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Ind. Eng. Chem. Res. 2008, 47, 493-501

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Melting Point Depression of Ionic Liquids with CO2: Phase Equilibria Aaron M. Scurto,*,†,‡ Elizabeth Newton,§,| Ross R. Weikel,§,| Laura Draucker,§,| Jason Hallett,§,| Charles L. Liotta,§,|,⊥ Walter Leitner,†,# and Charles A. Eckert§,|,⊥ Institut fu¨r Technische und Makromolekulare Chemie, RWTH Aachen, Worringerweg 1, D-52074 Aachen, Germany, Department of Chemical & Petroleum Engineering and NSF-ERC Center for EnVironmentally Beneficial Catalysis, UniVersity of Kansas, Lawrence, Kansas 66045, School of Chemical and Biomolecular Engineering, Georgia Institute of Technology, Atlanta, Georgia 30332, School of Chemistry and Biochemistry, Georgia Institute of Technology, Atlanta, Georgia 30332, Specialty Separations Center, Georgia Institute of Technology, Atlanta, Georgia 30332, and Max-Planck-Institut fu¨r Kohlenforschung, Mu¨lheim an der Ruhr, Germany

Development of ionic liquids for specific tasks is currently being pursued by many researchers as numerous cation/anion combinations are theoretically possible. However, only a small fraction of these combinations melt below 100 °C. Recently, large melting point depressions of several ionic solids with compressed carbon dioxide have been reported. This investigation details the melting point depression of a large number of ionic organic compounds (ionic liquids) with gaseous, liquid, and supercritical CO2. Large and previously unreported depressions were observed for some of the ionic solids. This methodology greatly expands the numbers of compounds and functional groups that can be employed in an ionic liquid/compressed gas system for various applications. Thermodynamic analysis indicates that even small amounts of CO2 can lead to substantial melting point depression, due to its very low melting temperature and negative deviations to Raoult’s law. Introduction Room-temperature ionic liquids (ILs) have been the subject of intense focus due to their demonstrated usefulness as solvents for extractions,1 reactions,2,3 and materials processing.4,5 Preliminary toxicology studies of ILs6-8 indicate low to moderate toxicity on average, and most ILs have no measurable vapor pressure, thus eliminating air emissions; an advantage over any other organic solvent. Due to these advantageous properties, ILs have been touted as “green” solvents. In addition, biphasic systems with ILs can be utilized for advanced reactions and extractions. In particular, there has been significant research in pairing CO2 with ionic liquids.9-11 Carbon dioxide is frequently used in green processing due to its benign characteristics.12 CO2 has excellent solubility in many ILs, while ILs have little to no solubility in CO2.9 In addition, Scurto et al.13,14 have used CO2 pressure as a separation “switch”, to induce immiscibility in ionic liquid solutions with organics and water. Brown et al.15 have used supercritical CO2 to extract products of a homogeneously catalyzed reaction from an IL. Leitner and co-workers have performed a variety of reactions in CO2/IL biphasic systems where the ionic liquid immobilizes a homogeneous catalyst in a continuous process.10,16,17 Sellin et al.18,19 have also used the IL to immobilize a homogeneous catalyst in an IL/ CO2 biphasic system for hydroformylation reactions. Catalytic reactions using enzymes have also been carried out in such biphasic media.11,20,21 A detailed overview of biphasic catalysis in IL/CO2 systems is available in the literature.22,23

The present definition of an ionic “liquid” is an organic salt that has a melting point below 100 °C.24 There are a myriad of cation/anion combinations that can be molecularly engineered with specific physical or chemical properties. However, designing organic salts with particular properties and functional groups that are also liquid near room-temperature is not a trivial task. Of the estimated 1018 possible cation-anion combinations (1014 unique cations and anions),25 only a very small fraction actually melt below 100 °C. Often useful functional groups yield solid salts, such as those shown with metal-extracting ionic compounds by Visser et al.,26 chiral ionic liquids,27,28 an IL that is both solvent and catalyst,29 etc. A methodology or process to increase the range of ionic compounds for use as solvents would be highly useful. It has been long known that organic solids can undergo melting below their normal melting points (melting point depression) in the presence of a number of compressed gases.30 Kazarian et al.31 observed liquid-crystal transitions induced by CO2 for surfactant-like imidazolium salts. Recently, Scurto and Leitner32 have demonstrated that CO2 can induce unprecedented melting point depressions in some types of ionic salts. Moreover, the induced melts can serve as an advantageous platform for metal-complex catalysis. This contribution will present the results for a larger variety of ionic solids and demonstrate the effect of temperature, pressure, and composition on the melting point. The melting point with composition data will be modeled using an ideal and advanced activity-coefficient model. Phase Behavior

