Effect of Acetonitrile on the Solubility of Carbon Dioxide in 1-Ethyl-3

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Effect of Acetonitrile on the Solubility of Carbon Dioxide in 1-Ethyl-3-methylimidazolium Bis(trifluoromethylsulfonyl)amide G. Hong, J. Jacquemin, P. Husson, M. F. Costa Gomes,* M. Deetlefs,† M. Nieuwenhuyzen,† O. Sheppard,† and C. Hardacre† Laboratoire de Thermodynamique des Solutions et des Polyme` res, UMR 6003 CNRS/UniVersite´ Blaise Pascal, Clermont-Ferrand, 24 aVenue des Landais, 63177 Aubie` re Cedex, France

The effect of the addition of acetonitrile on the solubility of carbon dioxide in an ionic liquid, the 1-ethyl3-methylimidazolium bis(trifluoromethanesulfonyl)amide, [C2mim][NTf2], was studied experimentally at pressures close to atmospheric and as a function of temperature between 290 and 335 K. It was observed that the solubility of carbon dioxide decreases linearly with the mole fraction of acetonitrile from a value of 2.6 × 10-2 in the pure ionic liquid at 303 K to a mole fraction of 1.3 × 10-2 in the mixture [C2mim][NTf2] + CH3CN with xCH3CN ) 0.77 at the same temperature. The gas solubility decreases with temperature, and the thermodynamic properties of solvation could be calculated. The vapor pressures of the [C2mim][NTf2] + CH3CN mixtures were measured in the same temperature range, and strong negative deviations from Raoult’s law were obtained: up to 36% for a mixture with xCH3CN ) 0.46 at 334 K. Negative excess molar volumes of approximately -1 cm3 mol-1 at equimolar composition could also be calculated from density measurements of the pure components and of the mixtures. These observations are confirmed by neutron diffraction studies and are compatible with the existence of strong ion-dipole interactions in the mixed liquid solvent. 1. Introduction Over the past decade, the interest in the use of ionic liquids as novel solvents and materials with negligible vapor pressure has increased significantly,1-4 with the vast majority of research concentrated on their implementation as new reaction media.5 In that context, the knowledge of their thermophysical properties, in particular their phase behavior with gases and molecular solvents, is still limited despite its importance if the use of ionic liquids as solvents is to become widespread commercially. The study of gas solubility is a particular case where the data can be both of fundamental and practical utility providing information about the structure and interactions in solution6 as well as allowing the calculation of (vapor + liquid) equilibria in systems of technological interest, for example, when the development of new separation and reaction media5 is to be performed. This study reports on the solubility of carbon dioxide in a typical ionic liquid based on the 1-ethyl-3-methylimidazolium ([C2mim]+) cation and the bis(trifluoromethanesulfonyl)amide anion ([(CF3SO2)2N]- or [NTf2]-)ssee Figure 1sin order to investigate the effect on this property of the addition of a molecular cosolvent, acetonitrile (CH3CN). This system is considered as closely approximate to a real reaction mixture and the characterization of its saturation and volumetric properties provides insights about the molecular mechanisms involved in the solvation of a gas in a mixture of liquid solvents one of which being an ionic liquid. Carbon dioxide is the gaseous solute most studied in several families of imidazolium based ionic liquids and in particular those based on the [NTf2] anion.7-11 By analyzing the data already available in the literature, it can be observed that carbon dioxide is one of the most soluble gases in ionic liquids with mole fraction solubilities of the order of 10-2, that decrease * To whom correspondence should be addressed. Tel.: +33473407205. Fax: +33473405328. E-mail: [email protected]. † QUILL Research Center, School of Chemistry and Chemical Engineering, Queen’s University of Belfast, United Kingdom.

Figure 1. Structural formula of the ionic liquid 1-ethyl-3-methylimidazolium bis(trifluoromethanesulfonyl)amide anion.

with temperature. It is also noteworthy that it is the anion in the ionic liquid that determines the solubility of carbon dioxide11 and that the effect of changing the cation of the ionic liquid is much less significant.10 The phase behavior of binary mixtures of ionic liquids with several molecular organic compounds (hydrocarbons,12 aldehydes and ketones,13 and several oxygenated organic compounds14,15) has been studied by different research groups. Heintz and co-workers have measured the activity coefficient at infinite dilution of the binary system [C2mim][NTf2] + CH3CN from 313 to 363 K16 and obtained a value of γ∞i of 0.444 at 298 K. Such a relatively low value is consistent with the high solubility of acetonitrile in the ionic liquid. Ternary systems involving ionic liquids have also been studied. Hert et al.17 have measured the enhancement of the solubility of methane and oxygen due to the presence of carbon dioxide dissolved in the ionic liquid [C6mim][NTf2]. Doker and Gmehling18 have studied the vapor-liquid equilibria of ternary systems involving one ionic liquid and two molecular compounds: acetone, methanol, or water. To date, the only studies involving ternary systems of an ionic liquid, a molecular solvent, and a gas concern the determination of the influence of the water content on the solubility of carbon dioxide in two ionic liquids: [C4mim][PF6]19,20 and [C4mim][NTf2].21 Herein, we report the determination of vapor-liquid equilibria data for the binary liquid [C2mim][NTf2] + CH3CN as a function of temperature and composition. The volumetric properties of the same mixtures (and of the two pure components) were also measured as a function of the composition at two temperatures (303 and 323 K), and the excess molar volumes are reported and discussed. Finally, vapor-liquid

10.1021/ie0605377 CCC: $33.50 © 2006 American Chemical Society Published on Web 10/24/2006

