Mean Activity Coefficients for NaCl in the Mixtures Containing Ionic

Article pubs.acs.org/jced

Mean Activity Coefficients for NaCl in the Mixtures Containing Ionic Liquids [Emim][MeSO3] + H2O and [Emim][EtSO4] + H2O at 298.15 K Eliseo Amado-Gonzalez,*,† Miguel A. Esteso,*,‡ and Wilfred Gomez-Jaramillo† †

Department of Chemistry, University of Pamplona, (57+7) 5685303−5685304, IBEAR FJ-207 Biofuels Laboratory, Pamplona, Colombia U.D. Química Física, Universidad de Alcalá, 28871 Alcalá de Henares (Madrid), Spain

‡

S Supporting Information *

ABSTRACT: Mean activity coefficient values for NaCl in the mixture (water + ionic liquid) (IL) as solvent (IL being 1-ethyl-3-methyl-imidazolium methanesulfonate, [Emim][MeSO3] or 1-ethyl-3-methyl-imidazolium ethyl sulfate, [Emim][EtSO4]) were determined from the electrode potential measurements of the cell: Na−(ISE) | NaCl (mA), ILs (w), H2O (1 − w) | Cl−(ISE) (ISE means ion selective electrode) at total ionic strengths ranging from 0.10 to 3.20 mol·kg−1 and at 298.15 K. Different weight fractions (w) of ILs (w = 0.01, 0.05, 0.1, 0.2, 0.3 and 0.4) were used. Data were fitted to the Pitzer model and the ion interaction parameters β0, β1, and Cγ were calculated. The osmotic coefficients of the solvent (water + IL), the solvent activity, the excess Gibbs free energy for the system were calculated. The negative sign in the excess free energy for the NaCl + IL+ H2O solvent mixture at 298.15 K suggests a spontaneous aggregation of {[Emim][MeSO3] + H2O} or {[Emim][EtSO4] + H2O}.

Table 1. Chemicals, Suppliers, and Stated Puritya

1. INTRODUCTION Nowadays, the measurements of electrochemical properties of acidic ionic liquids (ILs) become useful to evaluate new electrochemical theories1,2 to predict physical3−5 and thermodynamic properties of ILs to make possible novel industrial applications of ILs.6−10 The electrochemical stability of ILs gives the possibility to play important roles in biology, pharmacy, and medicine11,12 as well as to be used as solvents in liquid extraction and separation processes13,14 among other possibilities of application. Despite the scientific and technological importance of imidazolium ILs and activity coefficients of ternary systems involving ILs plus other electrolytes, the knowledge of the activity coefficient values for ILs in mixed electrolyte solutions is very important to understand the phase behavior of these salts, and data of these systems are not found in the literature either.15,16 Furthermore, in most of the papers the IL has been considered as the solute component, and now ILs are considered as new green solvents.17 The importance of our studies with IL + H2O solvent mixtures is related to future electrochemical model development and to evaluate the behavior of these mixtures as solvents in the interactions with salts. In this study, we use the mixture of [Emim][MeSO3] + H2O and [Emim][EtSO4] + H2O as solvents at different IL weight fractions (w = 0.01, 0.02, 0.10, 0.20, 0.30, and 0.40) for the determination of the activity coefficients of NaCl obtained from emf measurements at 298.15 K.

chemical name NaCl

Sigma-Aldrich Puriss. p.a.,Reag. ACS [Emim] Solvent Innovation [MeSO3] GmbH [Emim] Solvent Innovation [EtSO4] GmbH water a

purity 99.8% >99% >99%

purification method

.

no further purification no further ≤0.03% purification no further ≤0.03% purification double distillation

Determined by Karl Fisher method.

higher than 99% were purchased from Solvent Innovation GmbH, Cologne, Germany. They were used as received without any further purification. Nevertheless, their water content was determined by using a volumetric Karl Fischer method (TitroLine 7500 KF trace from Schott instruments) as ≤0.03%. Sodium chloride (Merck, pro analysis) was just dried under vacuum at 373.15 K for 3 days before use. All the chemicals were kept in a desiccator over silica gel. Millipore water {κ = (0.7−0.9) × 10−4 S m−1} was used. Solutions were freshly prepared by direct weighing each one of the components. All solutions were prepared by weight. For each set of experiments, corresponding to a molar fraction (w), the solutions were obtained by adding successive solid-weighed NaCl to the solution previously prepared for each IL and bidistilled water (κ < 5 × 10−7 S cm−1). The solutions were stirred with a magnetic stirrer.

2. EXPERIMENTAL SECTION 2.1. Chemicals and Solutions. In Table 1 are the list of materials used in this work: 1-ethyl-3-methyl-imidazolium ethylsulfate [Emim]-[EtSO4] and 1-ethyl-3-methyl-imidazolium methanesulfonate [Emim][MeSO3] ILs with a nominal purity © 2017 American Chemical Society

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Received: September 21, 2016 Accepted: December 22, 2016 Published: January 6, 2017 752

DOI: 10.1021/acs.jced.6b00820 J. Chem. Eng. Data 2017, 62, 752−761

Journal of Chemical & Engineering Data

Article

1 × 10−5 g cm−3 and accuracy of 5 × 10−5 g cm−3 in the range of 0−95 °C of temperature and 0−10 bar of pressure. The temperature inside the U-tube is controlled by a Peltier system with a precision of 0.01° and accuracy of 0.03°. The calibration of the densimeter was checked time to time with ultrapure water. The uncertainty of these density values was better than 2 × 10−5 g cm−3.