* To whom correspondence should be addressed. E-mail: ascurto@ ku.edu. Phone: +1 (785) 864-4947. Fax: +1 (785) 864-4967. † Institut fu¨r Technische und Makromolekulare Chemie. ‡ University of Kansas. § School of Chemical and Biomolecular Engineering, Georgia Institute of Technology. | Specialty Separations Center, Georgia Institute of Technology. ⊥ School of Chemistry and Biochemistry, Georgia Institute of Technology. # Max-Planck-Institut fu¨r Kohlenforschung.

Melting point depression is a thermodynamic condition that can be represented by two different equilibria: solid-liquid and solid-liquid-vapor (SLV). Figure 1 illustrates the phase transitions with pressure and temperature of a typical highly asymmetric system of a common organic solid with an ambient gaseous component: naphthalene and ethane or carbon dioxide, for example. In the presence of a second gaseous component, the melting curve of the pure solute can be lowered as shown

10.1021/ie070312b CCC: $40.75 © 2008 American Chemical Society Published on Web 09/28/2007

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Figure 1. Typical PT diagram of a highly asymmetric system for an organic compound and compressed gas.

by the dashed SLV curves labeled I and II. Solid-liquid-vapor (SLV) equilibria for a binary system is obtained by starting with higher molar loading of the compressed gas and maintaining temperature (1a to 2 in Figure 1) or pressure (1b to 2 in Figure 1) and varying the other variable until the very first instance that a solid, liquid, and vapor coexist in equilibria. The melting point depression (∆Tm) is thus Tm minus TSLV. For nonionic organic systems, the SLV line emanates from the triple point of the solute and terminates at the upper critical endpoint (UCEP). The SLV can proceed directly (curve I, e.g., naphthalene and ethane33) to the UCEP, or the SLV curve can have a minimum in temperature on a PT projection (curve II, e.g., naphthalene and carbon dioxide34-36). This implies that increasing pressure along certain isotherms will first melt the solid (SLV curve II) to produce vapor-liquid equilibrium and then at higher pressures solidify the liquid phase, returning to solidfluid equilibrium. However, some gaseous components do not induce much or any melting point depression and can actually increase the melting point beyond that of the pure component. This phenomenon is observed when naphthalene is exposed to helium.36 Solid-liquid-vapor equilibrium can also occur at lower temperatures near the vapor pressure of the less volatile component, as shown in Figure 1; however, this transition is primarily the “boiling point elevation” of the liquefied gas in the presence of the solute. At lower gas loading (moles gas: solute) and pressures above the SLV pressure, the melting point depression will be represented by solid-liquid equilibrium (SLE). The inset in Figure 1 illustrates in a pressurecomposition plot for the phase behavior regions of both SLV and SLE (P2 represents point 2 (SLV) in the larger plot). The melting point of naphthalene (Tm ) 80.1 °C) has been determined in the presence of a variety of gases: methane,37 ethane,33 ethylene,38 carbon dioxide,35 and xenon.39 The largest melting point depression was found in the presence of supercritical xenon (∆Tm ) 37.5 °C at 124 bar). Large melting point depressions with CO2 and inorganic minerals have also been discovered, however at conditions in excess of 1000 K and 1000 bar.40 CO2 has been known for several decades to lower the glass transition temperature of polymers.41 Only more recently has CO2 been shown to lower the melting temperature of polymers. Zhang and Handa42 report that CO2 lowers the melting point of syndiotactic polystyrene (Tm ) 277 °C) by approximately 12 °C at 72 bar. Kishimoto and Ishii43 found that CO2 depressed the melting point of isotactic polypropene (Tm ) 180 °C) by 11 °C at 94 bar. Weidner and co-workers44 determined the melting point depression of polyethylene glycol (PEG, MW ∼ 1500, Tm ) 44.9 °C) of approximately 16 °C at 150 bar.