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equilibria data were determined for the ternary system [C2mim][NTf2] + CH3CN + CO2 and the solubility of the gaseous component in the liquid mixtures with distinct compositions is reported and compared with that found in the pure ionic liquid (also determined in the present study). 2. Experimental Section Materials. The ionic liquid 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide, [C2mim][NTf2], was prepared using the standard synthesis procedures described by Bonhoˆte et al.22 The bromide content of the ionic liquid sample was determined using ion chromatography23 and was less than 5 ppm. The water content was determined before and after the solubility measurements by coulometric Karl-Fisher titration using a Mettler Toledo DL31 titrator. A reference value of 45 ppm (w/w) of water was obtained after drying the liquid in vacuo at 303 K for 8 h. The acetonitrile used, Janssen Chimica, (>0.995 mole fraction) was dried for 48 h over CaH2 and then distilled before use, ultimately containing 600 ppm (w/w) H2O, as determined by Karl-Fisher titration. The carbon dioxide was purchased from AGA/Linde Gaz (>0.99995 mole fraction) and was used as received. Gas Solubility Measurements. The experimental apparatus used for the low-pressure gas solubility measurements is based on an isochoric saturation technique that has been described in previous publications.24,25 In this technique, a known quantity of gaseous solute is put in contact with a precisely determined quantity of degassed solvent at constant temperature inside an accurately known volume. When thermodynamic equilibrium is attained, the pressure above the liquid solution is constant and is directly related to the solubility of the gas in the liquid. The vacuum/gas line and the equilibrium cell previously used15 had to be redesigned in the present arrangement in order to permit the work with multicomponent solvents. The experimental apparatus used in this work is schematically represented in Figure 2. The equilibrium cell (EC) together with the precision manometer (M) and the glass bulb (GB), limited by valve V2, constitute the equilibrium section of the apparatus which is bordered by valves V1 and the one leading to connection C3. This new design of the equilibrium cell allows the connection of the bulb OS to the vacuum/gas line for degassing the organic solvent used through connection C4, or to the equilibrium cell for preparing the liquid mixture through connection C3. The equilibrium cell permits the handling of total volumes of liquid solvent varying from 5 to 25 mL, and appropriate gas/liquid contact is guaranteed by means of good agitation using a glass coated magnetic follower. The whole equilibrium section is maintained at constant temperature to within (0.01 K using a liquid thermostat and a PID (proportional-integral-derivative) temperature controller, and its value is measured accurately by means of a calibrated (against a primary standand platinum resistance thermometer (PRT) from Tinsley with a precision of (0.02 K) 100 Ω platinum resistance sensor associated with a 61/2 digit multimeter (Keithley model 2000). The solubility measurement starts with the introduction of a known quantity of gas solute in the gas bulb (GB) whose volume has been calibrated as a function of temperature with a precision better than 0.1% (VGB ) (90.39 ( 0.01) cm3 at 303.15 K with a thermal expansion coefficient R ) 2.76 × 10-5). The exact amount of gas introduced is determined by measuring its pressure in the manometer (M) (Druck RPT 350 model, calibrated for pressures between 35 and 3500 mbar with a precision of 0.01% full scale) at constant temperature and

Figure 2. Solubility apparatus used in this work: (VP) vacuum pump; (TP) cold trap; (VG) vacuum gauge; (GB) gas bulb; (OS) molecular solvent bulb; (M) precision manometer; (TB) thermostated bath; (EC) equilibrium cell; (R) heating element; (V1, V2) constant volume glass valves; (C1, C2, C3, C4) vacuum o-ring connections.

correcting for gas imperfections. The gas is isolated from the rest of the installation by closing the glass valve (V2). The ionic liquid is then introduced in the equilibrium cell through connection C2 by means of a syringe, its mass being determined with a precision of 1 × 10-4 g. The ionic liquid is degassed and dried by evacuating it (pressure of 1 Pa) for 15 h at a temperature above 303 K. For the measurements of gas solubility in binary liquid mixtures as solvents, the more volatile liquid (acetonitrile in the present case) is introduced in the OS bulb and is degassed by successive melting/freezing cycles while vacuum pumping noncondensable gases though connection C4. This degassed liquid is then transferred to the equilibrium cell under its own vapor pressure through connection C3. The amount of liquid introduced in the equilibrium cell at this stage is determined by weighing the OS bulb right after degassing the solvent and then after its introduction into the equilibrium cell. The gas/liquid equilibrium process starts by bringing into contact the solvent mixture (previously equilibrated by stirring the two liquids at constant temperature and waiting for a constant pressure to be reached) with the gaseous solute by opening valve V2 (a constant volume valve). The total volume of the equilibrium section was previously calibrated by gas expansions from the gas bulb at different temperatures in order to appropriately take into account the thermal expansion corrections (Vtot ) (143.6 ( 0.1) cm3 at 303.15 K with R ) 1.35 × 10-4). To avoid condensation of the solvent (mixture of the ionic liquid with the volatile molecular organic solvent) in the manometer during the solubility measurements, care was taken to control the temperature in the manometer ca. 0.5 K above the liquid thermostat using a variable resistance (R in Figure 2). The determination of the solubility at different temperatures is done by changing the liquid thermostat set point and waiting for a new thermodynamic equilibrium at a different temperature. The mole fraction solubility, x2, is calculated from the amount of solute present in the liquid solution, nliq 2 through

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x2 )

nliq 2 liq (nliq solv + n2 )

(1)

liq liq where nliq solv ) n1 + n3 is the amount of solvent in the liquid liq phase with n1 being the amount of ionic liquid which, due to its negligible vapor pressure, is equal to the total quantity of ionic liquid introduced in the equilibrium cell and nliq 3 the amount of acetonitrile in the liquid solution (this value is, of course, zero in the case of the gas solubility measurements in pure ionic liquids). The amount of solute present in the liquid solution is calculated from two pVT measurements: first, when the gas is introduced in the calibrated bulb with volume VGB and, second, when thermodynamic equilibrium is reached. In this calculation, it is assumed that Raoult’s law is valid for the system gas/solvent:26,27

nliq 2

)

ntot 2

-

ngas 2

piniVGB

)

-

[Z2(pini,Tini)RTini] [peq - xsolvpsat solv](Vtot - Vliq) [Z23(peq,Teq)RTeq]

(2)

where pini and Tini are the pressure and temperature in the first pVT determination and peq and Teq are the pressure and temperature at equilibrium. Vtot is the total volume of the equilibration cell, and Vliq is the volume occupied by the liquid solution with vapor pressure psat solv. Vliq is always considered to be equal to the volume of the pure solvent as it is assumed that the gas dissolved does not change significantly the volume of the liquid phase. In the case of multicomponent liquid solvents, this volume has to be determined as, apart from the molar volumes of the two components of the solvent mixture, it is necessary to take into account the volume change on mixing. Z2 is the compressibility factor for the pure gas, and Z23 is the compressibility factor for the vapor phase in equilibrium with the solution (equals Z2 in the case of the solubility of gases in pure ionic liquids) given by

Z23 ) 1 +

p (y B + y3B33 + y2y3δ23) RT 2 22

(3)

where B22 and B33 are the second virial coefficients for CO2 and CH3CN, respectively, and δ23 ) 2B23 - B22 - B33 with B23 being the crossed virial coefficient considered, in the present case, as the average between B22 and B33. The amount of solvent present in the liquid phase, nliq solv, can be calculated from (by assuming, once again, the validity of Raoult’s law for the solute and the solvent in equilibrium) tot gas tot nliq solv ) nsolv - nsolv ) nsolv -

xsolvpsat solv(Vtot - Vliq) [Z23(peq,Teq)RTeq]

(4)

where ntot solv is the total amount of solvent initially introduced in the equilibrium cell and Z23 is the compressibility factor for the solvent calculated using eq 3. The equilibrium compositions are calculated iteratively with the process starting with initial guesses for the mole fractions in both phases and continuing with the determination of the quantities of solute and solvent present in the liquid solution and in the vapor in equilibrium with it (eqs 2 and 4), provided that the vapor pressure of the solvent is known as well as its molar volume. The mole fraction solubility, calculated by eq 1, can then be used for the determination of the Henry’s law constant, KH:28

f2(p,T,x2) φ2(peq,Teq)y2peq = x2f0 x2 x2

KH ≡ lim

(5)

where f2 is the fugacity of the solute and φ2 is its fugacity coefficient. The Henry’ law constants can be exactly converted to the Gibbs energy of solvation, ∆solvG, corresponding to the change in Gibbs energy when the solute is transferred, at constant temperature, from the pure perfect gas state at the standard pressure to the standard state of infinite dilution of the solute in the solvent (in the present case, the solvent is a mixture of an ionic liquid and a molecular solvent):