2.2. EMF Measurements. The potentiometric measurements were performed from electrochemical cells containing two ion selective electrodes (ISEs) in accordance to the scheme described elsewhere18 Na-ISE | NaCl (m); IL (w); water (1 − w) | Cl-ISE (I) where m is the molality of NaCl in the different (IL + H2O) mixtures and w is the mass fraction of IL in these mixtures (w = 0.01, 0.05, 0.10, 0.20, 0.30, and 0.40). The ISEs used were a Na-ISE (model 8611BNWP) and a Cl-ISE (model 9617BNWP) from Thermo Scientific Co. An Orion Versa Star 40 Dual channel Benchtop meter (model VSTAR40B) was used to carry out the EMF measurements with a resolution of ±0.01 mV. A double-wall cell, enabling the circulation of thermostated water, was used to place the electrodes and the solution. The temperature inside the cell was maintained constant {at T = (298.15 ± 0.02) K} with the help of a thermostat-cryostat (Polyscience Scientific Thermostat model 1156). This temperature was monitored by using a calibrated type K thermocouple. As usual, all the measurements were performed under stirring conditions. The EMF of cell (I) can be expressed by the Nernst equation

E = E0 + 2k ln mγ±

3. RESULTS AND DISCUSSION 3.1. Mean Ionic Activity Coefficients. Experimental values of E for various NaCl molal concentrations in the different (IL + water) mixtures studied are collected in Tables 4 and 5 for the two ILs used, together with those for the mean ionic activity coefficient of NaCl, γ±(NaCl), calculated by mean of eq 1. These mean activity coefficient values are also plotted against the molal concentration of NaCl in Figures 1 and 2 for both mixtures (IL + water) studied; both systems show similar shaped curves with a pronounced decrease in the diluted zone, the definition of a minimum (around 0.5 mol kg−1), and a later tendency to slightly increase when the molal concentration increases. In both (IL + water) mixtures, this initial decreasing is more pronounced as the presence of the IL is greater. This behavior is the usual for strong electrolytes in a high dielectric constant medium and it is attributed to a predominance of the ion−ion interactions (long-range interactions). Hence, such pronounced initial decreases have to be related to the existence of interactions between the NaCl ions themselves and also to interactions of the ions from NaCl with the ions from the IL due to the increase of the solute−solvent interaction, similarly to that we found for similar multicomponent aqueous system previously studied.18 However, when comparing the behavior of both the curves for the [Emim][EtSO4] + H2O solvent mixture and the curves of the [Emim][MeSO3] + H2O solvent mixture, sensible differences are found that indicate that the role of the ethyl group on the anion of the IL, [EtSO4]−, is much more marked than the methyl group on the [MeSO3]− anion. By taking into account all the above considerations, for the (NaCl + IL + H2O) ternary system it would be possible to conclude that interactions between the ions from NaCl and the ions of the IL exist and, undoubtedly, those concerning the [EtSO4]− anion are more important. On the other hand, some authors have considered hydrophobic solvation effect of the nonpolar groups on the IL by hydrogen-bonded water structures to explain the interaction.22,23 Our results suggest that the hypothesis of water preferentially hydrogen bonding to the anion may be an explanation of the behavior of the osmotic coefficients. The experimental mean activity coefficients (γ±) were calculated by using the Pitzer model24 that for a 1:1 electrolyte is described through the equations

(1)

where k = RT/F is the theoretical Nernst slope with R, T, and F being the gas constant, the absolute temperature, and the Faraday constant, respectively. E and E0 are, respectively, the cell potential and the apparent standard cell potential, the latter dependent on both the activity of the ions in the internal reference solution and the asymmetry potential of the two ion selective electrodes used. Table 2 shows the values of NaCl (m), Table 2. Experimental E Values, Mean Activity Coefficients (γ±) of NaCl at Different Molalities (m) at T = 298.15 K and p = 101 kPaa m mol·kg−1

E mV

γ±19

m mol·kg−1

E mV

γ±19

0.0629 0.3220 0.5184 0.7846 0.9145 1.1826 1.4022 1.6188

0.06 77.02 100.54 120.26 128.10 139.30 148.38 156.2

0.8061 0.7027 0.6774 0.6608 0.6567 0.6531 0.6538 0.6568

1.8339 2.0485 2.2625 2.4768 2.6916 2.9078 3.1232

163.08 169.38 175.06 180.50 185.44 190.41 194.94

0.6617 0.6682 0.6761 0.6851 0.6954 0.7069 0.7195

a Standard uncertainties are u(m) = 0.0001 mol·kg−1, u(w) = 0.0001, u(E) = 0.5 mV, u(T) = 0.03 K, u(p) = 10 kPa.

cell potential (E0), and mean activity coefficient of NaCl (γ±) used to calculate E0 value for w = 0 (that is, in pure water) were performed, because the values for the mean activity coefficients of NaCl in pure water at 298.15 K are known in a wide concentration range.19 A Nernstian response of the electrodes is found by using a least-squares method; a value of E0 = 153.35 mV was obtained from the intercept, with the experimental slope value k = 25.63 ± 0.10 mV, which perfectly agrees with the theoretical one, with a regression coefficient (R2 = 0.9999). Table 3 shows the average molar mass (M), relative permittivity (εr), density (ρ1), Debye− Hückel parameter (Aϕ), and cell potential (E0) for (IL + H2O) mixtures at 298.15 K. 2.3. Density Measurements. Density measurements were carried out by using an oscillating U-tube densimeter (Rudolph Research Analytical, model DDM 2911) with a resolution of

ln γ± = f γ + mBγ + m2C γ

(2)

where ⎡ ⎤ ⎛2⎞ I1/2 1/2 ⎜ ⎟ln(1 + bI ⎥ f γ = −Aϕ⎢ + ) ⎝b⎠ ⎣ (1 + bI1/2) ⎦

(3)

⎧ [1 − exp(− αI1/2)(1 + αI1/2 − 1/2αI1/2)] ⎫ ⎬ Bγ = 2β0 + 2β1⎨ (αI1/2) ⎩ ⎭ (4) ⎪

⎪

⎪

⎪

In eqs 2−5, β0, β1, and Cγ are the Pitzer parameters whose values depend on both the interactions between the solutes and the 753



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Table 3. Values of the Average Molar Mass (M), Relative Permittivity (εr), density (ρ1), Debye−Hückel Parameter (Aϕ), and Cell Potential (E0) for (IL + H2O) Mixtures at 298.15 K and p = 101 kPaa IL w [Emim][EtSO4] 0.01 0.05 0.10 0.20 0.30 0.40 [Emim][MeSO3] 0.01 0.05 0.10 0.20 0.30 0.40 a

