution of IL with two different solvents (isopropanol and n-butanol) at T = 323.15 K, respectively. ϕ = ln(as)/νmMs

(9)

ln(as) = ln(p /p*) + (Bs − Vs*)(p − p*)/RT

(10)

where ϕ, as, and Ms are the osmotic coefficient, activity, and molar mass of the solvent, respectively. ν is the sum of the stoichiometric numbers of the anion and the cation of the ionic liquid, m is the molality of the ionic liquid, and p and p* are the vapor pressure of the solution and the vapor pressure of the pure solvent, respectively. Bs, Vs, and R are the second virial coefficient, molar volume of the pure solvent, and the universal gas constant, respectively. The second term on the right side of eq 10 measures the nonideal behavior of the solvent, and the values of Bs and Vs are shown in Table 4, at T = 323.15 K. The nonideal behavior of the solution observed is due to the ion− ion interaction.

Figure 3. Change of vapor pressure, Δp, against with molality, m, at T = 323.15 K: ▲, {[BmPr][I] + isopropanol}; ○, {[BmPr][I] + nbutanol}.

= 323.15 K for both solvents. The depression of vapor pressure is a measure of the difference between the vapor pressure of the pure alcohol and the vapor pressure of the IL alcohol solution; it is a property directly related to the solute−solvent interactions. From Figure 3 the depression of vapor pressure of the mixture of IL with isopropanol is greater than that of the IL with n-butanol. It indicates that the solute−solvent interaction between the IL and isopropanol solvent molecules is greater than the IL and n-butanol molecule; for this reason the depression of vapor pressure is greater for isopropanol than for n-butanol in the presence of ionic liquid. 3.3. Determination of the Mean Molal Activity Coefficient. The mean molal activity coefficients, γ ± , of alcoholic solution of IL have been calculated using eq 11 for binary 1:1 electrolyte solution.17

Table 4. Second Virial Coefficients, Bs, and Molar Volumes, Vs, of Isopropanol and n-Butanol and the Vapour Pressure of the Pure Solvent, p*, at the Studied Temperature, Ta

a

106 V s*

p*

(m ·mol )

(m3·mol−1)

kPa

−0.161944 −0.394135

0.7779 0.9288

23.59 4.48

T

104 Bs

solvent

K

isopropanol n-butanol

323.15 323.15

−1

3

From ref 18.

The values of activity decreased with the increasing molality of the solution (Table 5). The nonideal behavior was observed predominantly at lower concentrations with a negative slope, which is the common phenomenon of the osmotic coefficients

⎡ m1/2 ⎤ 2 ln γ± = −Aϕ⎢ + ln(1 + bm1/2)⎥ 1/2 b ⎣ 1 + bm ⎦

Table 5. Change of Activity, as, Vapor Pressures, p, of the Binary Solution of Isopropanol with [BmPr][I] and the Binary Solution of n-Butanol with [BmPr][I] Systems at Temperature, T,a and Atmospheric Pressure [BmPr][I] + isopropanol T

mb

p −1

+ m(2β 0 + A1 + A 2 ) +

⎤ ⎛ α12m ⎞ 2β1 ⎡ 1/2 ⎢ ⎟ exp( −α1m1/2)⎥ ⎜ 1 1 m − + α − 1 2 ⎥⎦ 2 ⎠ α1 m ⎢⎣ ⎝

A1 =

p −1

K

(mol·kg )

as

kPa

(mol·kg )

as

kPa

323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15

0.0529 0.1022 0.1978 0.3344 0.5053 0.7375 1.0041 1.2316 1.5529 1.8390 2.2383

0.995 0.992 0.986 0.978 0.971 0.961 0.950 0.942 0.936 0.928 0.913

23.481 23.404 23.263 23.076 22.887 22.645 22.375 22.187 22.035 21.815 21.430

0.0483 0.1040 0.2081 0.2954 0.5164 0.7700 0.9771 1.2427 1.5288 1.8715 2.2662

0.995 0.990 0.983 0.978 0.966 0.955 0.946 0.940 0.928 0.913 0.900

4.457 4.436 4.403 4.380 4.323 4.275 4.231 4.203 4.145 4.074 4.007

(11)

where

[BmPr][I] + n-butanol mc

m2 (3C 0 + A3) 2

(12)