∆solvG ) RT ln

() KH

(6)

p0

where p0 is the standard state pressure. From the temperature variation of KH, other thermodynamic properties of solvation can be calculated, which in the case of gaseous solutes at low pressures can be approximated to the thermodynamic properties of solution. The differences in enthalpy and entropy between the two standard states referred above (∆solvH and ∆solvS, respectively) can be obtained by taking the corresponding partial derivatives of the Gibbs energy with respect to temperature:

[ ( )] [ ( )] ( )

∆solvH ) -RT2 ∆solvS ) -RT

KH ∂ ln 0 ∂T p

KH ∂ ln 0 ∂T p

- R ln

KH p0

(7)

Determination of the Molar Volumes. The molar volume of [C2mim][NTf2] was calculated as a function of temperature from density measurements performed at atmospheric pressure between 293 and 333 K at 10 K intervals. The molar volume of the [C2mim][NTf2] + CH3CN mixtures was determined from density measurements as a function of composition at 303 and 323 K. All the density measurements were performed using a U-shape vibrating tube densimeter (Anton Paar, model DMA 512 P) operating in a static mode. Temperature was maintained to within (0.01 K by means of a circulating bath equipped with a PID temperature controller (Julabo model FP40-HP) and was measured using a 100 Ω resistance platinum thermometer calibrated against a secondary reference standard with a precision of (0.02 K and an accuracy of (0.04 K. The calibration of the densimeter was made, in the present study, by measuring the period of vibration of the U-tube between 293 and 333 K filled with air, triply distilled water, and two aqueous solutions of sodium chloride with accurately known concentrations of approximately 1 and 3 M. From the density of the mixtures as a function of composition and that of the pure components, it is possible to calculate the excess volume, VEm, of the mixtures by

VEm ) Vmix m -

∑i

Mi xi

Fi

(8)

where Vmix m is the molar volume of the mixture and xi is the mole fraction of component i with molar mass Mi and density Fi. Measurement of the Vapor Pressure. The vapor liquid equilibria of the solvent mixtures were determined as a function of temperature and of the composition.

Ind. Eng. Chem. Res., Vol. 45, No. 24, 2006 8183 Table 1. Densities, Gm, and Excess Molar Volumes, VEm, for [C2mim][NTf2] + CH3CN Binary Mixtures as a Function of the Composition, Expressed in the Mole Fraction of CH3CN, x3, at Two Temperatures and at Atmospheric Pressure Fm VEm (kg m-3) (10-6 m3 mol-1)

x3 0.0000 0.1024 0.2554 0.4878 0.5514 0.7615 0.9441 1.0000

T ) 303.02 K 1513.7 0.000 1498.2 -0.242 1467.9 -0.444 1397.3 -0.597 1369.2 -0.578 1224.6 -0.453 939.1 -0.139 771.28 0.000

x3 0.0000 0.1196 0.2243 0.4481 0.5104 0.5850 0.8367 0.9354 1.0000

Fm VEm (kg m-3) (10-6 m3 mol-1) T ) 322.41 K 1494.2 0.000 1476.3 -0.432 1456.5 -0.691 1394.2 -0.961 1369.7 -0.985 1333.9 -0.952 1116.5 -0.572 937.95 -0.230 750.00 0.000

The equilibrium cell described above (represented in Figure 2) and used for the solubility measurements was also employed to measure the vapor pressure of the [C2mim][NTf2] + CH3CN mixtures following a static method. The liquid mixtures are prepared, with different total compositions, in the way described for the solubility measurements and are equilibrated with their vapor, at constant volume, at distinct temperatures. The pressure is measured by the manometer (M) that is kept at a temperature slightly above that of the liquid thermostat (TB) in order to avoid the condensation. It was proved that this small difference in temperature does not affect significantly the vapor pressures measured. 3. Results and Discussion The densities of the degassed and dried [C2mim][NTf2] measured at the different temperatures were fitted to the function (the average absolute deviation of the fit is 0.009%):

F1(kg m-3) ) 1543.55 - 1.00094(T(K) - 273.15) (9) Several authors have also measured the density of this ionic liquid.25,28,30 It is the data from Krummen et al.15 that fit best our values with a maximum relative deviation of 0.05% at 333 K. The densities measured by Tokuda et al.29 and Fredlake et al.30 are systematically higher than ours with a maximum deviation of 0.2% which is within their claimed experimental error. The density of pure acetonitrile was measured at 303.02 K (771.23 kg m-3) and at 322.41 K (749.94 kg m-3). Our data were compared with experimental values presented in the literature as the density of CH3CN was determined, as a function of temperature, by several authors,31-33 using pycnometers. The mean relative deviation between the linear fit of the literature data and our experimental values is of 0.008%. This value can be regarded as an estimation of the accuracy of our data. Because the reported density of acetonitrile covers a larger temperature range (data at four temperatures between 293 and 323 K), we have decided to consider the data obtained by Paez and Contreras,33 adjusted to a linear function of the temperature:

F3(kg m-3) ) 804.03 - 1.098(T(K) - 273.15)

(10)

The densities of the [C2mim][NTf2] + CH3CN mixtures were determined as a function of composition at atmospheric pressure and at two temperatures: 303.02 and 322.41 K. The values measured are presented in Table 1. By using the densities of the pure components and that of the mixture at different compositions, it was possible to calculate the excess molar volumes, VEm, at the two temperatures. These values are also

Figure 3. Excess molar volumes for [C2mim][NTf2] + CH3CN mixtures as a function of CH3CN mole fraction composition, x3: (b) T ) 303.02; (9) T ) 322.41 K. The lines represent the Redlich-Kister-type fittings with the parameters indicated in the text.

listed in Table 1 and are represented in Figure 3. The values were correlated, at each temperature, as a function of composition by a Redlich-Kister-type equation:34

VEm ) x1x2A0 + A1(x1 - x2) + A2(x1 - x2)2 + ‚‚‚

(11)