M g·mol−1

εr20

ρ1 g·cm−3

ρ g·cm−3 21

Aϕ kg1/2·mol1/2

E0 mV

18.1852 18.8897 19.8703 22.0967 24.9343 28.6013

77.86 76.47 74.62 70.93 67.06 63.01

0.99874 1.00544 1.01460 1.03417 1.05673 1.08198

0.9987 1.0055 1.0146 1.0342 1.0568 1.0819

0.3957 0.4079 0.4251 0.4631 0.5093 0.5658

138.21 143.50 150.86 167.57 188.86 216.39

18.1827 18.8789 19.8598 22.0811 24.7918 28.3873

77.73 75.65 72.95 67.56 62.07 56.17

0.99729 0.99820 0.99948 1.00239 1.00593 1.01063

0.3965 0.4131 0.4365 0.4905 0.5580 0.6497

153.31 130.52 138.24 152.39 166.97 181.76

Standard uncertainties are u(w) = 0.0001, u(ρ) = 5 × 10−5 g·cm−3, u(T) = 0.03 K, u(p) = 10 kPa, u(E) = 0.5 mV.

temperature and are determined by fitting the experimental data; I is the ionic strength; α and b are fixed parameters (assumed values are 2.0 and 1.2 kg1/2 mol−1/2, respectively), and Aϕ is the Debye−Hückel constant for the osmotic coefficient given by ⎛ e 2 ⎞1.5 1 Aϕ = 2πNoρ1 ⎜ ⎟ 3 ⎝ 4πεr kT ⎠

dissociated, it is considered as the number of ions released by dissociation, νi = ν+ + ν−), and mi is the molality of the solute i. In the Pitzer model,27 this osmotic coefficient is described by the set of equations ϕ − 1 = f ϕ + mBϕ +

(5)

fϕ =

N

−AϕI1/2 (1 + bI1/2)

Bϕ = β0 + β1 exp( −αI1/2)

(8) (9)

In Table 3, the values of Aϕ are shown for each weight fraction (w). The values of α and b are fixed parameters (assumed values are 2.0 and 1.2 kg1/2 mol−1/2, respectively). Therefore, the calculated osmotic coefficients values are collected in Tables 4 and 5 for both mixtures (IL + water) studied. In Figures 5 and 6, the osmotic coefficients are plotted against the NaCl molal concentration for [Emim][EtSO4] + H2O and [Emim][MeSO3] + H2O solvent mixtures, respectively at the different IL mass fractions studied. The dependence of the osmotic coefficient with the electrolyte concentration responds to the typical pattern for strong electrolytes. That is, an initial decrease (negative slope), a minimum (around m = 0.3 mol kg−1), which is more pronounced as the IL concentration is higher, especially in the case of the [Emim][EtSO4] + H2O solvent mixture, and afterward an increase (positive slope) at high NaCl concentrations. The results suggest that the size of the hydrocarbon chain of the anion in the IL plays an important role in the interactions taking place in these multicomponent systems as previously proposed.28 Thus, it has to be assumed that a different behavior in relation to the solvation of NaCl in [Emim][EtSO4] + H2O solvent mixture probably is caused by an interaction between water molecules around the anion.29−31 By using eq 6 values for the solvent activity (asolvent = as), the NaCl + [Emim][EtSO4] + H2O and NaCl + [Emim][MeSO3] + H2O ternary systems were calculated. The solvent activity values for [Emim][EtSO4] + H2O and [Emim][MeSO3] + H2O solvent mixtures are collected in Tables 7 and 8 and plotted against the NaCl molality in Figures 7 and 8. By taking into account that the solvents used were IL + water mixtures at different IL weight fractions (w = 0.01, 0.05, 0.10, 0.20, 0.30, and 0.40),

−ln asolvent νim

Msolvent ·∑i = 1 1000i

(7)

with

with ρ1 and εr being the density and the static dielectric constant (relative permittivity) of the mixed solvent. N0 and k are Avogadro’s number and Boltzmann’s constant, respectively, and the other symbols have their usual meaning. The density (ρ1) values of the mixed solvents for the IL mass fractions used in this work were measured. Those for the dielectric constant (εr) were estimated from literature data24,25 by applying the mixing rule suggested by Oster26 to the polarization per unit volume of the pure components of the mixture. Table 3 shows the values of the average molar mass (M), relative permittivity (εr), density (ρ1), Debye−Hückel parameter (Aϕ), and cell potential (E0) for (IL + H2O) mixtures at 298.15 K together with the density values from literature.25 By fitting the measured experimental EMF values to eqs 2, 3, 4, 5 and 9, the Pitzer parameters, β0, β1, and Cγ, were obtained for the different IL mass fractions in both IL + H2O mixtures studied. Their values are summarized in Table 6, as well as the standard deviations values found for these fittings. Figures 3 and 4 show the tendency of the Pitzer parameters for both mixtures (IL + water) studied. The results suggest that the effect of the anion (EtSO4−1) mainly on the behavior of β0, and β1 is higher than the effect of MeSO3−1 on Pitzer parameters. 3.2. Osmotic Coefficients and Activity of the Solvent. Another way to study these liquid systems consists in analyzing the behavior of the solvent mixture. However, because in the dilute region the values for the solvent activity coefficient are very close to 1, it is usual to use the osmotic coefficient of the solution instead. This osmotic coefficient is given by ϕ=

2 2 γ mC 3

(6)

with asolvent and Msolvent being, respectively, the activity and the molar mass of the solvent, νi is the stoichiometric coefficient of the solute i (for strong electrolytes, which are assumed to be fully 754



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Table 4. Experimental E Values, Mean Activity Coefficients (γ±) of NaCl, and Osmotic Coefficient (ϕ) of the Solution at Different Molalities (m) in {[Emim][EtSO4 ] + H2O} Mixtures with Variable Mass Fraction of IL (w) at T = 298.15 K and p = 101 kPaa m/mol·kg−1

