Phase Equilibrium, Excess Enthalpies, and Densities of Binary

Article pubs.acs.org/jced

Phase Equilibrium, Excess Enthalpies, and Densities of Binary Mixtures of Trimethylbutylammonium Bis(trifluoromethylsulfonyl)imide with Ethanol, 1‑Propanol, and Dimethylformamide Marjorie Massel, Anne-Laure Revelli, Ethan Paharik, Maribeth Rauh, Lesli O. Mark, and Joan F. Brennecke* Department of Chemical and Biomolecular Engineering, University of Notre Dame, 182 Fitzpatrick Hall, Notre Dame, Indiana 46556, United States S Supporting Information *

ABSTRACT: The molecular interactions between the ionic liquid trimethylbutylammomium bis(trifluoromethylsulfonyl)imide, [N1114][Tf2N], and three common solvents, including short chain alcohols (ethanol and 1-propanol) and an aprotic solvent (N,N-dimethylformamide or DMF) were investigated through experimental properties. The excess enthalpies were measured by calorimetry at temperatures from 308.15 K to 323.15 K over the whole composition range and excess volumes were calculated from density measurements. The vapor− liquid equilibrium (VLE) measurements of the [N1114][Tf2N] + DMF mixture were obtained using a headspace gas chromatography instrument. It was determined that hydrogen bonds between the alcohol molecules contribute to the positive excess enthalpies, while the favorable interactions between DMF and [N1114][Tf2N] might be explained by electrostatic interactions between the ionic liquid and the zwitterionic resonance structure of DMF. The nonrandom two liquid (NRTL) equation was successfully used to correlate the excess enthalpy and VLE results.

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INTRODUCTION The unique properties of ionic liquids (ILs) have opened up a plethora of new application opportunities, such as absorption refrigeration and extractive distillation.1−9 The importance of gaining knowledge of the thermodynamic properties of ILs is, therefore, increasing as their usage as solvent replacements becomes more widespread. Information on the thermodynamics of mixtures of ILs with water or organics is essential to improving process design and optimizing these technologies to reduce energy needs and remove the demand for toxic and corrosive solvents. Additionally, the thermodynamic properties of the mixtures provide insights into the molecular interactions between the IL and solvent, thus helping with the design of new solvents for these applications. Our group has recently conducted several studies of IL− water interactions based on thermodynamic measurements, thermophysical properties, and CHELPG atomic charges calculations.6,7,11 The thermodynamic measurements included excess enthalpies (HE) and vapor liquid equilibria (VLE) data. These are two different but excellent measures of the temperature dependence of the solution nonideality that tell us about the strength of the IL and solvent interactions. Several groups have previously reported excess enthalpy values of IL and organic binary mixtures, but published data remains relatively sparse compared to the abundance of available VLE data.12−14 © XXXX American Chemical Society

Here we investigate intermolecular interactions of binary mixtures of an IL with organic solvents via the direct measurement of thermodynamic and thermophysical properties. Trimethylbutylammomium bis(trifluoromethylsulfonyl) imide, [N1114][Tf2N], was mixed with ethanol, 1-propanol, and dimethylformamide (DMF), and the excess enthalpies for the three binaries were directly measured by calorimetry. Vapor−liquid equilibrium for [N1114][Tf2N] and DMF was also measured, along with densities for all three binaries. We chose the ionic liquid [N1114][Tf2N] to have a complete set of VLE, HE, and density data for these three mixtures (adding to VLE data and activity coefficients at infinite dilution previously reported for this ionic liquid).14 For comparison, we also performed some measurements with dimethyl sulfoxide (DMSO). Since experimental measurements of both excess enthalpies and VLE are now available for the three binary mixtures at two different temperatures, we will investigate the reliability of the nonrandom two liquid (NRTL) equation as a modeling tool. Received: July 24, 2014 Accepted: November 7, 2014

A

dx.doi.org/10.1021/je500696p | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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EXPERIMENTAL SECTION Chemicals. The sample information is given in Table 1. The [N1114][Tf2N] sample was checked for impurities with 1H

DMSO/IL mixture are significantly smaller than those for the three other systems, so the standard uncertainty for this mixture is u(HE) = 15 J·mol−1. To verify that the heat associated with the presence of mercury in the sample cell was negligible, several measurements were repeated with both neat IL and mercury-saturated IL. The difference in the two excess enthalpy values was within the uncertainty of the experiment. Vapor−Liquid Equilibria Measurements. The vapor− liquid equilibria measurements for [N1114][Tf2N] and DMF were taken at three temperatures (308.15 K, 313.15 K and 323.15 K) using a headspace gas chromatograph (HSGC). The apparatus consists of a Teledyne Tekmar HT3 static headspace (HS) autosampler and a Varian 450-GC gas chromatograph (GC) with a flame ionization detector (FID). The GC column was a Varian CP-Porabond Q capillary column (25 m × 0.32 mm i.d., df, 5 μm). Nine binary mixtures with varying compositions were prepared in a dry and inert environment by weighing both compounds (i.e., DMF and [N1114][Tf2N]) using a Mettler Toledo XS 205 DU analytical balance. The vials (capacity of 9 mL) were charged with 3 mL of mixture and sealed with white PTFE septa. The samples were stirred for 10 min in order to obtain a homogeneous mixture. They were placed in the carrousel of the HS and heated at constant temperature until equilibrium between the vapor and the liquid phase was reached. The optimal equilibrium time was found to be 60 min. Then, a sample of the vapor phase was taken by the HS sampler and analyzed by the GC. The bubble point pressures of the binary mixture at a constant temperature can be estimated using eq 1:

Table 1. Sample Information chemical name trimethylbutylammomium bis(trifluoromethylsulfonyl)imide ethanol 1-propanol N,N-dimethylformamide dimethyl sulfoxide

source

weight fraction

Iolitec

0.99

Merck Alpha Aesar Merck Merck

0.995 0.999 ≥0.998 ≥0.998

NMR spectroscopy, which confirmed the purity provided by the manufacturer. Nonetheless, we observed some physical changes in the [N1114][Tf2N] over time in the presence of acetone, which we typically use to clean glassware and recover ILs. 1H NMR spectroscopy confirmed the presence of small quantities of decomposition products in the IL after it had been in contact with acetone. This decomposition process was not fully understood. Therefore, we used fresh [N1114][Tf2N] for all measurements and did not attempt to recover the IL for reuse. Ethanol and 1-propanol were dried over 3 Å molecular sieves while DMF and DMSO were used as received. The IL sample was dried under reduced pressure (10 mbar) for at least 24 h before each measurement. The water content of the IL and the molecular solvents was measured with a Metrohm 831 coulometric Karl Fischer titrator before each measurement. Although all precautions were taken to keep the samples dry, the water content of the molecular solvents increased from ∼1· 10−4 to a maximum of 5·10−4 weight fraction during loading for the HE and VLE experiments. The IL water content remained below 1·10−4 weight fraction. The water contents of the mixtures used for the density measurements are listed in Tables S-1 to S-3 in the Supporting Information. Excess Enthalpy Measurements. Excess enthalpy measurements were performed in a C80 Setaram calorimeter with reversal mechanism, as described previously.6 The stainless steel mixing cells are totally surrounded by the 3D Calvet sensors. The heat flow to and from the sample cell compared to the identical empty reference cell is recorded and is then integrated over time to obtain the excess enthalpy. The calorimetric chamber was maintained at the desired temperature (308.15 K, 313.15 K, or 323.15 K) 3 h prior to the mixing and an hour after the mixing occurred. The molecular solvent was loaded in the bottom compartment of the cell in a dry and inert environment. Mercury was then used to completely separate the two liquids and, finally, [N1114][Tf2N] was added in the top compartment. The composition of the mixture was determined gravimetrically using a Mettler Toledo XS 205 DU analytical balance accurate to 1·10−4 g which translates to a standard uncertainty of u(x) = 0.0001 in the mole fraction of the mixtures. The mixing occurs by repeated inversion of the calorimeter until the maximum absolute value of the heat flux is reached. In the cell the lid separating the two compartments ensures good mixing occurs. The method was verified by measuring binary organic systems, including DMF + 1-propanol,15 which matched literature data within the estimated relative standard uncertainty of ur(HE) = 0.02. However, the excess enthalpies of the

2

P=

2

∑ pi = ∑ xi·γi·pi0 i

i

(1)

Because of the low volatility of ILs, we may assume that the vapor pressure of the IL (p02) is essentially zero, as confirmed by the lack of other peaks from the GC, and eq 1 can be reduced to P = p1 = x1·γ1·p10

(2)

In HSGC, only the volatile compounds are analyzed, which in our case is DMF. The peak area of DMF (1) is proportional to the partial pressure of the solute in the vial (Kolb):

A1 = c·p1

(3)

where c is a calibration factor. Similarly, if pure DMF is present in a vial, the peak area (A01) will be proportional to the vapor pressure of DMF at a temperature T:

A10 = c·p10

(4)

Thus, eq 1 can be rewritten as follows:

A1 = γ1·x1·A10

(5)

And the activity coefficient can be expressed as γ1 =

A1 A10 ·x1

(6)

where x1 refers to the liquid mole fraction of DMF, A1 refers to the peak area of DMF in the {DMF + [N1114][Tf2N]} sample at the equilibrium temperature T, and A01 refers to the peak area of DMF in a pure DMF sample at the same temperature. With this technique, the activity coefficients of DMF can be directly B



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Table 2. Experimental Excess Enthalpy, HE, for Three Binary Mixtures at Different Temperatures at Pressure p = 0.1 MPaa

measured. The bubble-point pressures were calculated using eq 1. Vapor pressures of pure DMF were obtained from the Antoine equation and are given in Supporting Information, Table S.4. The method was verified by measuring the vapor−liquid equilibrium of the binary mixture {methanol + [N1114][Tf2N]} at 313.15 K; the comparison with published data is shown in Figure 1.14 Our data matches the literature within the relative standard uncertainty of the pressures of ur(p) = 0.03.