In attempt to perform electrochemistry in wet supercritical carbon dioxide, Wightman and co-workers45-47 used the ionic solids tetrahexylammonium nitrate and hexafluorophosphate ([THexAm][PF6]) to aid in electrochemical oxidation. They found that the CO2/H2O mixture induced a melting point in the ionic solid (Tm ) 131 °C of [THexAm][PF6]) to 40 to 50 °C at about 90 bar. This is not surprising considering the highly hygroscopic nature of most quaternary ammonium salts. Kazarian et al.31 have used attenuated total reflection-infrared (ATR-IR) to detect the formation of a liquid crystal phase transition under 70 bar of pure CO2 of a long-chain (C16) methylimidazolium PF6; at ambient pressure, the ionic solid forms a smectic A phase at 75 °C and an isotropic liquid phase at 125 °C. They also inferred from the IR spectra that the CO2 has preferential interactions with the fluorinated anion in the liquid crystal phase. However, the general trend or magnitude of the melting point depression with compressed gases for the vast majority of ionic solids has not been elucidated. Thermodynamics and Modeling Thermodynamic modeling of solid-liquid equilibria gives the relationship for the mole fraction of the solid (2) in terms of the melting point (Tm), the triple point temperature (Tt), the enthalpy of fusion (∆hfus), the enthalpy of any solid-solid transition (∆htrans), the temperature of that transition (Ttrans), and the activity coefficient γ2. Neglecting second-order corrections, the solid-liquid equilibrium is given by eq 148

ln

1 ∆hfus (Tm - T) ∆htrans (Ttrans - T) + - ln γ2 (1) ) x2 R Tt T R TtransT

This relationship can be used to make predictions of the melting point curve, where the liquid may be either ideal or nonideal, modeled by an expression for Gibbs excess free energy/activity (gEX/γ). Deviations from Raoult’s law may be both positive or negative, but unless there is a reasonably strong compound formed (e.g., liquid metals), this procedure will give a single eutectic and the predictions of the melting point curve are generally quite accurate.48-51 For ideal solutions, the activity coefficient, γi, is set equal to one. Modeling procedures are discussed below. Experimental Section Materials. Naphthalene, [TBAm][BF4] and [THexAm][Br] were purchased from Sigma-Aldrich at a purity of 99% or greater. The solids were then further purified by drying at room temperature under vacuum of 10-3 Torr for 48 h and stored at room temperature under a nitrogen atmosphere. Tetrabutylphosphonium bromide >99% ([TBP][Br]), tetrabutylammonium bromide >99% ([TBAm][Br]), tetrabutylammonium tosylate >99% ([TBAm][Tosyl]), tetrabutylammonium tetrafluoroborate >99% ([TBAm][BF4]), tetrabutylammonium hexafluorophosphate >99% ([TBAm][PF6]), tetrapentylammonium bromide >99% ([TPAm][Br]), tetrahexylammonium bromide >99% ([THexAm][Br]), tetraheptylammonium bromide ([THepAm][Br]), tetraoctylammonium bromide >99% ([TOAm][Br]), methyltrioctylammonium bromide >97% ([MTOAm][Br]), benzyldodecyldimethylammonium bromide >99% ([BDMDDAm][Br]), and 1-hexadecylpyrdinium chloride >98% ([HDPy][Cl]) were purchased from Sigma-Aldrich (Fluka). 1-Butyl-3-methylimidazolium chloride 98% ([C4MIm][Cl]), 1-butyl-3-methylimidazolium methanesulfonate 98% ([C4MIm][CH3SO3]), and 1-butyl-3-methyl-imidazolium tosylate 98% ([C4MIm][Tosyl])