E/mV

w = 0.01 [Emim][EtSO4 ] 0.1723 38.25 0.3464 72.91 0.5208 92.93 0.6936 106.86 0.8641 117.79 1.0388 128.65 1.3942 143.23 1.7324 154.72 w = 0.05 [Emim][EtSO4 ] 0.1094 18.00 0.3303 71.23 0.5483 99.01 0.7689 115.88 0.9826 128.11 1.1967 138.45 1.4131 147.19 1.6292 154.91 w = 0.10 [Emim][EtSO4 ] 0.1092 29.09 0.3269 78.21 0.5410 101.81 0.7559 117.98 0.9665 130.34 1.1791 143.30 1.3940 151.60 1.6194 159.52 w = 0.20 [Emim][EtSO4 ] 0.2092 68.20 0.4310 101.73 0.6412 121.23 0.8567 140.22 1.0649 151.20 1.2793 160.79 1.4973 169.33 1.7156 177.01 w = 0.30 [Emim][EtSO4 ] 0.1056 44.03 0.3225 96.90 0.5374 122.68 0.7505 140.53 0.9654 153.89 1.1828 164.90 1.4012 174.62 1.6167 182.69 w = 0.40 [Emim][EtSO4 ] 0.1084 76.85 0.3254 113.70 0.4324 127.80 0.5466 139.58 0.6532 148.91 0.7609 156.75 0.8674 163.54 0.9822 170.41 a

γ±(NaCl)

ϕ

m/mol·kg−1

E/mV

γ±(NaCl)

ϕ

0.8128 0.8009 0.7968 0.7946 0.7933 0.7927 0.7937 0.7979

0.9633 0.9712 0.9760 0.9788 0.9804 0.9816 0.9839 0.9875

2.0713 2.4159 2.7629 3.1005 3.4407 3.7849 4.1206

164.47 173.30 181.25 188.39 195.14 201.67 207.55

0.8058 0.8178 0.8342 0.8545 0.8797 0.9101 0.9450

0.9931 1.0012 1.0119 1.0247 1.0400 1.0579 1.0775

0.7988 0.7603 0.7511 0.7494 0.7513 0.7556 0.7618 0.7697

0.9482 0.9542 0.9629 0.9713 0.9794 0.9876 0.9963 1.0055

1.8427 2.0576 2.2748 2.7077 3.1326 3.5431 3.9661

161.86 168.24 174.44 185.02 194.61 203.21 211.46

0.7791 0.7899 0.8023 0.8313 0.8652 0.9034 0.9483

1.0152 1.0255 1.0365 1.0603 1.0859 1.1126 1.1419

0.7850 0.7382 0.7235 0.7173 0.7155 0.7164 0.7197 0.7251

0.9424 0.9442 0.9503 0.9567 0.9632 0.9702 0.9778 0.9864

1.8316 2.0483 2.2678 2.4769 2.9013 3.3338 3.7446

166.01 172.20 178.15 183.79 193.28 202.48 210.89

0.7320 0.7406 0.7510 0.7624 0.7898 0.8238 0.8620

0.9953 1.0051 1.0157 1.0266 1.0506 1.0776 1.1056

0.6950 0.6608 0.6533 0.6558 0.6634 0.6744 0.6877 0.7023

0.9154 0.9252 0.9425 0.9631 0.9841 1.0061 1.0285 1.0507

1.9331 2.1459 2.3605 2.5680 2.7798 2.9844 3.3431

183.96 190.16 196.32 201.59 206.80 211.70 219.84

0.7177 0.7332 0.7489 0.7638 0.7788 0.7926 0.8152

1.0724 1.0930 1.1132 1.1321 1.1506 1.1678 1.1960

0.6428 0.5291 0.4998 0.4963 0.5045 0.5191 0.5369 0.5556

0.8662 0.8429 0.8680 0.9086 0.9557 1.0052 1.0542 1.1007

1.8266 2.0325 2.2455 2.4491 2.6696 2.7756 2.8735

189.97 196.77 203.17 209.01 215.04 217.78 220.11

0.5735 0.5895 0.6038 0.6144 0.6220 0.6240 0.6249

1.1433 1.1821 1.2186 1.2500 1.2798 1.2926 1.3035

0.5764 0.4445 0.4217 0.4089 0.4035 0.4023 0.4039 0.4079

0.8269 0.7900 0.8004 0.8205 0.8445 0.8718 0.9006 0.9328

1.0886 1.1945 1.3008 1.5387 1.7860 2.0258 2.4710

175.91 181.11 185.99 195.72 204.88 211.98 224.95

0.4129 0.4189 0.4253 0.4402 0.4541 0.4641 0.4686

0.9630 0.9929 1.0226 1.0861 1.1464 1.1981 1.2739

Standard uncertainties are u(m) = 0.0001 mol·kg−1, u(w) = 0.0001, u(T) = 0.01 K, u(p) = 10 kPa, u(E) = 0.5 mV.

i component and with Mwater = 18.0151 g/mol, M[Emim][EtSO4] = 236.29 g/mol, and M[Emim][MeSO3] = 206.26 g/mol. These average molar masses are summarized in Table 5.

it was previously necessary to determine their average molar masses (Msolvent = Ms). They were calculated from ref 32: Ms = XwaterMwater + XIL MIL, Xi being the mole fraction of the 755



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Table 5. Experimental E Values, Mean Activity Coefficients (γ±) of NaCl, and Osmotic Coefficient (ϕ) of the Solution at Different Molalities (m) in {[Emim][MeSO3] + H2O} Mixtures with Variable Mass Fraction of IL (w) at T = 298.15 K and p = 101 kPaa m mol·kg−1