At the lower temperature, the coefficients obtained in the fit are A0 ) -2.3944 × 10-6 m3 mol-1, A1 ) +6.50 × 10-8 m3 mol-1 with an average absolute deviation (AAD) ) 3.1%. At 322.41 K, the coefficients obtained are A0 ) -3.9403 × 10-6 m3 mol-1, A1 ) +3.25 × 10-8 m3 mol-1 with an AAD ) 2.0%. The excess molar volumes for this mixture are negative and decrease with temperature. Minimum values for VEm were calculated for a quasi-equimolar composition and are equal to -0.6 × 10-6 m3 mol-1 at 303.02 K and -1.0 × 10-6 m3 mol-1 at 322.41 K. Negative excess molar volumes were also calculated by Zafarani-Moattar and Shekaari.35 for the system [C4mim][PF6] + CH3CN studied from 298 to 318 K and by Wang et al.36 for the system [C4mim][BF4] + CH3CN studied at 298 K. In both cases, minimum values of the excess molar volumes were obtained for mole fractions of CH3CN close to 0.7, the values being more negative than those obtained in this work: -1.296 cm3 mol-1 at 298 K and -1.546 cm3 mol-1 at 318 K for [C4mim][PF6] + CH3CN and -1.122 cm3 mol-1 for [C4mim][BF4] + CH3CN. A more detailed analysis of the data seems to indicate that the [PF6]- anion leads to more negative excess volumes when mixed with CH3CN than for the [BF4]-; this observation has already been reported by Wang et al.36,37 By analyzing the published data on the volumetric behavior of mixtures of ionic liquids with organic solvents, it is observed that for [C4mim][PF6] + acetone, butanone, pentanone, cyclopentanone, ethyl acetate,37 or methanol;35 for 4-methyl-Nbutylpyridinium tetrafluoroborate + methanol;38 or for [C4mim][BF4] + dichloromethane, butanone, or dimethylformamide,36 the excess molar volumes are negative and always larger in absolute value when the anion of the ionic liquid is the hexafluorophosphate; on average, the cited values are 30% more negative for [C4mim][PF6]. In comparison, binary systems composed of an ionic liquid and water, for example, [C4mim][BF4] + water39,40 and [C2mim][BF4] + water,41 exhibit positive excess molar volumes. These observations suggest that water weakens the strong electrostatic attractions between the ion pairs,39 thus reducing in this way the overall cohesive energy of the system.41 This

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Ind. Eng. Chem. Res., Vol. 45, No. 24, 2006 Table 2. Vapor Pressures of [C2mim][NTf2] + CH3CN Binary Mixtures, psat solv, as a Function of the Composition, Expressed in the Mole Fraction of CH3CN, x3, at Temperatures between 293 and 334 K x3

2 psat solv (10 Pa)

2 psat solv (10 Pa)

x3

T ) 293.39 K

T ) 303.48 K

0.0000 0.4603 0.5338 0.6690 0.7792 1.0000

Figure 4. Comparison of the cation-anion radial distribution functions for [C2mim][NTf2] at 303 K in the absence and presence of 20 mol % acetonitrile derived from the EPSR model. Each radial distribution function is calculated from the center of the imidazolium ring and from the nitrogen atom of the [NTf2]- anion.

conclusion is reinforced by the observation of a lowering in the viscosity on the mixtures of ionic liquids with water. Furthermore, it has been reported, both experimentally and by molecular simulations, that water binds strongly to the anion, preferentially, and can lead to deactivation of the water with respect to reactions.42-46 In the case of the organic solvents, two major factors can be used to explain the negative excess molar volumes observed. First, the large difference in the molar volume of the two components which favors the packing of the molecular solvent in the natural cavities which exist in the ionic liquid will lead to a negative VEm. Second, the ion-dipole interactions in the mixture seem to dominate over both the dipolar order in the pure molecular solvent and the electrostatic interactions in the pure ionic liquid. This last argument is confirmed by the study of Wang et al. in which molecular solvents of different polarities were studied with [C4mim][PF6].37 Therein, the reduction in the viscosity can be explained by the relative importance of the ion-dipole interactions that are weaker than the ion-ion forces in the pure ionic liquid. The arguments of this discussion are adequate to explain the observed VEm for the [C2mim][NTf2] + CH3CN mixtures determined in the present work. The less negative values presented here are not compatible with the smaller difference in molar volume between the molecular solvent and the ionic liquid. Therefore, the difference between the [NTf2]- based ionic liquid and those based on [PF6]- or [BF4]- is more likely to be due to a decrease in the relative importance of the ion-dipole interactions compared with the ion-ion and dipole-dipole interactions. Preliminary neutron diffraction studies47 show no evidence of the formation of micelles for mole fraction concentrations of 0.2 of acetonitrile in the [C2mim][NTf2], a result explained by the relatively weak dipole-dipole interactions. The structure is dominated by the ion-ion interactions at this concentration, and little change in the ionic liquid structure is found when adding the acetonitrile. This is clearly seen by comparing the partial radial distributions of the anion around the cation with and without acetonitrile (Figure 4). In addition, in this structure, no solute-solute interactions were found. This fact is consistent with X-ray diffraction studies of tetraalkyl phosphonium ionic liquids where the salt structure was practically unchanged on dissolution of high concentrations of acetonitrile, the structure being considered as a liquid clathrate.48 Similar behavior has also been observed for the dissolution of benzene in 1,3-dimethylimidazolium hexafluorophosphate.49 The vapor pressure of the [C2mim][NTf2] + CH3CN mixtures

0 29.6 40.4 49.0 67.2 94.7 T ) 313.48 K 0.000 0 0.4596 69.4 0.5330 90.8 0.6684 117.8 0.7789 157.9 1.0000 230.6

0.0000 0.4600 0.5335 0.6688 0.7791 1.0000

0 45.4 61.2 78.4 104.0 150.5 T ) 323.46 K 0.0000 0 0.4590 103.7 0.5324 127.7 0.6680 172.3 0.7786 234.6 1.0000 342.8

T ) 333.95 K 0.0000 0.4582 0.6673 0.7783 1.0000

0 149.3 258.7 339.7 505.6

Table 3. Parameters A, B, and C of Antoine Equation 12 Used To Fit the Experimental Vapor Pressures from Table 2 as a Function of the Temperature at the Studied Compositions along with the Percent Average Absolute Deviation of the Fit (AAD)a x3

A

B (K)

C (K)

AAD %

0.4594 0.5332 0.6683 0.7788 1.0000

21.44 20.69 19.93 27.86 7.8262

-3952 -3635 -2757 -8432 -3268.53

0.0162 -0.0184 -51.82 150.3 -31.615

0.6 0.6 0.3 0.8

a The values for pure acetonitrile (x ) 1) are the ones by Ewing and 3 Sanchez Ochoa.50

was measured as a function of temperature and composition. The experimental data are presented in Table 2. At each composition studied, the vapor pressure was correlated as a function of temperature by an Antoine equation:

ln[psolv sat (Pa)] ) A +

B T(K) + C

(12)

The parameters A, B, and C are listed in Table 3 which also includes the data for pure acetonitrile (x3 ) 1) from the work reported by Ewing and Sanchez Ochoa.50 The vapor pressure results are represented in Figure 5: in the upper plot, the vapor pressure of the liquid mixture is represented for the different compositions measured, and in the lower plot, the same data are depicted but as a function of the mole fraction composition at each one of the temperatures studied (the dashed lines represent the values given by Raoult’s law). No immiscibility gap was found, which confirms that acetonitrile is miscible with the ionic liquid in the entire composition range. The mixture shows a strong nonideality with negative deviations from Raoult’s law at all temperatures and all compositions. The largest deviation from ideality was found for the lowest mole fraction of acetonitrile (x3 ) 0.46) and the highest temperature studied (36% deviation from Raoult’s law at T ) 333.95 K). The saturation behavior of the [C2mim][NTf2] + CH3CN mixtures can be explained, in the same manner as for the VEm, in terms of the interactions between molecules in the pure liquids and in the mixture.51 In the present case, the observed negative