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were generously donated by Solvent Innovation. 1,5-Diethyl2-methylpyridinium ethylsulfate ([EEMPy][EtSO4]), tetrabutylammonium trifluorotris(perfluoroethyl)phosphate ([TBAm][TFEPF3]), (1S,2R)-1-hydroxy-N,N,N-trimethyl-1-phenylpropan-2-ammonium triflate ([EP][Tfo]), (S)-1-hydroxy-N,N,Ntrimethylbutan-2-ammonium triflate ([AB][Tfo]), (S)-N-(1-hydroxybutan-2-yl)-N,N-dimethylbutan-1-ammonium +mentholsulfate ([MTOA][+MS]), (R)-1-methoxy-N,N,N-trimethylbutan2-ammonium bis(trifyl)imide ([IHETMAm][Tf2N]), methyltrioctylammonium triflate ([MTOAm][Tfo]), trimethylsulfonium bis(trifyl)imide ([TMSfn][Tf2N]), tetraethylammonium bis(trifyl)imide 99% ([TEAm][ Tf2N]), and methyl-tributylphosphonium triflate ([TBMP][Tfo]) were generously donated by Prof. Dr. Peter Wasserscheidt and were all of 98%+ purity (NMR). Methyl tributylammonium triflate ([TBMAm][Tfo]) and methyltris(2-methyl-propyl)ammonium triflate ([TiBMAm][TFo]) were synthesized by direct metathesis of the tributyl (isobutyl) amine by dropwise addition of methyltriflate in dichloromethane as a solvent. The reaction was allowed to finish overnight and was evaporated and then dried under vacuum. Purity was confirmed by NMR. All samples, regardless of source, were dried under vacuum and stored under argon in Schlenk tubes. In addition, each sample was dynamically extracted with supercritical CO2 prior to measurement. Two different methods were used: one to determine the pressure and temperature at the SLV and the other to determine solely the solid-liquid equilibrium (concentrations) at pressures slightly above the SLV pressure, to prevent the formation of a vapor phase. The melting point depressions (SLV) and the PT projection of binary systems of ionic solid and compressed gas were measured using a static high-pressure view cell. The socalled “first melting” method was used here, where the solid sample was slowly heated at constant pressure and visually observed for the first signs of melting. This method is opposed to the so-called “first freezing” method where the sample is liquefied, cooled, and crystallization is observed. The apparatus consists of a high-pressure view cell (autoclave) and a compressor. Figure 2a is a diagram of the apparatus. An air-operated compressor with back-pressure regulator was built in-house for the compression of the gases. The autoclave design is similar to that used by Leitner and co-workers52 for reactions, except using view-cell glass (KDH, GmbH, PN400.4462 B 0247) with a 400 bar pressure rating. Pressure was measured with a pressure transducer (Wika, GmbH, ECO-1) with a maximum pressure rating of 400 bar, with a nominal accuracy of 2 bar. The temperature was maintained by a heating plate (Ika Werke, GmbH, PN: RET Basic C) with electronic temperature control (Ika Werke, GmbH, PN: IKATRON ETS-d4-fuzzy) using a Pt-1000 RTD placed through the wall of the autoclave. The precision is 0.1 °C and was calibrated to an ice-bath and boiling water adjusted for atmospheric pressure for a resultant accuracy of approximately 0.2 °C. All samples were dried in vacuo (∼0.1 mbar) for approximately 48 h. Approximately 200 mg of the sample (most were powders) were placed into a small sample vial capable of fitting into the autoclave. The sample vial was placed in the autoclave while under argon and sealed. The autoclave was connected with the compressor and then purged with the desired gas. For the CO2 systems, the autoclave was then charged to supercritical conditions (∼40-60 °C and ∼150-200 bar) and allowed to equilibrate for about 10 min. The system was then dynamically extracted for about 5 min. This was to ensure removal of any residual organic contaminants or water. The system was depressurized and allowed to rest for approximately

Figure 2. Experimental apparatus used in these experiments: (a) autoclave for melting point depression data (SLV) [(1) heater/stirrer, (2) windowed autoclave vessel, (3) valve, (T) RTD, (P) pressure transducer]; (b) solubility at the melting point depression (SLE).

