Phase Equilibria of Gases and Liquids with 1-n-butyl-3

Chapter 10

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Phase Equilibria of Gases and Liquids with 1-n -butyl-3-Methylimidazolium Tetrafluoroborate Jennifer L . Anthony, Jacob M . Crosthwaite, Daniel G. Hert, Sudhir Ν. V. K. Aki, Edward J . Maginn, and Joan F. Brennecke Department of Chemical Engineering, University of Notre Dame, South Bend, IN 46556

The solubility of carbon dioxide, methane, nitrogen, carbon monoxide and benzene in 1-n-butyl-3-methylimidazolium tetrafluoroborate ([bmim][BF ]) at 25 °C are reported. In addition, we present densities, which compare very favorably with one set of literature values, for [bmim][BF ]. Accurate values of density are needed in the analysis of gas solubility data. Finally, we report liquid-liquid equilibrium data for [bmim][BF ] with 1-propanol and benzene. Information on vapor-liquid and liquid-liquid equilibrium of gases and liquids with ionic liquids is important for evaluating the use of ionic liquids in a wide variety of processes, including reactions and separations. 4

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Introduction Here we present preliminary results for the solubility of various gases in 1n-butyl-3-methylimidazolium tetrafluoroborate ([bmim][BF ]). Previously, we have measured the solubilities of a wide variety of gases in l-w-butyl-3methylimidazolium hexafluorophosphate ([bmim][PF ]) (1). The goal of this work is to determine the influence of the nature of the anion on gas solubilities. 4

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© 2003 American Chemical Society In Ionic Liquids as Green Solvents; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2003.

Ill We know of no other published values of the solubility of gases in [bmim][BF ], C 0 is the most soluble gas that we have found in [bmim][PF ] (almost 2 mole % at 1 bar and 25°C), with C2H4, C H6 and CH4 being progressively less soluble (1,2). 0 and Ar were just sparingly soluble and the solubility of H , CO and N were below the detection limit of our apparatus (1). Our motivation for measuring gas solubilities is three-fold. First, various researchers have shown that ionic liquids (ILs) are excellent solvents for a wide variety of reactions (3-5) involving permanent gases, such as hydrogénations, oxidations and hydroformylations. From our previous results (1), it appears that the solubility of H , 0 and CO in ILs is likely to be a limiting factor in these reactions. Second, the solubility of carbon dioxide (C0 ) in ionic liquids is important for evaluating the possibility of using supercritical C 0 to extract solutes from ionic liquids, as we have proposed previously (6,7). Finally, we are interested in exploring the possibility of using water stable ILs for performing gas separations. We also present measurements of the density of [bmim][BF ] at temperatures between 20 and 70 °C. Accurate densities are required in the analysis of the gas solubility data (1). As described below, we measure the gas solubilities with a gravimetric microbalance. The mass uptake must be corrected for buoyancy effects, which requires the liquid density. There were several papers available in the literature that give densities for [bmim][BF ] but the values were not consistent, necessitating the performance of our own measurements. Suarez et al. (8) reported values between 6 and 81 °C of 1. Μ Ι.18 g/mL. Seddon et al. (9) reported values between 1.16 and 1.21 g/mL for temperatures between 20 and 90 °C. Rogers and coworkers (10) report the density of [bmim][BF ] to be just 1.12 g/mL at 25 °C. One factor that could account for some variability is the water content. [bmim][BF ] is completely miscible with water at room temperature (11) and is very hygroscopic. Suarez et al. (8) do not report the water content of their sample. The sample used by Seddon et al. was reported to contain just 0.03 wt% water, while the sample tested by Rogers and coworkers (10) contained 0.45 wt% water. While water content can affect density (12), it is unlikely that water content can fully explain the large discrepancies in the reported values. Finally, we present the liquid-liquid equilibrium (LLE) of benzene and 1propanol with [bmim][BF ]. LLE is important for determining the solubility of reactants and products in ILs, as well as the amount of IL that will contaminate organic or aqueous phases that are brought into contact with the IL. Marsh has measured the LLE of a series of dialkylimidazolium hexafluorophosphates with ethanol, 1-propanol and 1-butanol (13). Najdanovic-Visak et al. have studied the ethanol/water/[bmim][PF ] system, as well as the ethanol/[bmim][PF ] and water/[bmim][PF ] binaries (14). We have provided a few numbers on water/IL LLE (17) and some preliminary estimates of the solubility of a wide variety of organic liquids in [bmim][PF ] (7). DatafromRogers and coworkers (15,16) for 4

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In Ionic Liquids as Green Solvents; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2003.

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some water/DL systems show lower mutual solubilities than other published results (14,17). All of these systems show upper critical solution temperature (UCST) behavior; i.e. the mutual solubilities increase with increasing temperature.