308.15 K x1 0.1229 0.2546 0.2623 0.2752 0.3943 0.5084 0.5133 0.5408 0.6267 0.7252 0.7354 0.7659 0.8645 0.1224 0.2622 0.2665 0.3740 0.4927 0.5007 0.5108 0.5118 0.6191 0.7440 0.7466 0.7505 0.8709

Figure 1. Vapor−liquid equilibria for the binary [N1114][Tf2N] + methanol as a function of mole fraction of methanol at 313.15 K. Open symbols, Heintz et al.;14 filled symbols, this work.

Density Measurements. Density measurements were carried out at atmospheric pressure in a DMA 4500 Anton Paar oscillating U-tube densitometer. As described previously,10 the instrument includes an automatic correction for the viscosity of the sample and two integrated Pt 100 platinum resistance thermometers to control the temperature. The relative standard uncertainty, taking sample impurities into account, is ur(ρ) = 0.002. The mixtures were prepared in a dry and inert environment; they were kept under nitrogen in sealed vials, and stirred. The composition was determined by mass using a Mettler Toledo XS 205 DU analytical balance (u(x) = 0.0001). To verify that the mixtures were perfectly homogeneous, the density of two samples from the same batch was measured. Modeling. To complement the excess enthalpies values and VLE data reported in this work, VLE data for {ethanol + [N1114][Tf2N]} and {1-propanol + [N1114][Tf2N]}, as well as the activity coefficients at infinite dilution for the three binary mixtures, are available in the literature.14 These experimental data were used to fit two binary interaction parameters for each binary pair to the nonrandom two-liquid (NRTL) equation. In this study, the procedure (objective function, equal weight coefficients) and the assumptions (vapor phase only consists of the organic vapor that is an ideal gas at low pressures) are the same as previously reported.7

0.1285 0.2488 0.2496 0.2543 0.3983 0.5182 0.5211 0.6283 0.7679 0.8890

E

313.15 K

H /J mol

−1

x1

E

323.15 K −1

H /J mol

x1

Ethanol (1) + [N1114][Tf2N] (2) 848 0.1130 781 0.1187 1627 0.2393 1564 0.2569 1658 0.2549 1610 0.2694 1721 0.3933 2167 0.3969 2133 0.4936 2379 0.5270 2337 0.4959 2373 0.5431 2306 0.4990 2375 0.6214 2330 0.6245 2358 0.7499 2278 0.7406 2123 0.8707 2131 0.8666 1543 2104 1955 1471 1-Propanol (1) + [N1114][Tf2N] (2) 935 0.1172 869 0.1124 1826 0.2447 1750 0.2409 1848 0.2560 1803 0.2540 2327 0.2592 1825 0.2635 2604 0.3759 2341 0.2641 2607 0.4936 2621 0.3785 2625 0.5026 2616 0.4933 2636 0.6534 2536 0.6239 2534 0.7474 2269 0.6628 2225 0.8660 1593 0.7281 2230 0.7324 2178 0.8711 1481 DMF (1) + [N1114][Tf2N] (2) −413 0.1226 −402 0.1290 −699 0.2487 −699 0.2459 −702 0.2610 −730 0.2473 −701 0.3939 −969 0.3807 −972 0.5209 −1144 0.5174 −1138 0.5249 −1150 0.5224 −1138 0.6250 −1198 0.6286 −1187 0.7717 −1033 0.7693 −1058 0.8796 −726 0.8804 −676

HE/J mol−1 773 1698 1765 2251 2510 2558 2504 2292 1729

825 1672 1800 1863 1882 2400 2689 2724 2653 2476 2481 1699

−425 −706 −708 −960 −1162 −1169 −1218 −1087 −735

a

The relative standard uncertainty is ur(HE) = 0.02 and the standard uncertainties are u(x1) = 0.0001 and u(T) = 0.1 K.

[N1114][Tf2N] with alcohols are endothermic, whereas the mixture with DMF is exothermic. It has been reported previously that positive excess enthalpy values can be caused by the energy required to break the hydrogen bonds between the alcohol molecules when mixing with ILs.12 The magnitude of this contribution can be estimated from the excess enthalpy values of binary mixtures of the corresponding alcohol and an inert compound such as n-hexane.15 The maximum values of 830 J·mol−1 and 910 J·mol−1 for mixing ethanol and 1propanol,16 respectively, with an inert account for only about half of the measured excess enthalpies in the present work. Although the systems [N1114][Tf2N] + ethanol and 1-propanol are fully miscible, it appears that the interactions between the IL and the alcohols are not favorable. Previously the presence of electron withdrawing groups, such as the CF3 of the