E mV

w = 0.01 [Emim][MeSO3] 0.0955 −6.21 0.3159 41.87 0.5316 70.49 0.7433 87.40 0.9631 99.32 1.1826 111.70 1.3976 120.36 1.6152 127.10 w = 0.05 [Emim][MeSO3] 0.1081 5.33 0.3223 57.86 0.5413 81.86 0.7610 102.05 0.9765 113.23 1.2002 123.20 1.4165 131.98 1.6384 139.49 w = 0.10 [Emim][MeSO3] 0.1049 15.05 0.3264 70.03 0.5398 95.23 0.7573 115.30 0.9747 127.02 1.1904 136.74 1.4063 146.41 1.6180 155.00 w = 0.20 [Emim][MeSO3] 0.1026 31.12 0.3199 82.13 0.5446 110.31 0.7645 128.86 0.9740 141.10 1.1871 150.26 1.4034 158.53 1.6187 167.34 w = 0.30 [Emim][MeSO3] 0.1165 50.01 0.2212 78.18 0.4294 107.52 0.6467 127.36 0.8621 147.27 1.0786 158.20 1.2939 167.42 1.5099 176.35 w = 0.40 [Emim][MeSO3] 0.1040 58.62 0.2065 88.41 0.3239 110.84 0.4295 121.91 0.5377 133.12 0.6454 142.09 0.8669 157.82 1.0875 170.13 a

γ± (NaCl)

ϕ

m mol·kg−1

E mV

γ± (NaCl)

ϕ

0.7166 0.6092 0.5837 0.5828 0.5936 0.6107 0.6305 0.6517

0.8980 0.8777 0.8996 0.9341 0.9752 1.0179 1.0593 1.0997

1.8357 2.0492 2.2694 2.4854 2.7031 2.9195 3.1343

135.59 142.06 147.28 151.38 157.51 161.80 167.62

0.6729 0.6920 0.7095 0.7235 0.7339 0.7401 0.7417

1.1383 1.1730 1.2058 1.2347 1.2605 1.2828 1.3014

0.7017 0.6056 0.5804 0.5788 0.5879 0.6029 0.6199 0.6382

0.8940 0.8783 0.9000 0.9342 0.9723 1.0131 1.0519 1.0900

1.8531 2.0693 2.2840 2.4992 2.7164 2.9319 3.1510

146.83 153.29 159.03 162.48 167.66 173.59 177.17

0.6551 0.6706 0.6835 0.6933 0.6995 0.7016 0.6992

1.1245 1.1565 1.1854 1.2113 1.2340 1.2532 1.2692

0.6991 0.6004 0.5760 0.5737 0.5819 0.5955 0.6117 0.6285

0.8940 0.8783 0.8990 0.9316 0.9687 1.0068 1.0444 1.0799

1.8515 2.0683 2.2835 2.4983 2.7079 2.9189 3.1336

161.51 167.28 173.00 176.86 181.00 186.00 189.76

0.6467 0.6622 0.6755 0.6861 0.6933 0.6972 0.6972

1.1168 1.1485 1.1772 1.2030 1.2253 1.2449 1.2616

0.6889 0.5934 0.5691 0.5675 0.5751 0.5875 0.6025 0.6180

0.8928 0.8799 0.9025 0.9352 0.9698 1.0059 1.0419 1.0762

1.8299 2.0509 2.2627 2.4911 2.7100 2.9267 3.1449

173.24 179.31 184.73 190.29 195.26 200.11 204.52

0.6329 0.6471 0.6587 0.6683 0.6742 0.6766 0.6751

1.1079 1.1387 1.1656 1.1916 1.2134 1.2320 1.2476

0.6612 0.6068 0.5664 0.5570 0.5609 0.5714 0.5850 0.6002

0.8868 0.8779 0.8909 0.9195 0.9531 0.9886 1.0238 1.0580

1.7244 1.9388 2.1487 2.3598 2.5719 2.7805 2.9947

185.01 191.59 197.82 204.81 209.65 213.49 218.22

0.6155 0.6300 0.6429 0.6540 0.6627 0.6686 0.6715

1.0905 1.1209 1.1485 1.1739 1.1970 1.2173 1.2354

0.6433 0.5839 0.5526 0.5384 0.5312 0.5286 0.5318 0.5416

0.8808 0.8705 0.8746 0.8845 0.8978 0.9129 0.9470 0.9827

1.3077 1.5256 1.7403 1.9575 2.1720 2.3815 2.5919

179.68 191.28 198.80 205.82 212.25 218.11 223.30

0.5550 0.5700 0.5855 0.6011 0.6158 0.6290 0.6406

1.0184 1.0529 1.0857 1.1172 1.1466 1.1734 1.1984

Standard uncertainties are u(m) = 0.0001 mol·kg−1, u(w) = 0.0001, u(E) = 0.01 mV, u(T)= 0.01 K, u(p) = 10 kPa, u(E) = 0.5 mV.

3.3. Excess Gibbs Free Energy. In Tables 7 and 8, the excess Gibbs free energy (GE) is calculated for the NaCl + [Emim][EtSO4] + H2O ternary system and for

Figure 6 shows that solvent activity for the system [Emim][EtSO4] + H2O solvent mixture is reduced by increasing molalities of NaCl in the ternary system. 756



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In eq 10, R is the gas constant, T is the absolute temperature, mNaCl is the molality of NaCl, and ϕ is osmotic coefficient of the solvent. On the basis of the collecting networks of ionic liquid, anions, cations, and water, the values of GE are produced by the

the NaCl + [Emim][MeSO3] + H2O ternary system from eq 1032 GE = RT[vNaclmNacl (1 − ϕ + ln γ±NaCl)]

(10)

Figure 1. Activity coefficients of NaCl in [Emim][EtSO4] + H2O solvent mixture at 298.15 K.

Figure 3. Pitzer parameters of NaCl in [Emim][EtSO4] + H2O solvent mixture at 298.15 K.

Figure 4. Pitzer parameters of NaCl in [Emim][MeSO3] + H2O solvent mixture at 298.15 K.

Figure 2. Activity coefficients of NaCl in [Emim][MeSO3] or + H2O solvent mixture at 298.15 K.