A2 =

⎤ ⎛ α 2m ⎞ 2β 2 ⎡ ⎢1 − ⎜1 + α2m1/2 − 2 ⎟ exp( −α2m1/2)⎥ 2 ⎥⎦ 2 ⎠ α2 m ⎢⎣ ⎝ (13)

A3 =

⎛ 4C1 ⎡ m2 ⎞ ⎢6 − ⎜6 + 6α3m1/2 + 3α32m + α33m2 − α3 4 ⎟ 4 2 2 ⎠ ⎝ α3 m ⎣ ⎤ × exp(− α3 4m2)⎥ ⎦

(14)

and where ϕ is the experimental osmotic coefficient value. The value of γ ± depends on how well the Archer extension of the Pitzer model correlated the osmotic coefficient value. The parameters of the Archer extension of the Pitzer model are

Standard uncertainties u are u(m) = 0.001 mol·kg−1, u(T) = 0.01 K, and the combined expanded uncertainty Uas is Uas(C) = 0.004 (0.95 level of confidence). bSolvent = isopropanol, cSolvent = n-butanol. a

D



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obtained from the data fitting (Table 3). From Figure 4 it is observed that with increasing concentration of the solution the value of γ ± decreases.

Figure 6. Excess Gibbs free energy, GE, plot against molality, m, at T = 323.15 K and atmospheric pressure: ▲, {[BmPr][I] + isopropanol}; o, {[BmPr][I] + n-butanol}. Figure 4. Mean molar activity coefficients, γ±, against molality, m, at T = 323.15 K and atmospheric pressure: ▲, {[BmPr][I] + isopropanol}; ○, {[BmPr][I] + n-butanol}.

micellar concentration in solvent medium by using the pseudophase separation model.30,31 This model is based on the experimental value of ϕ and the molality of the solution. By using this model, we can estimate the cac and n. Here instead of saying the critical micellar concentration, we prefer to say it as critical aggregation concentration.

In Figure 5, we showed the deviation of the IL and alcohol binary mixture system from the Debye−Huckel limiting law (DHLL). It is observed from the preceding figures that with increasing concentration deviation from the DHLL law increases. The possible reason for the deviation may be due to ion association of the IL within the solution.15,29 3.4. Excess Gibbs Free Energy. The excess Gibbs free energy, GE, can be expressed in terms of the activity coefficients and osmotic coefficients by the following equation.9,14 GE = νnILRT[ln(γ±) + 1 − ϕ]

ϕ=

⎛ 1 1 ⎞ cac + ⎜1 − ⎟ ⎝ n n⎠ m

(16)

where n is the aggregation number at which aggregation starts and cac is the concentration at which the ions start to aggregate. The values of n and cac are determined by extrapolation of the liner plots (Figure 7). In this model ϕ vs (1/m) is plotted to obtain cac as well as the aggregation number of the ILs in both solvents. We found the critical aggregation concentrations of (0.30 and 0.26) mol·kg−1 for the studied IL in isopropanol and n-butanol at T = 323.15 K, respectively. The value of cac gradually decreases when the solvent medium was changed from isopropanol to n-butanol due to the greater polarity of isopropanol as compared to n-butanol. The values of n were found to be 3.0 and 3.5 for the studied IL in isopropanol and nbutanol at T = 323.15 K, respectively.

(15)

where ν and nIL are the sum of the stoichiometric numbers of the anion and the cation of the IL and number of moles of IL, respectively. By using the preceding expression, the excess Gibbs free energy value, GE was calculated. In Figure 6, we showed the variation of GE with the molality of the IL in isopropanol and n-butanol solvents, respectively. It was observed that with increasing concentration the excess Gibbs free energy value decreased. From this figure it is also observed that the excess Gibbs free energy of IL in isopropanol is greater than that in n-butanol, due to presence of a short aliphatic chain in n-butanol. 3.5. Critical Aggregation Concentration. Attwood et al. and Desnoyers et al. suggested determination of the critical

4. COMPARISON WITH THE LITERATURE The thermodynamic parameters of [BmPr][I] + isopropanol and [BmPr][I] + butanol systems were obtained in this work. The thermodynamic parameters were compared with literature having the same cation but different anions in the same solvent,

Figure 5. Activity coefficients, γ±, against the square root of molality, m1/2, at T = 323.15 K and atmospheric pressure: (a) isopropanol}; (b) ○, {[BmPr][I] + n-butanol}. The straight lines were generated by using DHLL. E

▲,

{[BmPr][I] +



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Figure 7. Variation of osmotic coefficient, ϕ, with reciprocal of molality, 1/m at T = 323.15 K: (a) {[BmPr][I] + n-butanol}.