Ind. Eng. Chem. Res., Vol. 45, No. 24, 2006 8185 Table 4. Experimental Values of Gas Solubilities in Several [C2mim][NTf2] + CH3CN Binary Mixtures Expressed Both as Henry’s Law Constants, KH, and as CO2 Mole Fraction, x2, Corrected for a Partial Pressure of Solute of 0.1 MPaa peq (102 Pa)

KH (105Pa)

x2 (10-2)

Dev (%)

303.63 303.87 313.87 313.92 313.95 323.96 323.98 324.05 324.12 334.07 334.10 334.16 343.99 344.23

479.2 464.1 495.9 492.9 500.6 519.6 520.4 433.0 521.6 542.2 546.1 448.4 563.7 571.0

x3 ) 0.0000 38.08 38.61 45.85 46.22 45.43 54.15 54.64 53.92 54.47 62.93 63.43 63.90 72.51 72.90

2.614 2.578 2.172 2.155 2.192 1.840 1.823 1.848 1.829 1.584 1.571 1.560 1.375 1.368

-0.3 +0.6 -0.1 +0.6 -1.2 -0.1 +0.7 -0.7 +0.2 -0.5 +0.3 +0.9 -0.3 -0.1

293.76 303.71 313.72 324.01 334.07

534.8 579.0 635.4 696.7 765.6

x3 ) 0.4523 42.00 49.81 62.57 73.79 83.60

2.368 1.998 1.591 1.350 1.192

+1.1 -2.9 +1.4 +1.1 -0.9

293.53 303.62 313.59 323.85

500.0 548.7 601.9 662.7

x3 ) 0.5261 46.49 58.83 70.73 85.54

2.140 1.692 1.408 1.165

-0.2 +0.7 -0.7 +0.2

293.76 303.15 313.27 334.38

526.9 587.5 652.9 841.4

x3 ) 0.6610 49.36 61.99 73.11 96.66

2.015 1.605 1.362 1.031

-0.4 +1.1 -0.9 +0.2

292.66 303.03 312.79 322.86 333.37

493.7 556.7 627.4 721.1 847.8

x3 ) 0.7730 64.27 76.98 83.79 91.09 104.15

1.548 1.293 1.188 1.094 0.9568

-1.2 +2.7 -0.4 -2.5 +1.4

T (K)

Figure 5. Vapor pressure of the [C2mim][NTf2] + CH3CN mixtures. (upper plot) Representation as a function of temperature at different CH3CN mole fraction compositions, x3: ([) x3 ) 0.4594; (9) x3 ) 0.5332; (1) x3 ) 0.6683; (2) x3 ) 0.7788. (lower plot) Representation as a function of the CH3CN mole fraction composition x3: ([) T ) 293.39; (9) T ) 303.48; (1) T ) 313.48; (2) T ) 323.46; (b) T ) 333.95 K. The dashed lines represent Raoult’s law at each temperature.

deviations from Raoult’s law are probably due to the increasing importance of the ion-dipole interactions, compared with the ion-ion and dipole-dipole interactions present in the two pure liquids, and it cannot be discounted that a stronger chargetransfer effect may also exist. This is in agreement with structural data which show that acetonitrile is dispersed throughout the ionic liquid and few acetonitrile-acetonitrile interactions are found in the mixture. The activity coefficients at infinite dilution of binary systems involving acetonitrile and one ionic liquid have been studied by other authors,26,52 namely, mixtures of CH3CN + [C2mmim][NTf2],26 CH3CN + [C2mim][NTf2],26 and CH3CN + [C4mim][NTf2].52 The values measured are low and so compatible with the high affinity of the two species that they lead to the volumetric and saturation properties of the present study. The solubility of carbon dioxide was measured in the pure ionic liquid [C2mim][NTf2] and in four mixtures of [C2mim][NTf2] + CH3CN of different compositions as a function of temperature between 283 and 343 K in steps of approximately 10 K. The experimental data are reported in Table 4 in terms of mole fraction and a Henry’s law constant calculated by eq 5. The experimental mole fraction solubilities are calculated by eq 1 from the values obtained by eqs 2 and 4 assuming the validity of Raoult’s law for the solute-solvent mixture. The mole fraction solubilities reported in Table 4 are calculated from KH determined from the experimental values (measured at

a p is the equilibrium pressure, and the percent deviation is relative to eq the correlations of the data reported in Table 5.

slightly different equilibrium pressures) assuming a partial pressure of the gaseous solute equal to 0.1 MPa. The relative atomic masses used are those recommended by IUPAC;53 the values of the second virial coefficient of carbon dioxide were taken from the compilation by Dymond and Smith,54 and values for acetonitrile were from the work of Cholinsky et al.55 In each case, the density of the gas-liquid solution in equilibrium with its vapor is considered to be equal to that of the liquid solvent at the same temperature (at atmospheric pressure) and is given, for the pure ionic liquid, by eq 9 and, for the mixed liquid solvent at each composition in acetonitrile, by the correlation of the values Fm reported in Table 1. The values for the solubility of carbon dioxide, expressed as Henry’s law constants, were correlated as a function of temperature by an empirical equation of the type: n

ln[KH(105 Pa)] )

Ai(T(K))-i ∑ i)0

(13)

The coefficients Ai obtained are listed in Table 5 together with the average absolute deviations obtained. These values can be regarded as an estimation of the precision of the experimental data for gas solubility, which is in the present case better than 2%.

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Table 5. Parameters of Equation 13 Used To Smooth the Experimental Results on KH from Table 4 along with the Percent Average Absolute Deviation of the Fit (AAD) x3

A0

A1

A2 (104)

AAD %

0.0000 0.4523 0.5261 0.6610 0.7730 1.0000

+7.595 +4.177 +4.949 -3.895 +4.053 +10.86

-671.6 +1677 +1410 +6733 +1324 -2870

-16.04 -53.15 -50.92 -130.5 -37.71 +28.33

0.5 1.5 0.5 0.6 1.6 2.9

Figure 6. Henry’s law constants for CO2 in the [C2mim][NTf2] + CH3CN mixtures. (upper plot) Representation as a function of temperature at different CH3CN mole fraction compositions of the solvent, x3: (O) x3 ) 0 (pure ionic liquid); ([) x3 ) 0.4594; (9) x3 ) 0.5332; (1) x3 ) 0.6683; (2) x3 ) 0.7788. The upper line represents the extrapolated values for x3 ) 1 (pure acetonitrile) as explained in the text. (lower plot) Representation as a function of the CH3CN mole fraction composition x3: ([) T ) 293.39; (9) T ) 303.48; (1) T ) 313.48; (2) T ) 323.46; (b) T ) 333.95 K. The dashed lines correspond to the linear fit of the data at the different temperatures.