Table 6. Pitzer Parameters for the (NaCl + IL + H2O) Ternary System at 298.15 K and p = 101 kPaa w

a

β0 kg·mol−1

0.00 0.01 0.05 0.10 0.20 0.30 0.40

0.07722 0.03471 0.07597 0.06963 0.17398 0.37479 0.94623

0.00 0.01 0.05 0.10 0.20 0.30 0.40

0.07722 0.31100 0.30082 0.29014 0.28791 0.28791 0.28641

β1 kg·mol−1 [Emim][EtSO4] 0.25183 0.88905 0.64600 0.62636 0.19534 −0.97519 −1.64188 [Emim][MeSO3] 0.25183 −0.81485 −0.72655 −0.62673 −0.42673 −0.19004 0.04380

Cγ kg2·mol−2

σ

0.00106 0.00928 0.00311 0.00466 −0.01836 −0.07116 −0.11309

0.00064 [32] 0.00320 0.00179 0.01029 0.00244 0.01628 0.00977

0.00106 −0.05721 −0.05650 −0.05330 −0.05280 −0.04936 −0.04280

0.00064 [32] 0.00354 0.01280 0.00084 0.00453 0.00296 0.00110

Standard uncertainties are u(w) = 0.0001, u(T) = 0.01 K, u(p) = 10 kPa. 757



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Table 7. Solvent Activity and Excess Gibbs Free Energies GE at T = 298.15 K in the {[Emim][EtSO4] + H2O} mixturea mα mol·kg−1

αs

GE kJ mol−1

w = 0.01 [Emim][EtSO4] 0.1723 0.9970 −0.0821 0.3464 0.9940 −0.1950 0.5208 0.9910 −0.3168 0.6936 0.9879 −0.4414 0.8641 0.9849 −0.5661 1.0388 0.9818 −0.6945 1.3942 0.9755 −0.9531 1.7324 0.9693 −1.1914 w = 0.05 [Emim][EtSO4] 0.1094 0.9980 −0.0446 0.3303 0.9940 −0.1722 0.5483 0.9900 −0.3068 0.7689 0.9859 −0.4438 0.9826 0.9819 −0.5746 1.1967 0.9778 −0.7020 1.4131 0.9737 −0.8255 1.6292 0.9696 −0.9424 w = 0.10 [Emim][EtSO4] 0.1215 0.9980 −0.0597 0.3635 0.9942 −0.2619 0.6015 0.9903 −0.4993 0.8405 0.9862 −0.7540 1.0747 0.9821 −1.0107 1.3111 0.9779 −1.2721 1.5501 0.9734 −1.5360 1.8007 0.9685 −1.8107 w = 0.20 [Emim][EtSO4] 0.2092 0.9957 −0.1234 0.4310 0.9909 −0.2950 0.6412 0.9864 −0.4637 0.8567 0.9815 −0.6356 1.0649 0.9768 −0.7980 1.2793 0.9718 −0.9597 1.4973 0.9666 −1.1173 1.7156 0.9613 −1.2676 w = 0.30 [Emim][EtSO4] 0.1056 0.9976 −0.0579 0.3225 0.9925 −0.2380 0.5374 0.9873 −0.4352 0.7505 0.9821 −0.6337 0.9654 0.9766 −0.8311 1.1828 0.9709 −1.0248 1.4012 0.9650 −1.2110 1.6167 0.9591 −1.3853 w = 0.40 [Emim][EtSO4] 0.1084 0.9975 −0.1046 0.3254 0.9928 −0.5034 0.4324 0.9903 −0.7406 0.5466 0.9874 −1.0093 0.6532 0.9845 −1.2701 0.7609 0.9813 −1.5403 0.8674 0.9780 −1.8128 0.9822 0.9741 −2.1108

Figure 5. Osmotic coefficients of NaCl in [Emim][EtSO4] + H2O solvent mixture at 298.15 K.

Figure 6. Osmotic coefficients of NaCl in [Emim][MeSO3] + H2O solvent mixture at 298.15 K.

electrostatic interaction.33 In Figure 8, the values of GE are negative and consequently controlled by increase of ILs content in the solvent: mixture. Even though the analysis of alkyl branching cations has been proposed as the key principle in understanding the properties of ILs,34−42 the effect of anion of ILs on transport properties has been not found. However, from our results a systematic effect of the anion on the activity coefficients of NaCl could be found. The effect of anion on thermodynamic properties has been observed on conductivity of imidazoliumbased ILs.43−45 Figure 9 shows that a longer chain of the anion has a stronger effect over the GE at w = 0.30 and 0.40 for [Emim][EtSO4]. The overall uncertainties in the standard free energies were estimated to be less than 2.0%. The presence of the negative sign in the excess free energy for the NaCl + IL+ H2O solvent mixture at 298.15 K suggests a spontaneous aggregation of IL in water solution. This trend may be interpreted according by ion association between water and cation and anion of the ILs, as the alkyl chain length increases, the solvation of the ILs by water decreases, resulting in stronger association of the ions.

m mol·kg−1

αs

GE kJ mol−1

2.0713 2.4159 2.7629 3.1005 3.4407 3.7849 4.1206

0.9630 0.9564 0.9496 0.9428 0.9357 0.9282 0.9207

−1.4175 −1.6301 −1.8226 −1.9853 −2.1215 −2.2274 −2.2969

1.8427 2.0576 2.2748 2.7077 3.1326 3.5431 3.9661

0.9654 0.9611 0.9567 0.9477 0.9385 0.9291 0.9190

−1.0501 −1.1494 −1.2392 −1.3812 −1.4645 −1.4834 −1.4308

2.0366 2.2776 2.5217 2.7542 3.2260 3.7070 4.1637

0.9638 0.9589 0.9538 0.9488 0.9384 0.9275 0.9169

−2.0665 −2.3246 −2.5829 −2.8261 −3.3135 −3.8064 −4.2771

1.9331 2.1459 2.3605 2.5680 2.7798 2.9844 3.3431

0.9559 0.9505 0.9449 0.9394 0.9337 0.9281 0.9181

−1.4093 −1.5398 −1.6630 −1.7740 −1.8791 −1.9725 −2.1174

1.8266 2.0325 2.2455 2.4491 2.6696 2.7756 2.8735

0.9532 0.9472 0.9408 0.9346 0.9277 0.9243 0.9212

−1.5448 −1.6909 −1.8303 −1.9520 −2.0706 −2.1226 −2.1677

1.0886 1.1945 1.3008 1.5387 1.7860 2.0258 2.4710

0.9704 0.9665 0.9624 0.9529 0.9424 0.9320 0.9126

−2.3914 −2.6741 −2.9622 −3.6260 −4.3526 −5.1063 −6.6873

Standard uncertainties for u(m) = 0.0001 mol·kg−1 and u(w) = 0.0001.