▲,

{[BmPr][I] + isopropanol}; (b) ○,

activity coefficient, γ ± , were obtained by using experimental ϕ with the help of different mathematical expressions. We showed the variation of thermodynamic parameters with concentration and also showed the dependency of thermodynamic parameters of binary solution IL on the solvent medium. It was found that with increasing the molality the value of ϕ, γ ± , and p were decreased due to the ion association. The dependency of excess Gibbs free energy, GE, on concentration and the nature of the solvent medium has been shown in this work. The values of the critical aggregation concentration, cac, and the aggregation number, n, of the studied IL in isopropanol and n-butanol have been estimated by applying the pseudophase separation model on the experimental value of ϕ. The cac value of the studied IL in isopropanol was found to be greater as compared to the cac value in n-butanol at T = 323.15 K. By comparing with the literature, we showed that the interaction of IL with water molecule is greater as compared to the studied solvent molecules and also compared the different thermodynamic parameters of the studied systems with other binary systems reported in the literature.

where the alcoholic mixture of 1-butyl-1-methylpyrrolidinium trifluoromethanesulfonate, 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)imide, and 1-butyl-1-methylpyrrolidinium dicyanamide have been studied.7,18 In our system we obtained similar types of features that were reported for alcoholic solutions of IL having the same cation but having different anions. By using this method it is possible to analyze in a simple way the influence of structure on the thermodynamic properties such as osmotic coefficient, activity, and vapor pressure of binary mixtures. From the variation of osmotic coefficient (Table 2 and Figure 2) with concentration, it is observed that with increasing concentration the ϕ value gradually decreases. At lower concentration ϕ is greater as compared to at higher concentration. The probable reasons for such nonideal behavior may be due to ion-association interactions.7,12 It is implied that the interaction between the solvent molecules with IL molecules decreased with increasing concentration whereas, the ion−ion interaction increases. The osmotic coefficient value of 1-butyl-1-methylpyrrolidinium iodide [BmPr][I] is lower than 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)imide [BmPr][NTf2].12 The reason may be the higher solvation of NTf2 anion of [BmPr][NTf2] as compared to iodide anion of [BmPr][I]. In our previous study, we reported the thermodynamic properties of aqueous solution of 1-butyl-1-methylpyrrolidinium iodide [BmPr][I].8 By comparing with this, it is observed that the osmotic coefficient value of IL in a water system is higher compared to the osmotic coefficient value of 1-butyl-1methylpyrrolidinium iodide in alcohol systems. It is implied that there is greater interaction of IL with water as compared to alcohol. The solvation of the IL is in the following order:

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Fax: 91-612-2277383. Funding

This work was supported by a grant to D.S. from the Council of Scientific and Industrial Research (CSIR), New Delhi, Government of India, Scheme No. 01(2523)/11/EMR-II, dated Dec. 12, 2011. S.A.A. and B.M. are thankful to IIT Patna for research fellowship. A.C. is thankful to CSIR, New Delhi for a research fellowship.

water > isopropanol > n‐butanol

Notes

The authors declare no competing financial interest.

■ ■

In our study, we have found a similar trend for ϕ. Hence, we can say that the trend of osmotic coefficient value in water, isopropanol, and n-butanol is also in the same trend of solvation of [BmPr][I].

ACKNOWLEDGMENTS We are thankful to IIT Patna for the infrastructural and experimental facilities provided by IIT Patna.

5. CONCLUSIONS The experimental osmotic coefficient, ϕ, values of binary solutions of 1-butyl-1-methylpyrrolidinium iodide (IL) with isopropanol and n-butanol were obtained by using VPO technique at T = 323.15 K and atmospheric pressure. These values were correlated by using the Archer extension of the Pitzer model, and different ion interaction parameters were obtained. The different thermodynamic parameters such as activity, a, depression in vapor pressure, Δp, and mean molal

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