The solubility data, in the form of Henry’s law constants, for carbon dioxide in the different liquid solvents is plotted in the upper plot of Figure 6 as a function of temperature. In the lower plot of Figure 6, the same Henry’s law constants are represented as a function of acetonitrile composition in the liquid solvent. An estimate of the solubility of carbon dioxide in pure acetonitrile at the different temperatures studied can be obtained by extrapolating these values for x3 ) 1. This extrapolation is represented in the lower plot of Figure 6, and the values obtained at the different temperatures were correlated by eq 13, the coefficients obtained being also included in Table 5. Only a linear extrapolation, to obtain the estimation of the CO2 solubility in CH3CN, was compatible with the number and precision of the data point available. The linear fits, represented

by the dashed lines in the lower plot of Figure 6, are characterized by average absolute deviations below 1.5%, meaning that the solubility of carbon dioxide in acetonitrile, calculated in this way, can surely be considered to be precise to within 5%. The values were compared with the literature for the solubility of carbon dioxide in pure acetonitrile,56-58 and it was observed that the data are systematically above those reported herein, especially in the case of the extrapolation of the high-pressure data reported by Kordikowski et al.57 and Lazzaroni et al.58 probably due to the presence, in their case, of water dissolved in the acetonitrile sample. Clearly, the solubility of carbon dioxide decreases with the addition of acetonitrile; the solubility does seem to vary monotonically with the mole fraction of the molecular liquid but not in a similar manner for all the temperatures studied. In all the cases, the solubility decreases with temperature. The solubility of carbon dioxide in the pure ionic liquid using the modified apparatus showed good agreement with values measured in this laboratory using a different isochoric apparatus and another batch of [C2mim][NTf2].11 The deviations encountered are of the order of 0.6%, a result that lies within the precision and accuracy claimed for the measurements. Furthermore, these data have a maximum deviation of 5% when compared with other data reported for the solubility of CO2 in [C2mim][NTf2] as a function of temperature.9,59 From all these observations, it is possible to conclude that the extension of the isochoric saturation technique24,25,60 to measure the solubility of gases in mixtures of liquids is reliable. From the variation with temperature of the Henry’s law constants, it is possible to calculate the derived thermodynamic properties of solvation as explained above. For all the systems studied, an increase of the Gibbs energy of solvation (varying from 9 to 13 kJ mol-1) with temperature is observed corresponding to the increase in the Henry’s law constants. Negative enthalpies of solvation were calculated (between -8 and -16 kJ mol-1) corresponding to exothermic processes of solvation. Finally, strongly negative values (between -60 and -90 J K-1 mol-1) were obtained for the entropy of solvation, a situation also encountered for the solvation of carbon dioxide in other ionic liquids based in the imidazolium cation and the NTf2anion. 4. Conclusions This work permits the quantitative evaluation of the effect of acetonitrile, a common strongly polar molecular organic solvent, on the solubility of carbon dioxide in a widely used ionic liquid, the 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide, [C2mim][NTf2]. It is observed that the solubility decreases with the mole fraction concentration of acetonitrile in the liquid solvent from x2 ) 2.60 × 10-2 in the pure ionic liquid to x2 ) 1.29 × 10-2 in the solvent with 0.77 mole fraction concentration of acetonitrile at 303 K, a 50% reduction on the gas solubility. For all the systems, the carbon dioxide solubility decreases with temperature in the range studied from 290 to 335 K. This observation is compatible with exothermic processes of solvation, and the calculated thermodynamic properties confirm this fact quantitatively. The study of the saturation and volumetric properties of the ionic liquid + acetonitrile mixture was necessary to calculate the gas solubilities from the experimental data accurately. The vapor pressure of the different mixtures was measured as a function of temperature and strong negative deviations from Raoult’s law were obtained (up to 36% for a 0.46 mole fraction of CH3CN at the higher temperature studied). Negative excess

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molar volumes, with a maximum of approximately -1 cm3 mol-1 at an equimolar composition, were also measured at different temperatures. These two observations are compatible with the dominance of strong interactions between [C2mim][NTf2] and CH3CN probably of the type ion-dipole. These interactions are most probably more favorable than the dipole order in the molecular liquid. Acknowledgment M.D. thanks QUILL, the Royal Society, and EPSRC for funding through the Portfolio Partnership. O.S. acknowledges support from the Department of Education and Learning in Northern Ireland and Merck GmBH. CCLRC are thanked for the allocation of the beamtime. C.H., P.H., and M.F.C.G. thank the Royal Society and the Centre National de la Recherche Scientifique for funding. Literature Cited (1) Welton, T. Room-Temperature Ionic Liquids. Solvents for Synthesis and Catalysis. Chem. ReV. 1999, 99, 2071-2083. (2) Wasserscheid, P.; Keim, W. Ionic Liquids - New ×e3Solutions“ for Transition Metal Catalysis. Angew. Chem., Int. Ed. 2000, 39, 37723789. (3) Wilkes, J. S. Properties of ionic liquid solvents for catalysis. J. Mol. Catal. A, Chem. 2004, 214, 11-17. (4) Zhao, D. B.; Wu, M.; Kou, Y.; Min, E. Ionic liquids: applications in catalysis. Catal. Today 2002, 74, 157-189. (5) Ionic Liquids in Synthesis; Wasserscheid, P., Welton, T., Eds.; WileyVCH: Weinheim, Germany, 2003. (6) Costa Gomes, M. F.; Padua, A. A. H. Gas-liquid interactions in solution. Pure Appl. Chem. 2005, 77, 653-665. (7) Scovazzo, P.; Camper, D.; Kieft, J.; Poshusta, J.; Koval, C.; Noble, R. Regular solution theory and CO2 gas solubility in room-temperature ionic liquids. Ind. Eng. Chem. Res. 2004, 43, 6855-6860. (8) Baltus, R. E.; Culbertson, B. H.; Dai, S.; Luo, H.; DePaoli, D. W. Low-pressure solubility of carbon dioxide in room-temperature ionic liquids measured with a Quartz crystal microbalance. J. Phys. Chem. B 2004, 108, 721-727. (9) Camper, D.; Scovazzo, P.; Koval, C.; Noble, R. Gas solubilities in room-temperature ionic liquids. Ind. Eng. Chem. Res. 2004, 43, 30493054. (10) Anthony, J. L.; Anderson, J. L.; Maginn, E. J.; Brennecke, J. F. Anion effects on gas solubility in Ionic Liquids. J. Phys. Chem. B 2005, 109, 6366-6374. (11) Jacquemin, J.; Husson, P.; Majer, V.; Costa Gomes, M. F. Influence of the cation on the solubility of CO2 and H2 in ionic liquids based on the bis(trifluoromethylsulfonyl) imide anion. J. Sol. Chem., submitted for publication. (12) Kato, R.; Krummen, M.; Gmehling, J. Measurement and correlation of vapor-liquid equilibria and excess enthalpies of binary systems containing ionic liquids and hydrocarbons. Fluid Phase Equilib. 2004, 224, 47-54. (13) Verevkin, S. P.; Vasiltsova, T. V.; Bich, E.; Heintz, A. Thermodynamic properties of mixtures containing ionic liquids Activity coefficients of aldehydes and ketones in 1-methyl-3-ethyl-imidazolium bis(trifluoromethyl-sulfonyl)imide using the transpiration method. Fluid Phase Equilib. 2004, 218, 165-175. (14) Kato, R.; Gmehling, J. Measurement and correlation of vapor-liquid equilibria of binary systems containing the ionic liquids [EMIM][(CF3SO2)2N], [BMIM][(CF3SO2)2N], [MMIM][(CF3)2PO4] and oxygenated organic compounds respectively water. Fluid Phase Equilib. 2005, 231, 38-43. (15) Krummen, M.; Wasserscheid, P.; Gmehling, J. Measurement of activity coefficients at infinite dilution in ionic liquids using the dilutor technique. J. Chem. Eng. Data 2002, 47, 1411-1417. (16) Heintz, A.; Kulikov, D. V.; Verevkin, S. P. Thermodynamic properties of mixtures containing ionic liquids. 2. Activity coefficients at infinite dilution of hydrocarbons and polar solutes in 1-methyl-3-ethylimidazolium bis(trifluoromethyl-sulfonyl) amide and in 1,2-dimethyl-3-ethylimidazolium bis(trifluoromethyl-sulfonyl) amide using gas-liquid chromatography. J. Chem. Eng. Data 2002, 47, 894-899. (17) Hert, D. G.; Anderson, J. L.; Aki, S. N. V. K.; Brennecke, J. F. Enhancement of oxygen and methane solubility in 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl) imide using carbon dioxide. Chem. Commun. 2005, 2603-2605.