10 min. After this purging process, the autoclave was pressurized to the desired pressure and heated slowly. With the aid of a background light, the solid sample was observed for any sign of phase change while heating at a rate of roughly 2 °C/min. Once the first sign of melting occurred, the heating was stopped and the autoclave depressurized to cool below this melting point. The process was then repeated at a much slower heating rate of approximately 0.05 °C per minute; this was determined to be sufficient to allow adequate time for thermal equilibrium for the whole system. This was repeated for a total of three times to ensure that the temperature of this solid-liquid-vapor transition was reproducible to approximately 0.5 °C (usually to 0.2 °C). After each induced melting, the autoclave was turned upright and shifted to spread the liquid over the bottom of the sample vial containing the substance in the autoclave and then depressurized to crystallize the material as a thin film. Upon return of the autoclave to a horizontal position, the film of material is now vertical and directly within the plane of the light path, which allows easier detecting of the melting point. As the material changes from a crystalline state to a liquid state, the film evolved from an opaque solid to a clear liquid accompanied by downward flow due to gravity. Many of the studied materials are known to have solid-solid transitions, amorphous, and liquid crystal phases. At times, some of these materials became more transparent but did not flow within a reasonable time-scale (∼10 min). Others began to flow before becoming completely transparent. Only when transparency and

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flow are observed is the solid-liquid-vapor temperature taken unless otherwise noted. Solid-Liquid Equilibrium. In a separate experiment, the solid-liquid equilibrium was determined at known compositions of the liquid phase. A modest hydrostatic overpressure was used to ensure our knowledge of liquid composition, but the overpressures (tens of bars) are too small to have any measurable effect on the melting point depression observed. All of these measurements of CO2 were carried out in a variable-volume windowed vessel (1.59 cm i.d., 20 cm3 maximum volume) similar to that used by McHugh53 which was used previously to measure liquid-liquid equilibrium (LLE) systems.54 A schematic of this vessel is located in Figure 2b. The vessel window and variable-volume piston were sealed with Buna-N O-rings. Phase boundaries were measured by visually observing the freezing point through a 2.54 cm diameter sapphire window (1.27 cm viewable area) with a charge-couple device (CCD) camera (Sony) mounted on a 0.635 cm borescope (Olympus). The borescope and video camera not only allowed for safe observation of the phase equilibria but also provided a significant magnification of the viewable area. The binary mixtures were stirred with a Teflon-coated stir bar coupled with an external magnet. The entire cell was placed in a thermostated air bath (modified Varian 3400 gas chromatograph) with temperature control better than (0.5 K. Precise temperature control was not required as freezing points were induced and observed while cooling the vessel. The temperature was measured with a handheld readout (HH-22 Omega) and thermocouple (Omega type K) inserted into the center of the phase equilibria vessel. The thermocouple response time was on the order of seconds. The combination of thermocouple and readout was accurate to (0.2 K and calibrated for each experiment against a platinum RTD (Omega PRP-4) with a DP251 precision RTD bench-top thermometer (DP251 Omega) accurate to (0.025 K and traceable to NIST. Back-pressure was applied to the piston with a syringe pump (ISCO 100D) operated at constant pressure. To avoid any vapor phase, the pressure was held constant at 210 bar measured with a Druck DPI 260 gauge with a PDCR 910 transducer accurate to (0.1 bar. To ensure that this method would be applicable to melting point depression measurements, experimental data were taken to compare with literature. Cheong et al.34 measured the temperature composition of naphthalene-CO2 using what they called the “first freezing point method”. This method involved observing the initial appearance of solid, followed by sampling from a high-pressure view cell. An equation of state, paired with pressure, temperature, and volume data, provided the gasphase composition. Gravimetric measurement of the solid phase after dissolution and drying gave the other compositions. Figure 3 shows the comparison of our data to the literature, where the two methods show excellent correlation. Thus, our novel freezing point method was validated, where the composition could be determined without sampling. First, a known mass of organic solid was loaded into a variable-volume high-pressure view cell, and the pure solid melting point was confirmed before adding any CO2. Then, the cell with a known quantity of CO2 from a syringe pump was loaded and allowed to reach equilibrium with temperature and pressure such that there was a single liquid phase. Next, the mixture was cooled isobarically until the first crystal formed. Generally, the entire contents froze quickly thereafter. The composition of the single liquid phase is known since there are only two components and one phase. The first crystal of pure

Figure 3. Temperature-composition (T-x) diagram for naphthaleneCO2: experimental data compared to literature values (ref 34), ideal solubility, and modeling. Table 1. Physical and Model Properties of Select Ionic Solids ionic solid

Tm [°C]

