pH-Dependent partitioning in room temperature ionic liquids provides a link to traditional solvent extraction behavior

Green Context

Room temperature ionic liquids are among the most promising new “greener” solvent systems since their lack of vapour pressure removes vapour loss and atmospheric pollution. While there is an expanding interest in their use, there remains little knowledge about their physical properties and this is holding back their exploitation and application in areas such as separation science. In particular the workhorse of solvent extraction, pH dependent reversible switching, has not been addressed for these liquids. Here for the first time, reversible pH-dependent liquid/liquid partitioning for such systems is proven. Furthermore, it is shown how this partitioning can be fine-tuned through structural variations in the ionic liquids. This new knowledge should greatly facilitate the use of room temperature ionic liquids in separations and indeed facilitate their development in other application areas.

JHC

Summary

Room Temperature Ionic Liquids (RTILs) are emerging as novel solvent replacements for Volatile Organic Compounds (VOCs) traditionally used in liquid/liquid separations, however, the basic science involved with fully characterizing these systems (physical properties, solubilities, partitioning of solutes, water miscibilities, etc.) may be artificially holding back utilization of these green solvents. By demonstrating the compatibility of RTILs with fundamental principles used in solvent extraction (SX), separation scientists may use the considerable expertise developed in SX techniques for novel separations with RTILs. Here, using a simple indicator dye, thymol blue, the following are demonstrated: (a) reversible pH-dependent liquid/liquid partitioning, (b) the use of CO₂(g) and NH₃(g) to activate the ‘proton switching’ of phase preference, (c) that structural variation within the RTIL ions may be utilized to fine-tune partitioning behavior, and (d) solid/liquid separations are possible with low melting ionic liquids.

Introduction

Over 90% of hazardous waste is aqueous,¹ and thus much of the industrial reliance on VOCs is based on the need for efficient separations from liquid media. Traditional solvent extraction employs partitioning of a solute between two immiscible phases, typically an organic solvent and aqueous solution.^2,3 One area of opportunity for new chemical science and engineering technology which will help meet the goals of sustainable industrial technology is the development of new separations technologies that eliminate the use of flammable, toxic VOCs as solvents.⁴ Used in conjunction with, or instead of appropriate current manufacturing processes, such technologies would help to prevent pollution and increase safety.

Room temperature ionic liquids⁵ are gaining worldwide academic and industrial attention^6–9 as replacements for organic solvents in catalysis,¹⁰ synthesis^11–15 and separations,^16–19 in addition to their established use in electrochemistry.^20–22 This interest is understandable given the major industrial reliance on VOCs as solvents and the ever increasing regulation of toxic release of these solvents.^23,24 The unique properties of several RTILs enable their use as solvent alternatives and may speed the introduction of these potentially ‘green’ solvents into sustainable industrial processes.

Three key communications have appeared involving the utility of RTILs in separations. The partitioning of a series of substituted aromatic solutes in RTIL/aqueous systems has been correlated to their hydrophobicity, partitioning to the RTIL increasing with 1-octanol/water partition coefficient (log P),¹⁸ but charged solutes studied tended to stay in the aqueous phase. Stripping of nonvolatile aromatic solutes from a RTIL-loaded phase was demonstrated using supercritical CO₂.¹⁶ (Although this technology may be too exotic for immediate industrial application, sc-CO₂ technologies are gaining wider acceptance.) In a third paper, a crown ether extractant was used to enhance metal ion partitioning to a RTIL phase from water, however, stripping was not demonstrated and the reported behavior appeared to be opposite to accepted trends in organic solvents.¹⁷

Despite these interesting developments and growing interest in the field, the workhorse of solvent extraction, pH-dependent reversible partitioning,²⁵ has not been addressed for these ionic fluids. In these techniques, the receiving phase is loaded at one pH and stripped at another. Such systems may involve uptake of ionizable organic solutes or extraction of metal ions using a lipophilic ligand which complexes the metal ion and transfers it to the organic phase at one pH and releases the metal ion to an aqueous phase at another pH. Adaptation of these principles of traditional liquid/liquid separation are also often made to solid/liquid and liquid membrane separations.

Results and discussion

In order to qualitatively and quantitatively study reversible pH-dependent separations in RTILs, we have investigated the partitioning of an indicator dye, thymol blue, as a function of aqueous phase pH for a series of 1-alkyl-3-methylimidazolium hexafluorophosphates ([C_nmim][PF₆]) RTIL¹⁸ (Fig. 1). The butyl, hexyl and octyl derivatives are liquid at room temperature,²⁶ stable to moisture, and form a two phase system upon contact with water. The decyl derivative is a solid at room temperature and its cyrstal structure¹⁹ indicates solid-state behavior similar to analogous higher order alkyl derivatives.^26,27

Fig. 1 Structural variation in the 1-alkyl-3-methylimidazolium hexafluorophosphate salts used in this study.