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RESULTS AND DISCUSSION Excess Enthalpy. Experimental excess enthalpies are listed in Table 2 and are represented graphically in Figure 2. The solid lines correspond to the NRTL model correlations (the NRTL binary interaction parameters can be found in the Supporting Information Table S-5). Binary mixtures of C



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Figure 2. Excess enthalpy, HE, for the binary systems solvent (1) + [N1114][Tf2N] (2) as a function of mole fraction of solvent, at different temperatures. (a) The four binary systems at 308.15 K: ◊ magenta, DMF; □ blue, DMSO; ○ red, ethanol; and △ green, 1-propanol. (b) Ethanol, (c) 1-propanol, and (d) DMF with ○ green, 308.15 K; △ blue, 313.15 K; □ red, 323.15 K. The solid curves represent NRTL model correlations.

shown in Figure 3, as this rearrangement is more likely to happen in an ionic environment. Since dimethyl sulfoxide,

trifluoromethanesulfonate anion, have been shown to reduce the interactions of ILs with water, therefore increasing the excess enthalpy.11 The delocalized charge on the [Tf2N]− anion seems to have a similar effect on the alcohol/IL interactions by increasing the positive excess enthalpy. The cation can also have an effect on these interactions since the excess enthalpies of 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl) imide ([emim][Tf2N]) and ethanol were smaller than that of [N1114][Tf2N] at lower temperature.17 This is possibly because the tetraalkylammonium cation cannot form hydrogen bonds with the alcohol unlike the [emim]+ cation, which has an acidic proton at the C2 position. The exothermicity of the binary mixture DMF + [N1114][Tf2N] indicates that, to a first approximation, the IL/DMF interactions outweigh the IL/IL and DMF/DMF interactions. Even though the DMF molecules do not form hydrogen bonds as strong as alcohols, evidence of hydrogen bonding in DMF dimers has been reported in the literature. For instance, Vargas et al. showed the presence of interactions between the formamide hydrogen and the oxygen atom of DMF by calculating the distances between the atoms for different configurations and comparing them to van der Waals contact distances.18 Out of the four geometries studied, three exhibited shorter distances. Since the amide nitrogen group makes the proton more acidic, the formamide hydrogen of DMF can act as a hydrogen bond donor forming a hydrogen bond with the carbonyl oxygen of DMF. The relatively strong IL/DMF interactions might involve the DMF zwitterion structure,

Figure 3. DMF resonance structures.

which is similar to DMF, does not have a zwitterion resonance structure, excess enthalpy measurements were taken for the DMSO + [N1114][Tf2N] mixture at 308.15 K. The excess enthalpies values were significantly less negative than with DMF, confirming the hypothesis of zwiterrion/IL interactions playing a role in the exothermic nature of the DMF + [N1114][Tf2N] system. The raw excess enthalpy data for the DMSO + [N1114][Tf2N] binary mixture are listed in Table 3 and graphically represented in Figure 2a. Interestingly, published excess enthalpy values for 1-butyl-3-methylimidazolium chloride and acetate with DMF and DMSO show a different trend, suggesting stronger interactions between the IL and DMSO than with DMF.19 Perhaps the zwitterion/IL interactions are in that case minimized since the electrostatic interactions between the anion and cation are stronger than for [N1114][Tf2N]. For the three systems the temperature dependence of the excess enthalpies is very small and even inexistent for the [N1114][Tf2N] + DMF mixture. There is a distinct lack of excess enthalpy data for mixtures of ionic liquids and organic solvents in the literature. Therefore, D



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Table 3. Experimental Excess Enthalpy, HE, for the Binary System DMSO (1) + [N1114][Tf2N] (2) at 308.15 K and at Pressure p = 0.1 MPaa x1

HE/J mol−1

x1

HE/J mol−1

0.1288 0.2255 0.2609 0.3095 0.3626 0.4382 0.4941 0.4983 0.5007

−113 −174 −192 −214 −248 −263 −279 −267 −269

0.5590 0.5642 0.5646 0.6132 0.6267 0.7533 0.7693 0.8684

−286 −284 −289 −263 −259 −167 −155 −65

Table 4. Vapor−Liquid Equilibria for the Binary DMF (1) + [N1114][Tf2N] (2) System at Three Different Temperaturesa x1 0.1098 0.2071 0.3161 0.4067 0.5090 0.6071 0.7031 0.8024 0.9002

a The standard uncertainties are u(HE) = 15 J·mol−1, u(x1) = 0.0001 and u(T) = 0.1 K.