a

4. CONCLUSION The analysis of anion on the thermodynamics properties is an important field of study. The thermodynamic investigation of the

NaCl + ILs + H2O ternary system was studied by electrode potential method using Na-ISE and Cl-ISE at 298.15 K. The activity coefficients of NaCl, the osmotic coefficients Φ, 758



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Table 8. Solvent Activity and Excess Gibbs Free Energies GE at T = 298.15 K in the {[Emim][MeSO3] + H2O} Mixturea mα mol·kg−1

αs

GE kJ mol−1

w = 0.01 [Emim][MeSO3] 0.0955 0.9984 −0.0548 0.3159 0.9950 −0.2924 0.5316 0.9913 −0.5771 0.7433 0.9875 −0.8733 0.9631 0.9831 −1.1860 1.1826 0.9784 −1.4982 1.3976 0.9734 −1.8036 1.6152 0.9682 −2.1137 w = 0.05 [Emim][MeSO3] 0.1081 0.9982 −0.0665 0.3223 0.9947 −0.3033 0.5413 0.9908 −0.5957 0.7610 0.9867 −0.9072 0.9765 0.9822 −1.2189 1.2002 0.9773 −1.5444 1.4165 0.9723 −1.8610 1.6384 0.9668 −2.1895 w = 0.10 [Emim][MeSO3] 0.1049 0.9981 −0.0655 0.3264 0.9943 −0.3143 0.5398 0.9904 −0.6030 0.7573 0.9861 −0.9147 0.9747 0.9814 −1.2327 1.1904 0.9765 −1.5497 1.4063 0.9713 −1.8682 1.6180 0.9659 −2.1831 w = 0.20 [Emim][MeSO3] 0.1026 0.9980 −0.0675 0.3199 0.9938 −0.3186 0.5446 0.9892 −0.6293 0.7645 0.9843 −0.9508 0.9740 0.9794 −1.2628 1.1871 0.9740 −1.5825 1.4034 0.9682 −1.9085 1.6187 0.9623 −2.2365 w = 0.30 [Emim][MeSO3] 0.1165 0.9974 −0.0868 0.2212 0.9952 −0.2070 0.4294 0.9906 −0.4889 0.6467 0.9854 −0.8090 0.8621 0.9798 −1.1353 1.0786 0.9739 −1.4659 1.2939 0.9677 −1.7956 1.5099 0.9612 −2.1279 w = 0.40 [Emim][MeSO3] 0.1040 0.9974 −0.0830 0.2065 0.9949 −0.2092 0.3239 0.9920 −0.3755 0.4295 0.9893 −0.5362 0.5377 0.9864 −0.7070 0.6454 0.9834 −0.8806 0.8669 0.9770 −1.2434 1.0875 0.9701 −1.6065

m mol·kg−1

αs

GE kJ mol−1

1.8357 2.0492 2.2694 2.4854 2.7031 2.9195 3.1343

0.9627 0.9572 0.9515 0.9457 0.9399 0.9342 0.9285

−2.4320 −2.7486 −3.0882 −3.4398 −3.8182 −4.2243 −4.6630

1.8531 2.0693 2.2840 2.4992 2.7164 2.9319 3.1510

0.9614 0.9558 0.9502 0.9445 0.9387 0.9330 0.9273

−2.5144 −2.8527 −3.2046 −3.5782 −3.9824 −4.4162 −4.8967

1.8515 2.0683 2.2835 2.4983 2.7079 2.9189 3.1336

0.9598 0.9539 0.9480 0.9421 0.9362 0.9304 0.9245

−2.5366 −2.8745 −3.2237 −3.5905 −3.9710 −4.3817 −4.8334

1.8299 2.0509 2.2627 2.4911 2.7100 2.9267 3.1449

0.9562 0.9497 0.9434 0.9366 0.9300 0.9235 0.9170

−2.5645 −2.9179 −3.2707 −3.6717 −4.0817 −4.5176 −4.9925

1.7244 1.9388 2.1487 2.3598 2.5719 2.7805 2.9947

0.9545 0.9475 0.9407 0.9336 0.9265 0.9195 0.9124

−2.4615 −2.8012 −3.1435 −3.5014 −3.8789 −4.2720 −4.7030

1.3077 1.5256 1.7403 1.9575 2.1720 2.3815 2.5919

0.9629 0.9554 0.9478 0.9398 0.9317 0.9237 0.9156

−1.9684 −2.3259 −2.6786 −3.0385 −3.3998 −3.7613 −4.1363

Figure 7. The as for the system [Emim][EtSO4] + H2O solvent mixture at 298.15 K.

Figure 8. The as for the system [Emim][MeSO3] + H2O solvent mixture at 298.15 K.

a Standard uncertainties for u(m) = 0.0001 mol·kg−1 and u(w) = 0.0001.

Figure 9. Excess Gibbs free energy ΔGE for the NaCl + IL+ H2O solvent mixture at 298.15 K. IL 1, [Emim][MeSO3]; and IL 2, [Emim][EtSO4].

solvent activity as, and the excess Gibbs free energy GE of the solvent mixtures, [Emim][MeSO3]+ H2O or [Emim][EtSO4] + H2O were calculated using the Pitzer model. The behavior of the

osmotic coefficient and the activity of the [Emim][EtSO4] + H2O may be explained by the effect of the alkyl chain length of the anion with water. We may suppose that the behavior of 759