Fig. 2 illustrates the qualitative partitioning of thymol blue in its three forms between aqueous (top) and [bmim][PF₆] (bottom) phases as aqueous phase pH is changed from very acidic to very basic. The behavior of this molecule can be explained in terms of its ionization constants; 1.65 and 8.9.²⁸ At low pH thymol blue exists in its red form as a neutral zwitterion which prefers the RTIL phase. As the pH is increased via the addition of NaOH, the yellow monoanion forms with some detectable increase in concentration in the aqueous phase. The blue dianion, above pH = 10, partitions quantitatively to the aqueous phase.

Fig. 2 The phase preference of the three forms of thymol blue in [bmim][PF₆].

Adjustment of system pH, and thus phase preference of the solute, can be accomplished using mineral acids, and in this study both H₂SO₄ and HPF₆ were utilized with no observable differences based on the acid used.²⁹ A more environmentally-benign approach to pH adjustment may be accomplished using recyclable gases which can modify aqueous phase pH [e.g. CO₂(g) for acidic adjustments and NH₃(g) for basic adjustments].³⁰ To test this latter approach, a bubbler was loaded with 3 mL each of [bmim][PF₆] and a 10 mM aqueous thymol blue solution at pH = 13 [Fig. 3(a)] where the dye in the blue form remains in the upper aqueous phase. Bubbling CO₂ through the system lowered the pH of the aqueous phase to 8.3 observable by the change in dye color from blue to yellow [Fig. 3(b)]. Agitation of the system by the bubbles resulted in thorough mixing and transport of the yellow form of the dye to the RTIL phase [Fig. 3(c)]. The yellow form can be converted back to the blue dianion and stripped from the RTIL phase by bubbling NH₃ through the system [Fig. 3(d,e)]. (Further control of the pH, and thus extraction efficiency, can be accomplished by operating at increased pressure, where, for example, bubbling CO₂ through the system could lower the pH enough to observe the red zwitterionic form of the dye and further enhance distribution to the RTIL phase.)

Fig. 3 Utilizing recyclable gases to activate the proton switch: (a) Thymol blue in the upper aqueous phase at pH = 13; (b) Bubbling CO₂ through the RTIL/aqueous system, note the change in dye color; (c) The yellow form of thymol blue in the RTIL phase; (d) Bubbling NH₃ through the RTIL/aqueous system, note the darker color of the dye; (e) The blue form of thymol blue settling into the top aqueous phase.

The overall properties of an RTIL are a result of the composite cations and anions and can range from superacidic in, for example, the moisture sensitive [C_nmim][AlCl₄] salts,³¹ to water miscible (e.g. [C_nmim][BF₄] salts),³² to water immiscible and hydrophobic.^32,33 The anion is currently used to control the water miscibility, but the cation can also influence the hydrophobicity or hydrogen bonding ability. The properties of both ions are useful tools for fine-tuning an RTIL for particular ‘solvent’ properties.

To demonstrate the ability to fine-tune extraction properties, quantitative analysis of the partitioning of thymol blue was carried out using [bmim][PF₆], [hmim][PF₆], and [omim][PF₆] (Fig. 1). The subtle differences in the distribution of the solute in the three ionic liquids is shown in Fig. 4. Although the general trend is the same in each RTIL, distribution ratios increase with increasing length of the alkyl substituent on the cation, with differences in distribution ratios of over an order of magnitude.

Fig. 4 Distribution ratios of thymol blue in [C_nmim][PF₆]/aqueous systems as a function of aqueous phase pH.

Interestingly, at pH 12 the phase preference for the blue form is reversed as R is increased from C4, to C6, to C8. Distribution ratios show a distinct preference for the aqueous phase when using [bmim][PF₆], essentially no phase preference using [hmim][PF₆], and a preference for the RTIL phase when [omim][PF₆] is utilized. Thus, phase preference for specific solutes can be engineered into these designer solvents.

The decyl derivative of the RTIL used in this study, [dmim][PF₆], is a solid at room temperature, melting at 38 °C and illustrates yet another potential separation strategy. (The crystal structure of this compound was determined¹⁹ and it exhibits similar packing properties to its C12²⁶ and C14³⁵ analogs; forming bilayers with interdigitated alkyl chains.) A 10 mM solution of thymol blue at pH = 1.8 (red form) was added to a test tube containing crystalline [dmim][PF₆] [Fig. 5(a)]. When heated just above its melting temperature, the solute immediately partitions to the RTIL phase [Fig. 5(b)]. Separation of the RTIL phase, followed by crystallization at reduced temperature leads to pure [dmim][PF₆] crystals and deposits of the dye [Fig. 5(c)]. This material can be easily removed by washing with pH = 13 NaOH [Fig. 5(d)] and decanting to obtain the pure crystalline solid [Fig. 5(e)]. (The crystal structure of [dmim][PF₆] was determined both before and after melting, contact with an acidic thymol blue solution, and recrystallization, indicating no incorporation of thymol blue or water into the crystalline lattice and no change in the structural properties of the solid.)

Fig. 5 Solid/liquid separations utilizing the low melting [dmim][PF₆]: (a) An acidic thymol blue solution residing over solid [dmim][PF₆] prior to melting; (b) After melting the salt, thymol blue partitioned to the [dmim][PF₆]; (c) Recrystallization of [dmim][PF₆] leaves behind a red residue; (d) An alkaline wash is added to remove the dye from the solid surfaces; (e) The colorless crystalline [dmim][PF₆] solid after washing.