0.1098 0.2153 0.3069 0.4086 0.5100 0.6059 0.7074 0.7430 0.8027 0.8997

our excess enthalpy values were converted to partial molar enthalpies for comparison to available data and are shown in Supporting Information, Figure S-1.20 The data for ethanol is in reasonable agreement with Balantseva et al. but the 1-propanol data does not compare well. However, the difference in measurement conditions (temperature and composition) may explain this discrepancy. Vapor−Liquid Equilibria. The experimental P−x−γ data at three different temperatures are listed in Table 4 and plotted in Figure 4. Not surprisingly {DMF + [N1114][Tf2N]} shows a negative deviation from Raoult’s law, indicating strong interactions between DMF and the IL (i.e., experimental activity coefficients for DMF are lower than 1). The pressure values increase with the temperature (Figure 4); however, this is due to the vapor pressure of DMF, which is the dominant property, increasing while the activity coefficients remain relatively insensitive to temperature (Table 4). This means that there is no thermodynamic inconsistency between the negative enthalpy values obtained by calorimetry and the VLE data. Density. The experimental density data for the three mixtures and the pure components are listed in Tables 5 to 7. The density of our [N1114][Tf2N] sample is in remarkably good agreement with previously reported data.21−24 The density of the mixtures decreases as the temperature and the mole fraction of the solvent (alcohol or DMF) increase. The excess molar volumes were calculated from the density measurements and are presented graphically in Figure 5 as a function of the composition. As expected, the excess molar volumes are very small for the three systems with the values between −0.4 cm3· mol−1 and 0.2 cm3·mol−1. Interestingly the signs of the molar excess volumes are the opposite of the signs of the excess enthalpies. They are negative for the mixtures with alcohols and positive for the mixture with DMF. This trend was seen before for a series of binary mixtures of IL and ethanol.13 As it is easily seen in Figure 5a the dependence of the excess volumes on composition is different for the three binary mixtures. The volumes of the [N1114][Tf2N] + ethanol system are negative over the entire range of composition with minimum values around 0.9 mole fraction ethanol. Although the same behavior is seen for 1-propanol in the dilute region, the overall shape is sinusoidal with positive values in the IL-rich region up to 0.25 mole fraction. Similar S-shape curves were previously reported for binary mixtures of [emim][Tf2N] + 1-propanol and 1butanol and [hmim][Tf2N] + methanol, ethanol, 1-propanol, and 2-propanol.24,25 The magnitude of the excess volumes is

0.1017 0.2155 0.3405 0.4102 0.4948 0.6046 0.7032 0.8048 0.9003

p/Pa 308.15 K 14 33 62 105 199 331 477 659 840 313.15 K 18 43 74 139 258 427 630 707 861 1093 323.15 K 28 74 157 247 397 707 1091 1479 1892

γ1 0.130 0.161 0.198 0.260 0.392 0.547 0.680 0.824 0.936 0.124 0.150 0.181 0.254 0.378 0.525 0.665 0.710 0.800 0.907 0.118 0.146 0.197 0.258 0.343 0.501 0.664 0.786 0.899

a

The relative standard uncertainty is ur(p) = 0.03 and the standard uncertainties are u(x1) = 0.0001 and u(T) = 0.1 K.

Figure 4. Vapor−liquid equilibria for the binary [N1114][Tf2N] + DMF system as a function of mole fraction of DMF, at three different temperatures: ○ green, 308.15 K; △ blue, 313.15 K; □ red, 323.15 K. Solid lines: NRTL model correlations.

significantly smaller for 1-propanol than for ethanol but an increase in the length of the alkyl chain tends to give more positive excess molar volume.24−26 The excess molar volumes for the [N1114][Tf2N] + DMF mixtures are positive over the entire range of composition, except in the dilute region at E



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Table 5. Experimental Density, ρ, and Calculated Excess Molar Volume, VE, for the Binary System Ethanol (1) + [N1114][Tf2N] (2) at Different Temperatures and at Pressure p = 0.1 MPaa T/K x1

293.15

295.15

298.15

303.15

308.15

313.15

318.15

323.15

328.15

333.15

ρ/g·cm−3 0.0000 0.1284 0.2585 0.3814 0.5031 0.6487 0.7823 0.8561 0.8763 0.9039 0.9500 0.9745 1.0000 0.1284 0.2585 0.3814 0.5031 0.6487 0.7823 0.8561 0.8763 0.9039 0.9500 0.9745 a

1.3971 1.3795 1.3571 1.3297 1.2939 1.2321 1.1420 1.0663 1.0406 1.0001 0.9164 0.8602 0.7896 −0.055 −0.107 −0.138 −0.195 −0.248 −0.287 −0.298 −0.305 −0.285 −0.229 −0.159

1.3953 1.3777 1.3553 1.3278 1.2920 1.2303 1.1401 1.0645 1.0388 0.9983 0.9146 0.8584 0.7879 −0.054 −0.104 −0.135 −0.194 −0.246 −0.288 −0.300 −0.307 −0.287 −0.229 −0.159

1.3927 1.3751 1.3526 1.3251 1.2893 1.2275 1.1373 1.0617 1.0360 0.9956 0.9119 0.8558 0.7853 −0.051 −0.103 −0.134 −0.192 −0.247 −0.291 −0.304 −0.311 −0.292 −0.234 −0.163

1.3883 1.3706 1.3481 1.3206 1.2847 1.2228 1.1327 1.0571 1.0314 0.9910 0.9074 0.8514 0.7810 −0.049 −0.102 −0.132 −0.193 −0.249 −0.297 −0.311 −0.320 −0.300 −0.241 −0.167