∆Hfus [kJ/mol]

Ttrans [°C]

∆Htrans [kJ/mol]

naphthalenea [TBAm][BF4]b [THexAm][Br]b

80.2 160 100

18.8 10.5 15.9

68 32

2.5 2.5

a

Reference 34. b Reference 65.

solid has a negligible impact on the overall composition, and thus, the composition at the freezing point is known. After one freezing point is measured, the system was reheated back into one phase, and the process was repeated until two consecutive data points were identical to ensure the system was at equilibrium, which can take up to 24 h, depending on the solid. The high-pressure cell was also loaded under nitrogen atmosphere, thus preventing air contamination of the solids. After secondary purification, the ionic solids were tested for water content. Samples were prepared by dissolving 1 g of solid in 0.5 mL of methanol. Methyl alcohol, extra dry with molecular sieves, ∼100 °C).

Table 4. Melting Point Depression of Quaternary Ammonium Bromide Compounds with CO2

a

Reference 32. b Possible solid-solid/liquid crystal transitions.

has only a moderate depression of roughly 20 °C, and the tosylate, with a depression of 34 °C. The highly fluorinated trifluorotris(perfluoroethyl)phosphate anion produces a tetrabutylammonium compound with a normal melting point of 54 °C

and ∆Tm of 37 °C at only 35 bar of CO2 pressure. Table 3 also indicates that analogous phosphonium compounds (tetrabutyl phosphonium vs ammonium bromides) seem to have larger melting point depressions with CO2.

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Table 5. Various Cations with Sulfates or Bis(trifyl)imide Anions

a

Reference 32.

For quaternary ammonium compounds, the cation may also have a large effect on the SLV behavior. With a common bromide anion, symmetric tetra-alkyl compounds (four equal substituents) have rather similar normal melting points at approximately 98 °C (see Table 4), from tetrabutyl- to tetraoctylammonium bromide. However, the maximum ∆Tm occurs with tetra-hexyl-ammonium bromide and CO2 with a decrease of 71 °C, while smaller and larger alkyl chains produce significantly less-pronounced ∆Tm values at similar pressures. Replacing one of the alkyl chains with a methyl group (methyl-trialkyammonium) seems to simultaneously decrease the normal melting point and increase the degree of ∆Tm (see [MTOAm][Br], [BDMDDAm][Br], and [MTOA][Br] in Table 4). Table 5 illustrates various cations with fluorinated sulfonate or bis(trifluorylmethysulfonyl)imide [Tf2N] anions. These cations include chiral ionic solids ([EP][Tfo], [AB][Tfo], [MTOA][+MS], and [IHETMAm][Tf2N]) for potential use in chiral reactions55 and separations. Of special note is the relatively simple ionic solid, tetraethylammonium [Tf2N]. This compound has a normal melting point of 102 °C but melts at 20 °C with only 35 bar of gaseous CO2. This ∆Tm at 82 °C would probably represent the largest melting point depression of all of the compounds tested if cryogenic equipment were available to

permit measurement and comparison at 150 bar. This represents an approximately 2.4 °C depression per bar of CO2 pressure. Overall, several trends can be established to increase the melting point depression with CO2 and organic ionic solids. Fluorinated anions universally increase the melting point depression. This follows the observation of Kazarian et al.,31 who found enhanced CO2 interaction with fluorinated anions by ATR-IR. Bis(trifyl)imide ([Tf2N]) anions always seem to yield the highest melting point depressions, even considering the usually lower normal melting points. Asymmetric ammonium cations which have more nonequal substituents (e.g., methytributyl ammonium [TBMAm], trimethyl-hydroxyisobutyl ammonium [AB][Tfo], etc.) tend to have lower melting points with CO2 pressure compared with equal substituents. Straight chain alkyl groups tend to have lower ∆Tm values to a greater extent than similar branched salts, e.g., [TBMAm][Tfo] vs [TiBMAm][Tfo] in Table 5. For the more planar cations, i.e., imidazolium and pyridinium, the anion seems to have a much greater effect than the cation. SLV with Pressure and Temperature. The melting point depression of several ionic solids was taken over larger pressure ranges. Figure 4 and Table 6 illustrate the global phase behavior of four ionic solids to 350 bar of CO2 pressure. As seen from

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Figure 4. Experimental and modeling of the SLV of several ionic solids and naphthalene (ref 35) and octacosane (ref 34) [smoothed data]. Table 6. SLV Equilibrium for Several Ionic Solids [TBAm][BF4]/CO2

Figure 5. T-x diagram for tetrabutyl ammonium tetrafluoroborate-CO2 with ideal solubility prediction.