Conclusions

The results reported in this work if derived from organic/aqueous extraction data, are not surprising yet they dramatically illustrate the similarities of specific RTILs to organic diluents in traditional SX. Given the ever increasing regulatory demands on the use of VOCs, RTILs offer a potentially green alternative worthy of additional study. If the use of RTILs can be put into the general context of previous work in SX, RTILs may offer drop-in replacements of current VOC-based technologies. Such efforts, coupled with the development of RTIL libraries and physical properties databases (such as those under development at the Queen’s University Ionic Liquid Laboratory³⁶) will lead to new sustainable industrial technologies which help to eliminate the health and safety concerns which arise from the pervasive use of VOCs.

Experimental

All reagents were at least 99% pure and used as received from Aldrich (Milwaukee, WI, USA). The preparation of any aqueous solutions used 18.3 Ω cm deionized water (Nanopure). Thymol blue was used as the disodium salt. Adjustments in pH were carried out using NaOH, H₂SO₄, or HPF₆.

Partitioning of thymol blue in the three ionic liquids was monitored via UV–VIS with a Cary 3 spectrophotometer. A calibration curve for the absorbance of thymol blue in each phase was constructed over a region of concentrations that corresponded to a linear response and obeyed Beer’s Law. The partitioning of thymol blue was carried out by contacting equal volumes of ionic liquid and aqueous phase, thoroughly mixing them with two repetitions of vortexing (2 min) and centrifugation (2 min, 2000 g) followed by separation of the two phases. The absorbance was measured in each phase. The concentrations of thymol blue were determined from the calibration curve for each phase and the partitioning of the solute was calculated as a distribution ratio (D), as follows:

The solid/liquid partitioning in [dmim][PF₆] was performed by adding 3 mL of thymol blue in H₂SO₄ (pH 1.07) to solid [dmim][PF₆] (1.56 g) which was then heated to 50 °C to melt the solid. After thorough mixing, the dye partitioned to the ionic liquid phase, the aqueous phase was removed and the RTIL crystallized by cooling in a freezer (−5 °C). The red dye appeared to coat the colorless [dmim][PF₆] crystals and 3 mL of NaOH (pH 13.6) was added to remove the dye. The resulting crystalline material was colorless and determined by crystallographic analysis to be [dmim][PF₆].

The partitioning of several organic solutes was measured in a [bmim][PF₆]/water system to determine whether using H₂SO₄ or HPF₆ to adjust the pH of the aqueous solution would have an effect on their affinity for the RTIL phase. Using an aqueous phase pH of 3.3, the behavior of acetophenone, p-toluic acid, dichlorobenzene, salicylic acid, and phthalic acid were examined and indicated no observable difference in distribution ratios. (The partitioning of these organic solutes was determined radiometrically according to a method detailed in a previous report.¹⁹)

The experiments using CO₂ and NH₃ to adjust the pH were carried out in a glass tube with a frit at the bottom to allow the gases to permeate the solution. Rubber tubing was connected to the bottom of the tube and CO₂ (from an Erlenmeyer flask containing dry ice) or NH₃ (produced by passing air over concentrated aqueous ammonia) was bubbled through the biphasic system. The reported pH values are those for the aqueous phase and were measured using a Corning 220 pH meter.

Synthesis of [1-alkyl-3-methylimidazolium][PF₆]

The ionic liquids used in these experiments were synthesized according to the method outlined in a previous report.¹⁸ For the synthesis of 1-butyl-3-methylimidazolium hexafluorophosphate, equal molar amounts of methylimidazole and 1-chlorobutane in a round-bottomed flask fitted with a reflux condenser were refluxed for 24–72 h at 70 °C with stirring until two phases formed. The top phase (unreacted starting material) was decanted and ethyl acetate added (a volume approximately equal to half that of the bottom phase) followed by thorough mixing. The ethyl acetate was decanted and the procedure repeated twice using fresh ethyl acetate to ensure that any unreacted starting material was removed from the bottom phase. After the third was decanted, any remaining ethyl acetate was removed by heating the liquid bottom phase to 70 °C under vacuum with stirring.

The resulting liquid, [bmim]Cl, was poured from the reaction vessel to a 2 L plastic container lined with a perfluorinated material followed by addition of 500 mL of deionized water. HPF₆ (60% aqueous solution; 1.1∶1 stoichiometry) was added. (This addition must be made slowly to prevent the reaction temperature from rising too high.) As the addition proceeds, two phases form: the [1-butyl-3-methylimidazolium][PF₆] forms the lower phase and the upper phase is acidic. Repeated washings of the ionic liquid was carried out until the upper aqueous phase was no longer acidic.

The hexyl, octyl and decyl derivatives were prepared in a similar fashion, although the yield from the first step dropped to ca. 85–95%. All other parameters remained as described.

Acknowledgements

This research is supported by the Division of Chemical Sciences, Office of Basic Energy Sciences, Office of Energy Research, U.S. Department of Energy (Grant No. DE-FG02-96ER14673).

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