1.3840 1.3662 1.3437 1.3161 1.2801 1.2182 1.1280 1.0524 1.0268 0.9864 0.9029 0.8469 0.7767 VE/cm3·mol−1 −0.049 −0.101 −0.132 −0.197 −0.255 −0.305 −0.320 −0.328 −0.308 −0.247 −0.171

1.3796 1.3618 1.3392 1.3116 1.2756 1.2136 1.1233 1.0478 1.0221 0.9818 0.8984 0.8425 0.7723 −0.052 −0.103 −0.137 −0.201 −0.262 −0.314 −0.330 −0.339 −0.318 −0.256 −0.177

1.3752 1.3574 1.3348 1.3071 1.2710 1.2090 1.1186 1.0431 1.0174 0.9771 0.8938 0.8380 0.7679 −0.052 −0.109 −0.141 −0.208 −0.273 −0.324 −0.342 −0.350 −0.329 −0.266 −0.184

1.3709 1.3531 1.3304 1.3026 1.2665 1.2044 1.1140 1.0384 1.0127 0.9724 0.8892 0.8334 0.7635 −0.054 −0.113 −0.148 −0.215 −0.283 −0.337 −0.355 −0.363 −0.341 −0.276 −0.191

1.3666 1.3487 1.3259 1.2981 1.2619 1.1997 1.1092 1.0337 1.0080 0.9677 0.8845 0.8288 0.7589 −0.059 −0.121 −0.155 −0.226 −0.296 −0.351 −0.370 −0.378 −0.356 −0.286 −0.197

1.3623 1.3443 1.3215 1.2937 1.2574 1.1951 1.1045 1.0289 1.0033 0.9630 0.8798 0.8242 0.7544 −0.064 −0.129 −0.166 −0.238 −0.312 −0.370 −0.387 −0.395 −0.371 −0.299 −0.207

The relative standard uncertainty is ur(ρ) = 0.002 and the standard uncertainties are u(x1) = 0.0001 and u(T) = 0.05 K.

Table 6. Experimental Density, ρ, and Calculated Excess Molar Volume, VE, for the Binary System 1-propanol (1) + [N1114][Tf2N] (2) at Different Temperatures and at Pressure p = 0.1 MPaa T/K x1

293.15

295.15

298.15

303.15

308.15

313.15

318.15

323.15

328.15

333.15

−3

0.0000 0.1273 0.2596 0.3787 0.4996 0.6368 0.7624 0.8787 0.9498 0.9799 1.0000 0.1273 0.2596 0.3787 0.4996 0.6368 0.7624 0.8787 0.9498 0.9799 a

1.3971 1.3750 1.3469 1.3150 1.2739 1.2101 1.1260 1.0087 0.9039 0.8473 0.8036 0.016 −0.007 −0.015 −0.068 −0.076 −0.088 −0.102 −0.101 −0.068

1.3953 1.3732 1.3451 1.3132 1.2721 1.2083 1.1242 1.0069 0.9022 0.8457 0.8021 0.020 −0.003 −0.001 −0.067 −0.073 −0.087 −0.102 −0.101 −0.069

1.3927 1.3706 1.3424 1.3106 1.2694 1.2056 1.1214 1.0043 0.8997 0.8432 0.7997 0.021 0.000 −0.015 −0.063 −0.069 −0.079 −0.102 −0.103 −0.069

1.3883 1.3662 1.3380 1.3061 1.2649 1.2011 1.1170 0.9999 0.8955 0.8391 0.7956 0.021 0.002 −0.012 −0.059 −0.066 −0.085 −0.105 −0.105 −0.072

ρ/g·cm 1.3840 1.3618 1.3335 1.3016 1.2604 1.1965 1.1125 0.9955 0.8912 0.8350 0.7916 VE/cm3·mol−1 0.024 0.006 −0.011 −0.055 −0.061 −0.082 −0.105 −0.107 −0.072

1.3796 1.3574 1.3291 1.2972 1.2559 1.1920 1.1079 0.9911 0.8870 0.8308 0.7875 0.019 0.004 −0.013 −0.064 −0.064 −0.087 −0.112 −0.115 −0.077

1.3752 1.3530 1.3247 1.2927 1.2514 1.1874 1.1034 0.9867 0.8826 0.8266 0.7834 0.020 0.004 −0.014 −0.066 −0.065 −0.089 −0.117 −0.119 −0.080

1.3709 1.3486 1.3203 1.2882 1.2469 1.1829 1.0988 0.9822 0.8783 0.8224 0.7792 0.018 −0.002 −0.019 −0.069 −0.072 −0.091 −0.123 −0.126 −0.083

1.3666 1.3443 1.3159 1.2838 1.2424 1.1783 1.0943 0.9777 0.8739 0.8180 0.7749 0.013 −0.007 −0.028 −0.078 −0.078 −0.106 −0.137 −0.136 −0.090