Experimental The gas solubility measurements reported here were made using a gravimetric microbalance (IGA 003, Hiden Analytical) that is normally used to measure gas adsorption onto solids. However, since the ILs are nonvolatile they could be used with this apparatus. This microbalance and the technique used to measure gas solubilities in ILs has been described in detail elsewhere (1,17,18). Briefly, a small IL sample is placed on the balance and the system is thoroughly evacuated to remove any volatile impurities, including water. The gas is introduced and the gas solubility is determined from the mass uptake. The [bmim][BF ] density was determined using a 1 mL pycnometer. At temperatures other than ambient, the sample was equilibrated for at least 30 - 40 minutes in a thermostatted gas chromatography oven. Care was taken to avoid any exposure to air that could result in water uptake. The water content of the sample tested was 0.2 wt%, as determined by Karl-Fisher titration. 4

Figure 1. Experimental apparatus for determining liquid-liquid solubility of ionic liquid-organic systems; TO Temperature controller for immersion heater, 77: Temperature indicator for samples.



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A cloud point apparatus was used for determining the liquid-liquid equilibrium of IL/organic mixtures. This apparatus is shown schematically in Figure 1. In this technique, a sample of known composition was placed in a viewcell. The "cloud point temperature" is the temperature at which the sample initially changes from one phase to two phases; that phase transition was determined visually. In our apparatus, four solutions of ionic liquid and an organic liquid at different concentrations were prepared gravimetrically in 5 mL viewcells in a glovebox, sealed from the atmosphere, and placed in a water bath. The uncertainty in the compositions is estimated to be +0.0001 mole fraction. The samples were heated in the water bath using a 1000 W immersion heater to about S °C above the highest expected cloud point temperature and maintained at that temperature for 10 minutes. The samples were cooled using a recirculating water chiller at a cooling rate of about 0.01 °C/s. The temperature of each sample was individually measured using a T-type thermocouple. The experiment was repeated four additional times to determine accurate and reproducible values (+0.5 °C). The CH4 was from Matheson Gas Products with a purity of 99,99+%. The C 0 was from Scott Specialty Gases with a 99.99% purity. The CO and N were both from Mittler Supply Co. and had purities of 99.97% and 99.99%, respectively. Benzene and 1-propanol were purchased from Aldrich with reported purities of 99.9+% and 99+%, respectively. Both benzene and 1propanol were redistilled prior to use. [bmim][BF ] was synthesized according to standard procedure (19) and was analyzed by NMR spectroscopy. The residual chloride content was < 10 ppm, as measured by a Cole-Parmer chloride ion specific electrode, calibrated with [bmim][Cl], 2

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Results and Discussion The solubility of C 0 , CH4, N , and CO were measured in [bmim][BF ] at 25 °C and pressures to 13 bar. The solubilities of N and CO were below the detection limit of our instrument. The results for C 0 and CH4 are shown in Figures 2 and 3, below, where the solubilities are plotted as mole fraction of the gas in the IL at various pressures. The solubilities in [bmim][BF ] are compared with those in [bmim][PF ] (1). The solubility of C 0 in the ILs is significantly greater than the solubility of C H ^ Moreover, it is obvious that the solubilities of each of these two gases in [bmim][BF ] and [bmim][PF ] are very similar. In other words, the nature of the anion (BF versus PF ) has very little effect on the solubility of these gases. The simplest way to analyze gas solubility is in terms of a Henry's law constant. The Henry's law constant is defined as: 2

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Where H is the Henry's law constant, η is the fiigacity of the solute in the liquid phase and x, is the solute mole fraction in the liquid phase. For vapor/liquid equilibrium of IL/gas mixtures, the equifugacity criterion can be written as:

Since the IL is nonvolatile, we assume that the vapor phase is pure gas; therefore, the molefractionof gas, >% is unity. Since the Henry's law constant is determined at the lowest pressures, there is no need to correct for vapor phase nonidealities with a fugacity coefficient, 0/. With these simplifications, the Henry's law constant is just the slope of the solubility versus pressure plots shown in Figures 2 and 3, evaluated in the limit of low solubility.

0.25 0.20

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Pressure (bar) Figure 2. Solubility isotherms at 25 °Cfor C0 in [bmim][PF ] and [bmim][BF ]. 2

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At 25 °C, the Henry's law constant for C 0 in [bmim][BF ] is 56.5 ± 0.3 bar and for CH4 in [bmim][BF ] is 1560 ± 325 bar. By comparison, the values in [bmim][PF ] are 53.4 ± 0.3 bar for C 0 , and 1690 ± 180 bar for C H * These Henry's law constants, as well as those for all the other gases studied in [bmim][PF ] are shown in Figure 4. The striped bars in Figure 4 represent the estimated detection limits of the instrument for CO, N , and H . As mentioned above, accurate density measurements are required for the buoyancy correction needed to determine the values ih Figures 2 and 3 from the raw data of mass uptake. Our measurements for the density of [bmim][BF ] at temperatures between 20 and 70 °C are shown in Figure 5, along with the data of Seddon et al. (9) and Suarez et al. (8). 2

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Pressure (bar) Figure J. Solubility isotherms at 25 °C for CH in [bmim][PF ] and [bmim][BF ]. 4

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Figure 4. Henry's law constants for different gases in [bmim][PF ] and [bmim][BF ]at25°C 6

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• Suarez et al., J. Chlm. Phys., 1998. 1.12 20