[TBAm][BF4]/Ethylene

TSLV [°C]

PSLV [bar]

∆T [°C]

TSLV [°C]

PSLV [bar]

∆T [°C]

27.2 30.0 33.4 36.1 40.5 87.3 100.0 156.0

335 250 200 150 100 85 78 1

128.8 126 122.6 119.9 115.5 68.7 56 0

73 81.1 89.8 109 156.0

250 200 150 100 1

83 74.9 66.2 47 0

[TOAm][Br]/CO2

[TBMPhos][Tfo]/CO2

TSLV [°C]

PSLV [bar]

∆T [°C]

TSLV [°C]

PSLV [bar]

∆T [°C]

59.4 60.3 61.3 61.6 61.8 64.3 66.8 71.7 75.5 97.5

350.0 300.0 250.0 200.0 150.0 125.0 100.0 86.0 75.0 1.0

38.1 37.2 36.2 35.9 35.7 33.2 30.7 25.8 22.0 0.0

33.1 35.0 36.7 38.8 40.4 44.3 52.5 68.0 77.5 119.0

350.0 300.0 250.0 200.0 150.0 125.0 100.0 90.0 80.0 1.0

85.9 84.0 82.3 80.2 78.6 74.7 66.5 51.0 41.5 0.0

Table 7. T-P-Composition Data of Melting Point Depression with CO2 system with CO2

T [°C]

naphthalene -56.6 55.5 59.0 59.4 59.7 60.0 60.1 60.8 61.6 63.2 63.7 66.1 67.2 68.5 69.5 69.7 70.8 73.3 73.5 74.3 76.1 77.1 80.3

P [bar]

xCO2

1 172.4 181.3 222.6 241.3 195.2 77.6 215.8 195.1 137.9 174.6 139.9 105.0 105.1 153.8 105.1 105.1 84.3 105.3 77.1 77.2 181.3 1

1 0.838 0.556 0.549 0.536 0.498 0.438 0.528 0.448 0.365 0.371 0.311 0.262 0.239 0.239 0.239 0.203 0.143 0.154 0.098 0.062 0.096 0

system with CO2

T [°C]

-56.6 28.1 45.6 59.5 73.9 120.1 140.1 144.1 160 [THexAm][Br] -56.6 34.7 50.4 64.8 79.2 100 [TBAm][BF4]

P [bar]

xCO2

1 102.9 239.3 232.7 102.9 198.4 116.9 168 1 1 231.8 177.5 136.8 202.9 1

1 0.539 0.504 0.458 0.393 0.244 0.145 0.097 0 1 0.569 0.412 0.286 0.171 0

the figure, the change in the SLV temperature of the ionic solids with pressure is extremely large and negative in the lower pressure range, i.e., P < 150 bar. [TBAm][BF4] shows the most dramatic decrease, while [TOAm][Br] shows much more