1.3623 1.3399 1.3115 1.2794 1.2379 1.1738 1.0897 0.9731 0.8695 0.8137 0.7706 0.008 −0.020 −0.035 −0.085 −0.088 −0.112 −0.143 −0.144 −0.095


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Table 7. Experimental Density, ρ, and Calculated Excess Molar Volume, VE, for the Binary System DMF (1) + [N1114][Tf2N] (2) at Different Temperatures and at Pressure p = 0.1 MPaa T/K x1

a

293.15

295.15

298.15

303.15

0.0000 0.1268 0.2011 0.3028 0.3824 0.5078 0.6249 0.7548 0.8764 0.9157 1.0000

1.3971 1.3797 1.3677 1.3487 1.3313 1.2974 1.2558 1.1916 1.1013 1.0619 0.9487

1.3953 1.3779 1.3659 1.3469 1.3295 1.2956 1.2539 1.1897 1.0994 1.0600 0.9468

1.3927 1.3753 1.3632 1.3442 1.3268 1.2928 1.2512 1.1868 1.0965 1.0571 0.9439

1.3883 1.3708 1.3588 1.3397 1.3222 1.2882 1.2465 1.1821 1.0918 1.0524 0.9391

0.1268 0.2011 0.3028 0.3824 0.5078 0.6249 0.7548 0.8764 0.9157

0.065 0.119 0.167 0.184 0.207 0.195 0.145 0.062 0.030

0.067 0.121 0.171 0.187 0.210 0.197 0.145 0.061 0.028

0.070 0.126 0.173 0.190 0.210 0.197 0.143 0.058 0.026

0.071 0.129 0.178 0.195 0.213 0.197 0.139 0.052 0.020

308.15

313.15

ρ/g·cm−3 1.3840 1.3796 1.3664 1.3620 1.3543 1.3499 1.3352 1.3307 1.3177 1.3131 1.2836 1.2790 1.2418 1.2372 1.1774 1.1727 1.0870 1.0823 1.0476 1.0428 0.9344 0.9296 VE/cm3·mol−1 0.072 0.073 0.130 0.129 0.180 0.182 0.197 0.197 0.213 0.211 0.194 0.191 0.134 0.128 0.044 0.037 0.012 0.006

318.15 1.3752 1.3576 1.3454 1.3262 1.3086 1.2745 1.2326 1.1680 1.0775 1.0381 0.9248 0.076 0.130 0.180 0.194 0.207 0.185 0.120 0.026 −0.003

323.15 1.3709 1.3532 1.3410 1.3218 1.3041 1.2699 1.2279 1.1633 1.0728 1.0333 0.9200 0.074 0.129 0.179 0.190 0.203 0.180 0.111 0.018 −0.014

328.15 1.3666 1.3488 1.3366 1.3173 1.2997 1.2654 1.2233 1.1586 1.0680 1.0286 0.9152 0.072 0.125 0.175 0.186 0.196 0.171 0.101 0.005 −0.024

333.15 1.3623 1.3445 1.3323 1.3129 1.2952 1.2608 1.2187 1.1539 1.0633 1.0238 0.9104 0.071 0.122 0.169 0.181 0.188 0.161 0.088 −0.007 −0.035


Figure 5. Excess molar volume, VE, for the binary systems solvent (1) + [N1114][Tf2N] (2) as a function of mole fraction of solvent, at different temperatures. (a) The three binary systems at 308.15 K: △ green, ethanol; □ magenta, 1-propanol; and □ blue, DMF. (b) Ethanol, (c) 1-propanol, and (d) DMF with ○, 293.15 K; □, 303.15 K; △, 313.15 K; □, 323.15 K; ◊, 333.15 K.

temperatures higher than 318.15 K. Although the excess volumes are very small, it is somewhat surprising that they are

positive. The negative deviation from ideal behavior for VLE of this system and the negative excess enthalpies show evidence of G



Article

[N1114][Tf2N] + DMF system (AARD values below 15 % for the mixtures with alcohols and between 74 % and 88 % for the DMF mixtures). The measured pressures of DMF + [N1114][Tf2N] are 1 order of magnitude smaller than the pressures of the mixtures with alcohols, and might be more difficult to predict. The AARD values using this method (fitting NRTL parameters to HE and γ∞ data and predicting VLE) are shown in Table 9. In a previous study the nonrandomness parameter was fit as a third temperature-dependent parameter,14 whereas here it is set to 0.1 for the [N1114][Tf2N] + alcohol systems. Interestingly, if the fit α parameters from Heintz et al. are used to predict the HE values from NRTL parameters fit to the VLE and γ∞ data, the results improve drastically. For ethanol the AARD for the HE data predictions decreases from 89 % to 13 % at 308.15 K, and from 97 % to 22 % for 1-propanol at the same temperature. Attempts to model the [N1114][Tf2N] + ethanol, [N1114][Tf2N] + 1-propanol and [N1114][Tf2N] + DMF systems with the NRTL equation suggest that it is not an appropriate model for these IL + solvent systems. It is possible to correlate the data for any given property (e.g., VLE or HE), but it is not possible to adequately predict one property from parameters obtained from another property. Nonetheless, the dramatic improvement in the HE predictions when using the nonrandomness parameters fit to the VLE data suggests that at least part of the problem may be in the magnitude of the local composition enhancement, which is what α represents. All the binary parameters can be found in the Supporting Information in Tables S-5 and S-6.