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Temperature, °C Figure 5 Temperature dependence of the density of [bmim][BF ]. 4

A linear regression through our data yields a temperature dependence given by the following equation: ρ (g/mL) = -6.43 χ ΙΟ" Τ (°C) + 1.22. Our data is completely consistent with that of Seddon et al. (9). However, our values are significantly higher than those reported by Suarez et al. (8). Since those researchers did not report water content in that relatively early paper, it is likely that their densities are low due to large water content. As mentioned above, [bmim][BF ] is totally miscible with water and is highly hygroscopic. However, the value reported by Rogers and coworkers (10) of 1.12 g/mL at room temperature cannot be explained by the water content, which was reported to be just 0.45 wt%. This value is inconsistent with all the other recent measurements. We used the linear regression through our data to determine the densities used in the buoyancy corrections to the raw gas solubility data. Using the gravimetric microbalance, we have also measured the solubility of benzene vapor in [bmim][BF ] at 25 °C. Measurements were made up to about 75% of the vapor pressure of benzene (P = 0.126 bar at 25 °C). These results are shown in Figure 6, where the solubility of the benzene in the IL in mole fraction is plotted as a function of the pressure, which has been normalized by the vapor pressure. Benzene is very soluble in [bmim][BF ], with a Henry's law constant of 0.25 bar. This value is shown in comparison to other gases above in the bar graph in Figure 4. Also shown in Figure 6 is the solubility of liquid benzene in [bmim][BF ]. The vapor-liquid equilibrium data match the liquid-liquid equilibrium data extremely well when extrapolated to the saturation pressure (i.e., P/P -> 1), as expected. 4

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Figure 6. Solubility of benzene in [bmim][BF ], 4

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β 20]

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Figure 7 Liquid-liquid equilibrium of 1-propanol and [bmim][BF ] as a function ofwt% IL 4

In addition, we have used the cloud point apparatus to measure the L L E of fbmim][BF ] with 1-propanol. These results are shown in Figures 7 and 8, where the data are plotted both as a function of weight % and mole % EL. For comparison, the data of Marsh (13) for 1-propanol with [bmim][PF ] is also shown in the figures. Here the nature of the anion has a dramatic effect on the phase equilibria. The upper critical solution temperature is significantly lower for the [bmim][BF ] system and at a given temperature the mutual solubilities are much greater. 4

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Figure 8 Liquid-liquid equilibrum of 1-propanol and [bmimllBFd as a function of mol % IL

Summary We have presented measurements for the solubility of gases in l-#i-butyl-3methylimidazolium tetrafluoroborate ([bmim][BF ]). The values are very similar to those for [bmim][PF ], indicating that the nature of the anion has very little effect on the solubility of permanent gases in ILs. The Henry's law constants for C 0 and CH4 in [bmim][BF ] at 25°C are 56.5 ± 0.3 bar and 1560 ± 325 bar, respectively. Our measurements of densities of [bmim][BF ] as a function of temperature, which are needed to analyze the gas solubility data, closely match those of Seddon et al. (7) but are significantly greater than other published values (6, 8). Vapor-liquid and liquid-liquid equilibria measurements for benzene in [bmim][BF ] are internally consistent and indicate that benzene is highly soluble in this IL. We also present liquid-liquid equilibria data for 1propanol and [bmim][BF ] that, when compared with L L E data for 1propanol/[bmim][PF ], suggest that the anion does plays a very important role in determining liquid-liquid equilibria. 4

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Acknowledgments Financial support from the National Science Foundation (grant CTS 9987627), a Bayer Predoctoral Fellowship, and a National Science Foundation Graduate Research Traineeship Fellowship (grant, 9452655) are gratefully acknowledged.


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behaviour of room temperature ionic liquid solutions: an unusually large co-solvent effect in (water + ethanol). Phys. Chem. Chem. Phys. 2002, 4, 1701-1703. Visser, A. E.; Holbrey, J. D.; Rogers, R. D. Hydrophobic ionic liquids incorporating N-alkylisoquinolinium cations and their utilization in liquid-liquid separations. Chem. Commun. 2001, 2484-2485. Swatloski, R. P.; Visser, A. E.; Reichert, W. M . ; Broker, G. Α.; Farina, L. M . ; Holbrey, J. D.; Rogers, R. D. On the solubilization of water with ethanol in hydrophobic hexafluorophosphate ionic liquids. Green Chem. 2002, 4, 81-87. Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. Solution Thermodynamics of Imidazolium-Based Ionic Liquids and Water. J. Phys. Chem. Β 2001, 105, 10942-10949. Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. Gas Solubilities in 1-n -butyl-3-methylimidazolium hexafluorophosphate. In Ionic Liquids; Rogers, R. D., Seddon, K. R., Eds.; American Chemical Society: Washington D.C., 2002; ACS Symp. Ser. No. 818, pp 260-269. Cammarata, L.; Kazarian, S. G.; Salter, P. Α.; Welton, T. Molecular states of water in room temperature ionic liquids. Phys. Chem. Chem. Phys. 2001, 3, 5192-5200.


Phase Equilibria of Gases and Liquids with 1-n-butyl-3

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