moderate depression. At higher pressures, the increase in melting point depression is minimal with added pressure. The melting point depression of [TBAm][BF4] was also measured with ethylene and found to be large compared with most nonionic organics (naphthalene35 and octacosane34 with CO2), but still much lower than the same salt with CO2. This difference is almost surely due to variations in the solution behavior. The phase behavior of CO2 and the long-chain wax, octacosane (C28), from the work of Cheong et al.34 and the polyaromatic, naphthalene, from the work of McHugh and Yogan35 are shown for comparison. Initially, both of the solids undergo moderate melting point depression but not nearly to the same magnitude as the ionic solids. Octacosane and naphthalene both have a minimum in TSLV with pressure between 100 and 150 bar. After this minimum, the trend reverses and the melting point increases rather than decreases. At approximately 650 bar in the octacosane/CO2 system,34 the TSLV equals Tm, i.e., no melting point depression, followed by a melting point elevation at higher pressures. For the ionic solids, the largest rate in depression occurs between 0 and 150 bar. After which, the depression seems to level off approaching some asymptotic value. Whether these ionic solids experience a minimum in TSLV similar to octacosane and naphthalene is uncertain. This study was limited to 400 bar. Modeling Discussion Solid-Liquid Equilibrium. A number of differences in the results can be shown to be a function of solution nonideality. For example, the melting point depressions observed for [TBAm][BF4] with ethylene were far less than those for the same salt with CO2. Moreover, the melting point depression for naphthalene in CO2 is substantially less profound than that of the ionic liquids in CO2. The reason is almost surely due to specific interactions. Naphthalene interacts with CO2 almost entirely by dispersion (van der Waals) forces, as does ethylene with the ionic liquid. Such interactions always give positive deviations from Raoult’s law (γ > 1), and this term would become important in eq 1. Conversely, CO2 would interact with the ionic liquids in a different way, with strong unlike pair specific attractions due to the chemistry of CO2, probably due to weak Lewis acid-Lewis base interactions of the acidic carbon in CO2 with basic moieties of the saltsand similar interactions have been observed often in other situations.56-58 These would give negative deviations from Raoult’s law (γ < 1), and this term would also become important in eq 1. As previously mentioned, Figure 3 shows a comparison of literature results to model predictions for the composition versus

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Scurto and Leitner32 and elucidate the effect the CO2 has on [TBAm][BF4]. Figure 6 (Table 7) demonstrates the results for [THexAm][Br] which are very similar to that of [TBAm][BF4]. The ideal solubility is again included and it is clear that for both ionic liquids, the T-x curve falls well below the ideal solubility. This is in stark contrast to the naphthalene/CO2 data, where there are strong positive deviations from Raoult’s law. Conclusions

Figure 6. T-x diagram for tetrahexyl ammonium bromide-CO2 with ideal solubility prediction.

melting temperature and pressure for the naphthelene-CO2 system. The pressure, temperature, and composition data are listed in Table 7. Also included on this figure are the ideal solution prediction and the MOSCED (modified separation of cohesive energy density) prediction.59,60 It is clear that the ideal prediction (eq 1, with γ2 set at unity) for this system is a poor predictor for this system due to the large positive deviation from ideality, as one would expect for CO2-naphthalene interactions. The MOSCED model (eq 2), which has been recently shown to work well at predicting solid solubility,59,61 fits four parameters (hydrogen bond acidity and bascity, polarizability, and dipolarity) to experimental data to calculate infinite dilution activity coefficients.

ln γ∞2 )

[

V2 q12q22(τ1 - τ2)2 (λ1 - λ2)2 + + RT ψ1

] ()

V2 (R1 - R2)(β1 - β2) + ln ξ1 V1

aa

()

V2 +1V1

aa

(2)

The calculated activity coefficients were used in the Wilson equation48 to predict the solubility line shown in Figure 3. Several different modeling techniques have been used to predict the interaction between CO2 and ionic liquids, including the tPC-SAFT equation of state,62 the irregular ionic lattice model,63 and molecular dynamic simulations.64 However, as with most thermodynamic models, some experimental solubility data are needed for parameter tuning. For the tPC-SAFT model, vapor pressure data is needed; for the irregular ionic lattice model, solubility data is needed. MOSCED model parameters have been successfully regressed for a few ionic liquids in various solvents.59 However, there is currently insufficient data for any of the ionic solids studied here to allow MOSCED prediction with CO2. Future work in the area of thermodynamic modeling of ionic liquids will require an extensive solubility study over a wide range of ionic liquid solutes. Although no modeling is reported for our CO2-IL data, the results were compared to the ideal solubility. CO2 Solubility at the Melting Point. The results for [TBAm][BF4] are shown in Figure 5 and listed in Table 7. The ideal solubility, shown for reference, is calculated from eq 1 with each of the respective activity coefficients set equal to 1. The low-temperature region is enlarged to show the simple eutectic behavior expected in a melting point depression T-x diagram. The results show a melting point depression of 130 °C at a mole fraction near 0.5. The results seem to agree with the work of

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ReceiVed for reView March 1, 2007 ReVised manuscript receiVed May 18, 2007 Accepted May 24, 2007 IE070312B