favorable IL/DMF interactions. Apparently, the IL and DMF molecules do not pack as efficiently as the IL and alcohol molecules. Modeling. Fitting either the experimental VLE or the experimental HE data to the NRTL model equation yielded excellent results. The obtained binary interaction parameters, which are shown in Table S.5 of Supporting Information, were used to correlate the experimental values. Nonrandomness factors, α, of 0.1 for the two [N1114][Tf2N] + alcohol systems and 0.3 for the [N1114][Tf2N] + DMF or DMSO mixtures were found to be the most suitable. As seen in Figures 2 and 4, the correlated bubble-point curve and excess enthalpy values are in good agreement with the experimental data. The absolute average relative deviation (AARD) percentages (calculated with eq 7 and listed in Table 8) indicate that the model performs Table 8. Absolute Average Relative Deviations for Excess Enthalpy and Pressure Correlations for the Binary Solvent (1) + [N1114][Tf2N] (2) ethanol (1)

1-propanol (1)

DMF (1)

T/K

AARD % for HE

AARD % for HE

AARD % for HE

AARD % for p

308.15 313.15 323.15

1.91 2.03 3.80

1.59 2.10 2.38

1.97 2.36 2.33

4.97 4.93 4.83

similarly for the three systems (AARD values are less than 2.50 %), with the exception of {ethanol + [N1114][Tf2N]} at 323.15 K, that exhibits an AARD of 3.80 %. Note that the binary interaction parameters needed to fit the HE and VLE data for the [N1114][Tf2N] + DMF mixture are very different in magnitude (see Table S.5 of Supporting Information). Infinite dilution activity coefficients, γ∞, are available for ethanol, 1-propanol, and DMF in [N1114][Tf2N], as well as VLE for the [N1114][Tf2N] + ethanol and [N1114][Tf2N] + 1propanol systems.14 Therefore, we investigated whether it is possible to predict the HE values measured here from NRTL parameters fit to just VLE and γ∞ data. These values are shown in Table S.6 of Supporting Information. Since HE values are a temperature derivative of the excess Gibbs energy (and, therefore, the activity coefficients that determine VLE), they are difficult to predict;7 therefore, a large absolute average standard deviation was expected. This is what was observed: around 100 % AARD for the three binary mixtures, as seen in Table 9, which demonstrates the inability of the NRTL model to accurately predict HE from experimental VLE and γ∞ data. The NRTL model performs better for VLE prediction using parameters fit to HE and γ∞ data (parameters also shown in Table S.6), but the representation is still not good for the

AARD % =

1-propanol (1)

DMF (1)

AARD %

AARD %

AARD %

T/K

HE

p

HE

p

HE

p

308.15 313.15 323.15

90 84

9.9 10

97 103

14 15

123 133 141

74 77 88

n

∑ i=0

M* − M M*

(7)

where M can be p or H, n denotes the number of experimental data, and the asterisk denotes experimental quantity.

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CONCLUSION New experimental excess enthalpies, densities, and VLE data were reported for binary mixtures of [N1114][Tf2N] with ethanol, 1-propanol, and DMF. These thermodynamic properties were used to probe the molecular interactions between the species. Hydrogen bonds between the alcohol molecules are broken upon mixing, resulting in positive excess enthalpies. The negative excess enthalpies and activity coefficients less than 1 for the [N1114][Tf2N] + DMF system suggest favorable interactions between the IL and the aprotic solvent. There might be indeed electrostatic interactions between the DMF zwitterion structure and the ionic liquid. The excess volumes were small but have opposite signs from the HE values for ethanol and DMF; the VE values of the mixture with 1-propanol were negligible. Attempts to model the three binary mixtures with the NRTL equation suggest that it is an appropriate model to correlate HE data but it is not possible to predict a property using parameters fit to another property.

Table 9. Absolute Average Relative Deviations for Excess Enthalpy and Pressure Predictions for the Binary Solvent (1) + [N1114][Tf2N] (2) When the NRTL Parameters are Fit to γ∞ and VLE or HE Data ethanol (1)

100 n

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ASSOCIATED CONTENT

S Supporting Information *

Water content of the binary mixtures used for the density measurements, NRTL parameters, and comparison with published partial molar enthalpy data. This material is available free of charge via the Internet at http://pubs.acs.org. H



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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Funding

This material is based upon work supported by the Department of Energy under Award No. DOE FG02-05CH11294. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Luke Simoni for his help with programming and Aruni DeSilva for her help with NMR analysis.

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REFERENCES

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