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(7) Taige, M.; Hilbert, D.; Schubert, T. Mixtures of ionic liquids as possible electrolytes for Lithium ion batteries. Z. Phys. Chem. 2012, 226, 129−139. (8) Lewandowski, A.; Swiderska-Mocek, A. Ionic Liquids as Electrolytes for Li-ion Batteries -. An Overview of Electrochemical Studies. J. Power Sources 2009, 194, 601−609. (9) Wei, D.; Ivaska, A. Applications of Ionic Liquids in Electrochemical Sensors. Anal. Chim. Acta 2008, 607, 126−135. (10) Safa, M.; Chamaani, A.; Chawla, N.; El-Zahab, B. Polymeric Ionic Liquid Gel Electrolyte for Room Temperature Lithium Battery Applications. Electrochim. Acta 2016, 213, 587−593. (11) Mikkola, S. K.; Robciuc, A.; Lokajová, J.; Holding, A. J.; Lämmerhofer, M.; Kilpeläinen, I.; Holopainen, J. M.; King, A. W. T.; Wiedmer, S. K. Impact of Amphiphilic Biomass-Dissolving Ionic Liquids on Biological Cells and Liposomes. Environ. Sci. Technol. 2015, 49, 1870−1878. (12) Marrucho, I. M.; Branco, L. C.; Rebelo, L. Ionic Liquids in Pharmaceutical Applications. Annu. Rev. Chem. Biomol. Eng. 2014, 5, 527−546. (13) Requejo, P. F.; Calvar, N.; Domínguez, A.; Gomez, E. Application of the ionic liquid tributylmethyl ammonium bis(trifluoromethylsulfonyl)imide as solvent for the extraction of benzene from octane and decane at T = 298.15 K and atmospheric pressure. Fluid Phase Equilib. 2016, 417, 137−143. (14) Kubiczek, A.; Kamiński, W.; Górak, A. Modeling of single- and multi-stage extraction in the system of water, acetone, butanol, ethanol and ionic liquid. Fluid Phase Equilib. 2016, 425, 365−373. (15) Marsh, K. N.; Boxall, J. A.; Lichtenthaler, R. Room temperature ionic liquids and their mixturesa review. Fluid Phase Equilib. 2004, 219, 93−98. (16) Niedermeyer, H.; Hallett, J.; Villar-Garcia, I.; Hunt, P.; Welton, T. Mixtures of ionic liquids. Chem. Soc. Rev. 2012, 41, 7780−7802. (17) Ashassi-Sorkhabi, H.; Kazempour, A. Thermodynamic Study of Aqueous Solutions of 1-Butyl-3 methylimidazolium Tetrafluoroborate Ionic Liquid Using Potentiometric Measurements at Different Temperatures. J. Chem. Eng. Data 2016, 61, 3542−3547. (18) Hernández-Luis, F.; Amado-González, E.; Esteso, M. Activity coefficients of NaCl in trehalose-water and maltose-wáter mixtures at 298.15 K. Carbohydr. Res. 2003, 338, 1415−1424. (19) Archer, D. Thermodynamic properties of the NaCl + H2O System ll. Thermodynamic properties of NaCl (aq), NaCl.2 H2O(cr) and Phase equilibria. J. Phys. Chem. Ref. Data 1992, 21, 793−829. (20) Huang, M.; Jiang, Y.; Sasisanker, P.; Driver, G.; Weingärtner, H. Static Relative Dielectric Permittivities of Ionic Liquids at 25 °C. J. Chem. Eng. Data 2011, 56, 1494−1499. (21) Yang, J.; Lu, X.; Gui, J.; Xu, W.; Li, H. Volumetric properties of room temperature ionic liquid 2 The concentrated aqueous solutions of {1-methyl-3-ethylimidazolium ethyl sulfate + water} in a temperature range of 278.2 to 338.2 K. J. Chem. Thermodyn. 2005, 37, 1250−1255. (22) Head-Gordon, T. Is water structure around hydrophobic groups clathrate-like. Proc. Natl. Acad. Sci. U. S. A. 1995, 92, 8308−8312. (23) Stark, A.; Zidell, W.; Hoffmann, M. Is the ionic liquid 1-ethyl-3methylimidazolium methanesulfonate [Emim][MeSO3]capable of rigidly binding water? J. Mol. Liq. 2011, 160, 166−179. (24) Pitzer, K. S.; Simonson, J. M. Thermodynamics of Multicomponent, Miscible, Ionic Systems: Theory and Equations. J. Phys. Chem. 1986, 90, 3005−3009. (25) Wang, P.; Anderko, A. Computation of dielectric constants of solvent mixtures and electrolyte solutions. Fluid Phase Equilib. 2001, 186, 103−122. (26) Oster, G. The dielectric properties of mixture liquids. J. Am. Chem. Soc. 1946, 68, 2036−2040. (27) Pitzer, K.; Mayorga, G. Thermodynamics of Electrolytes. I I. Activity and Osmotic Coefficients for Strong Electrolytes with One or Both Ions Univalent. J. Phys. Chem. 1973, 77, 2300−2308. (28) Tokuda, H.; Hayamizu, K.; Ishii, K.; Susan, Md.; Watanabe, M. Physicochemical Properties and Structures of Room Temperature Ionic Liquids. 1. Variation of Anionic Species. J. Phys. Chem. B 2004, 108, 16593−16600.

activity coefficients, osmotic coefficients, activity, and excess Gibbs free energy may be explained by a highly structured ionic liquid framework. The results showed that the Pitzer model can be used to describe NaCl + ([Emim][MeSO3] or [Emim][EtSO4]) + H2O ternary system satisfactorily. The results show that the Pitzer model could correlate the experimental results and provide basic thermodynamic reference data for further research applications.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.6b00820. Figure of electrodes calibration with aqueous solution of NaCl at 298.15 K (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*(E.A.-G.) E-mail: [email protected]. *(M.A.E.) E-mail: [email protected]. ORCID

Eliseo Amado-Gonzalez: 0000-0003-4523-1323 Funding

Supported by Grant PR1300-017 (GA 150-BP-II-2013-2.1.22.1) of the Pamplona University. Notes

The authors declare no competing financial interest.

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LIST OF SYMBOLS E: cell potential GE: excess Gibbs free energy m: molality M: molecular weight Pm: polarization w: molar fraction xi: mole fraction of component

Greek Letters:

αs: solvent activity ρ: density ϕ: osmotic coefficient γ: activity coefficient εr: dielectric constant vi: molar volume

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Mean Activity Coefficients for NaCl in the Mixtures Containing Ionic

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