(18) Doker, M.; Gmehling, J. Measurement and prediction of vaporliquid equilibria of ternary systems containing ionic liquids. Fluid Phase Equilib. 2005, 227, 255-266. (19) Blanchard, L. A.; Gu, Z.; Brennecke, J. F. High-pressure phase behaviour of ionic liquid/CO2 systems. J. Phys. Chem. B 2001, 105, 24372444. (20) Fu, D.; Sun, X.; Pu, J.; Zhao, S. Effect of water content on the solubility of CO2 in the ionic liquid [bmim][PF6]. J. Chem. Eng. Data 2006, 51, 371-375. (21) Aki, S. N. V. K.; Mellein, B. R.; Saurer, E. M.; Brennecke, J. F. High-pressure phase behaviour of carbon dioxide with imidazolium-based ionic liquids. J. Phys. Chem. B 2004, 108, 20355-20365. (22) Bonhoˆte, P.; Dias, A. P.; Papageorgiou, N.; Kalyanasundaram, K.; Gratzel, M. Hydrophobic, highly conductive ambient-temperature molten salts. Inorg. Chem. 1996, 35, 1168-1178. (23) Villagra´n, C.; Deetlefs, M.; Pitner, W. R.; Hardacre, C. Quantification of halide in ionic liquids using ion chromatography. Anal. Chem. 1996, 68, 2118-2123. (24) Jacquemin, J.; Husson, P.; Majer, V.; Costa Gomes, M. F. Lowpressure solubilities and thermodynamics of solvation of eight gases in 1-butyl-3-methylimidazolium hexafluorophosphate. Fluid Phase Equilib. 2006, 240, 87-95. (25) Jacquemin, J.; Costa Gomes, M. F.; Husson, P.; Majer, V. Solubility of carbon dioxide, ethane, methane, oxygen, nitrogen, hydrogen, argon and carbon monoxide in 1-butyl-3-methylimidazolium tetrafluoroborate between temperatures 283 and 343 K and at pressures close to atmospheric. J. Chem. Thermodyn. 2006, 38, 490-502. (26) Benson, B. B.; Krause, D., Jr. A thermodynamic treatment of dilute solutions of gases in liquids. J. Solution Chem. 1989, 18, 803-821. (27) Costa Gomes, M. F.; Grolier, J.-P. Determination of Henry’s law constants for aqueous solutions of tetradeuteriomethane between 285 and 325 K and calculation of the H/D isotope effect. Phys. Chem. Chem. Phys. 2001, 3, 1047-1052. (28) Prausnitz, J. M.; Lichtenthaler, R. N.; Gomes de Azevedo, E. Molecular thermodynamics of fluid-phase equilibria, 3rd ed.; Prentice Hall: New Jersey, 1999. (29) Tokuda, H.; Hayamizu, K.; Ishii, K.; Susan, M. A. B. H.; Watanabe, M. Physicochemical properties and structures of room-temperature ionic liquids. 2. Variation of alkyl chain length in imidazolium cation. J. Phys. Chem. B 2005, 109, 6103-6110. (30) Fredlake, C. P.; Crosthwaite, J. M.; Hert, D. G.; Aki, S. N. V. K.; Brennecke, J. F. Thermophysical properties of imidazolium based ionic liquids. J. Chem. Eng. Data 2004, 49, 954-964. (31) Ritzoulis, G.; Papadopoulos, N.; Jannakoudakis, D. Densities, viscosities, and dielectric constants of acetonitrile + toluene at 15, 25, and 35 °C. J. Chem. Eng. Data 1986, 31, 146-148. (32) Nikam, P. S.; Shirsat, L. N.; Hasan, M. Density and viscosity studied of binary mixtures of acetonitrile with methanol, ethanol, propan-1-ol, propan-2-ol, butan-1-ol, 2-methylpropan-1-ol, and 2-methylpropan-2-ol at (298.15, 303.15, 308.15, and 313.15) K. J. Chem. Eng. Data 1998, 43, 732-737. (33) Paez, S.; Contreras, M. Densities and viscosities of binary mixtures of 1-propanol and 2-propanol with acetonitrile. J. Chem. Eng. Data 1989, 34, 455-459. (34) Rowlinson, J. S.; Swinton, F. L. Liquids and Liquid Mixtures; 3rd ed.; Butterworth: London, 1982. (35) Zafarani-Moattar, M. T.; Shekaari, H. Volumetric and speed of sound of ionic liquid, 1-butyl-3-methylimidazolium hexafluorophosphate with acetonitrile and methanol at T ) (298.15 to 318.15) K. J. Chem. Eng. Data 2005, 50 (5), 1694-1699. (36) Wang, J.; Tian, Y.; Zhao, Y.; Zhuo, K. A volumetric and viscosity study for the mixtures of 1-n-butyl-3-methylimidazolium tetrafluoroborate ionic liquid with acetonitrile, dichloromethane, 2-butanone and N, Ndimethylformamide. Green Chem. 2003, 5, 618-622. (37) Wang, J.; Zhu, A.; Zhao, Y.; Zhuo, K. Excess molar volumes and excess logarithm viscosities for binary mixtures of the ionic liquid 1-butyl3-methylimidazolium hexafluorophosphate with some organic compounds. J. Solution Chem. 2005, 34, 585-596. (38) Heintz, A.; Klasen, D.; Lehmann, J. K. Excess molar volumes and viscosities of binary mixtures of methanol and the ionic liquid 4-methylN-butylpyridinium tetrafluoroborate at 25, 40, and 50 °C. J. Solution Chem. 2002, 31, 467-476. (39) Rebelo, L. P. N.; Najdanovic-Visak, V.; Visak, Z. P.; Nunes da Ponte, M.; Szydlowski, J.; Cerdeirina, C. A.; Troncoso, J.; Romani, L.; Esperanc¸ a, J. M. S. S.; Guedes, H. J. R.; de Sousa, H. C. A detailed thermodynamic analysis of [C4mim][BF4] + water as a case study to model ionic liquid aqueous solutions. Green Chem. 2004, 6, 369-381.

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(40) Seddon, K. R.; Stark, A.; Torres, M. J. Influence of chloride, water, and organic solvents on the physical properties of ionic liquids. Pure Appl. Chem. 2000, 72, 2275-2287. (41) Zhang, S.; Li, X.; Chen, H.; Wang, J.; Zhang, J.; Zhang, M. Determination of physical properties for the binary system of 1-ethyl-3methylimidazolium tetrafluoroborate + H2O. J. Chem. Eng. Data 2004, 49, 760-764. (42) Cammarata, L.; Kazarian, S. G.; Salte, P. A.; Welton, T. Molecular states of water in room-temperature ionic liquids. Phys. Chem. Chem. Phys. 2001, 3, 5192-5200. (43) Hanke, C. G.; Lynden-Bell, R. M. A Simulation Study of WaterDialkylimidazolium Ionic Liquid Mixtures. J. Phys. Chem. B 2003, 107, 10873-10878. (44) Farmer, V.; Welton, T. The oxidation of alcohols in substituted imidazolium ionic liquids using ruthenium catalysts. Green Chem. 2002, 4, 97-102. (45) Doherty, S.; Goodrich, P.; Hardacre, C.; Luo, H. K.; Rooney, D. W.; Seddon, K. R.; Styring, P. Marked enantioselectivity enhancements for Diels-Alder reactions in ionic liquids catalysed by platinum diphosphine complexes. Green Chem. 2004, 6, 63-67. (46) Amigues, E.; Hardacre, C.; Keane, G.; Migaud, M. E.; O’Neill, M. Ionic liquids - media for unique phosphorous chemistry. Chem. Commun. 2006, 72. (47) The neutron diffraction data were taken on protiated and deuterated forms of [C2mim][NTf2] and acetonitrile at 303 K using the SANDALS diffractometer at the ISIS pulsed neutron source, Rutherford Appleton Laboratory, U.K. The data were analyzed using the standard ATLAS package to produce a differential scattering cross section for each sample. Soper, A. K.; Howells, W. S.; Hannon, A. C. ATLAS-Analysis of Timeof-Flight Diffraction Data from Liquid and Amorphous Samples; Report RAL-89-046, Rutherford Appleton Laboratory: U.K., 1989. The neutron diffraction data was fitted using the empirical potential structure refinement process, EPSR (Soper, A. K. Mol. Phys. 2001, 99, 1503-1516). The charges used for the solute and anion were Lo¨wdin charges from optimised structures at the 6-31++G*//B3LYP level of theory in GAMMES. The cation charges were those derived by Lopes and Padua (Lopes, J. N. C.; Padua, A. A. H. J. Phys. Chem. B 2004, 108, 16893-16898). The Lennard-Jones parameters for interactions between unlike atoms were calculated by employing the usual Lorentz-Berthelot mixing rules. (48) Gowda, G. A. N.; Chen, H.; Khetrapal, C. L.; Weiss, R. G. Amphotropic Ionic Liquid Crystals with Low Order Parameters. Chem. Mater. 2004, 16, 2101-2106. (49) (a) Deetlefs, M.; Hardacre, C.; Nieuwenhuyzen, M.; Sheppard, O.; Soper, A. K. Structure of ionic liquid-benzene mixtures J. Phys. Chem. B 2005, 109, 1593-1598. (b) Holbrey, J. D.; Reichert, W. M.; Nieuwenhuyzen, M.; Sheppard, O.; Hardacre, C.; Rogers, R. D. Liquid clathrate formation in ionic liquid-aromatic mixtures. Chem. Commun. 2003, 476477.

(50) Ewing, M. B.; Sanchez Ochoa, J. C. Vapor pressures of acetonitrile determined by comparative ebulliometry. J. Chem. Eng. Data 2004, 49, 486-491. (51) Murrell, J. N.; Jenkins, A. D. Properties of Liquids and Solutions, 2nd ed.; Wiley: Chischester, 1994. (52) Heintz, A.; Casas, L. M.; Nesterov, I. A.; Emel’yanenko, V. N. Thermodynamic properties of mixtures containing ionic liquids. 5. Activity coefficients at infinite dilution of hydrocarbons, alcohols, esters, and aldehydes in 1-methyl-3-butyl-imidazolium bis(trifluoromethyl-sulfonyl) imide using gas-liquid chromatography. J. Chem. Eng. Data 2005, 50, 1510-1514. (53) IUPAC Commission on Atomic Weights and Isotopic Abundances. Atomic Weights of the Elements 1985. Pure Appl. Chem. 1986, 58, 16771692. (54) Dymond, J. H.; Smith, E. B. The Virial Coefficients of Pure Gases and Mixtures; Clarendon Press: Oxford, 1980. (55) Cholinsky, J.; Szafranski, A.; Wyrzykowska-Stankiewicz, D. Computer-Aided Second Virial Coefficient Data for Organic IndiVidual Compounds and Binary Systems; PWN-Polish Scientific Publishers: Warsaw, 1986. (56) Fujita, E.; Szalda, D. J.; Creutz, C.; Sutin, N. Carbon dioxide activation: Thermodynamics of CO2 binding and the involvement of two cobalt centers in the reduction of CO2 by a cobalt(I) macrocycle. J. Am. Chem. Soc. 1998, 110, 4870-4871. (57) Kordikowski, A.; Schenk, A. P.; Van Nielen, R. M.; Peters, C. J. Volume expansions and vapor-liquid equilibria of binary mixtures of a variety of polar solvents and certain near-critical solvents. J. Supercrit. Fluids 1995, 8, 205-216. (58) Lazzaroni, M. J.; Bush, D.; Brown, J. S.; Eckert, C. A. Highpressure vapor-liquid equilibria of some carbon dioxide + organic binary systems. J. Chem. Eng. Data 2005, 50, 60-65. (59) Cadena, C.; Anthony, J. L.; Shah, J. K.; Morrow, T. I.; Brennecke, J. F.; Maginn, E. J. Why is CO2 so soluble in imidazolium-based ionic liquids? J. Am. Chem. Soc. 2004, 126, 5300-5308. (60) Bonifacio, R. P.; Costa Gomes, M. F.; Filipe, E. J. M. Solubility of xenon in n-hexane between 257 and 333 K. Fluid Phase Equilib. 2002, 193, 41-51.

ReceiVed for reView April 27, 2006 ReVised manuscript receiVed September 12, 2006 Accepted September 18, 2006 IE0605377