Gas Solubility - American Chemical Society

Chapter 1

Downloaded by UNIV OF CALGARY on September 5, 2013 | http://pubs.acs.org Publication Date (Web): December 18, 2009 | doi: 10.1021/bk-2009-1030.ch001

Gas Solubility (and Related High-Pressure Phenomena) in Systems with Ionic Liquids Gerd Maurer and Dirk Tuma Department of Mechanical and Process Engineering University of Kaiserslautern, Kaiserslautern, Germany

A review is presented on the solubility of single gases in pure ionic liquids and on the volume expansion of the liquid phase induced by a dissolved gas. Furthermore, the phenomenon of “salting-out” by a near-critical gas (SONG) is discussed. The SONG phenomenon describes the liquid-liquid phase split that might be observed when a binary liquid solvent mixture – here with an ionic liquid component – is pressurised by a gas at a temperature close to the critical temperature of that gas. Due to the large scope of literature on gas solubility in ionic liquids, this review is mainly restricted to work by the authors’ research group.

Introduction Gas solubility in ionic liquids is of interest in many areas, for example, in chemical reaction engineering when dissolved gases react with other solutes (such as in biphasic hydrogenation and hydroformylation), or for the storage of hazardous gases. The design of processes/equipment requires reliable data on gas solubility as well as on the phenomena that possibly occur in such processes. Two of the most important phenomena are the volume expansion of the liquid © 2009 American Chemical Society

In Ionic Liquids: From Knowledge to Application; Plechkova, N., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2010.

1

2


phase and the appearance of a liquid-liquid phase split. In previous (and ongoing) laboratory research, we have investigated the solubility of various nonreacting (single) gases in some pure ionic liquids and the volume expansion of the liquid phase caused by a dissolved gas. In theoretical work, the solubility of single gases in ionic liquids has been predicted by Gibbs Ensemble Monte Carlo simulations. In many applications, gases are dissolved in mixtures that contain an ionic liquid. Thus, experimental work has been extended into that field, in particular to the liquid-liquid phase split that might occur when a completely miscible solvent mixture is pressurised by a gas close to its critical temperature. Here, we review the progress that has been achieved in these research areas

Experimental Arrangement for Measuring the Solubility of a Gas in an Ionic Liquid The most important experimental techniques used to determine the solubility of a gas in a liquid are the gravimetric, the volumetric, the NMRspectroscopic, and the synthetic methods (1,2). In the gravimetric method, a high-resolution balance is applied to determine the increase of mass when a gas is dissolved in a liquid at preset temperature and pressure. This method needs only small amounts of the solvent, and it can provide data of excellent quality, in particular when the gas solubility is sufficiently high and the volume expansion of the liquid remains small. The volumetric method generally requires more solvent, and the amount of dissolved gas is determined volumetrically either during absorption or desorption. There are a lot of high-precision gas solubility studies known where this method was applied. The NMR-spectroscopic method is a very fast, but less accurate technique. It is mostly used to screen solvents for gases. In the synthetic method, the pressure is determined that is required to dissolve a known amount of gas in a known amount of solvent at a preset temperature. This technique is capable of delivering high-precision data for scarcely, as well as for very, soluble gases. It requires approximately the same amount of solvent as the volumetric method. We have a long-standing expertise in applying the synthetic method for measuring the solubility of single gases in pure solvents and solvent mixtures at temperatures from below room temperature up to about 200 °C (3). Figure 1 shows a scheme of such an apparatus as it was used to determine the solubility of a single gas in a pure ionic liquid (3). The heart of this apparatus is a thermostatted, cylindrical highpressure view cell (with a cell volume of about 30 cm3) with a sapphire window on each side. In an experiment, the evacuated view cell is charged with an exactly known amount of gas (determined either volumetrically or gravimetrically) before the solvent (here, the ionic liquid) is added via a calibrated high-pressure spindle press. The amount of solvent added is


3


sufficiently large to generate a single liquid phase in the view cell. This phase is equilibrated and, afterwards, very small amounts of the mixture are withdrawn via the spindle press until the first stable gas bubbles appear. After each withdrawal (i.e., each small pressure drop), the liquid mixture is equilibrated again. When the pressure drop in the last step (i.e., the step that induces a twophase system in the cell) is sufficiently small, the solubility pressure is reliably determined as the arithmetic mean of the pressure before and after that last step.

F

E D

D

D

D

A G

C

B

K

J

I

H

Figure 1. Scheme of the apparatus for measuring the solubility of a single gas in a solvent at elevated pressures: (A) cylindrical high-pressure equilibrium view cell with two sapphire windows and a magnetic stirrer, (B) thermostat, (C) container for the gas, (D) pressure transducers, (E) tank for rinsing water, (F) tank for the solvent, (G) high-pressure spindle press, (H) AC-bridge with three platinum resistance thermometers, (I) solution outlet, (J) cooling trap, and (K) vacuum pump. As both the volume of the cell (from a calibration experiment with pure gas or by filling the cell with pure liquid water) and the amounts of solvent (through the volume displacement of the calibrated spindle press and the density of the solvent) and solute are exactly known, a gas solubility experiment also provides volumetric data that can be used to determine the volume expansion of the ionic liquid by the dissolved gas.


4 Table 1 provides an overview of the (gas + ionic liquid) systems that were investigated so far in our group. For a comparison with literature data, the reader is referred to the review article by Anderson et al. (2) and the references cited therein.


Table 1. References for Experimental Work on Gas Solubility in Ionic Liquids in the Authors’ Group CO2 CO H2 O2 CH4 CF4 Xe

[C4mim][PF6]a

[C6mim][NTf2]b

[C4mim][CH3SO4]c

[C4mpyr][NTf2]d

(4, 10) (5) (7) (6) — — —

(8) (15) (9) (15) (13) (14) (13)

(10) (11) (11) — (12) — (12)

(16) — (16) — — — —

a

[C4mim][PF6] = 1-butyl-3-methylimidazolium hexafluorophosphate

b

[C6mim][NTf2] = 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide

c

[C4mim][CH3SO4] = 1-butyl-3-methylimidazolium methylsulfate

d

[C4mpyr][NTf2] = 1-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide

As a typical example, Figure 2 shows the experimental gas solubility data for the solubility of CO2 (left) and H2 (right) in the ionic liquid [C4mpyr][NTf2] as a plot in which for temperatures between 293 K and 413 K the solubility pressure is shown versus the gas molality (i.e., the number of moles of dissolved gas per kilogram of ionic liquid). At constant temperature, the solubility pressure increases nearly linearly with increasing gas molality – a behaviour typically observed for “physical gas solubility”. Although – at first glance – both parts of Figure 2 look rather similar, there are large differences in the solubility of CO2 and H2 in [C4mpyr][NTf2]. For example, at the same temperature and pressure, the gas molality differs by a factor of about ten, hydrogen being the less soluble gas. Furthermore, the influence of temperature is substantially different for both gases in that ionic liquid. The solubility of CO2 in [C4mpyr][NTf2] decreases with increasing temperature, whereas that of H2 increases. The experimental results are correlated by an extended version of Henry’s law on the molality scale.

kH,G,IL(T, p) aG(T, mG) = fG(T, p)


(1)

5 kH,G,IL(T, p) is Henry’s constant on the molality scale of a gas (G) in the ionic liquid (IL) at temperature T and pressure p. aG(T, mG) is the activity of the gas in the liquid that is related to the gas molality (mG) and the activity coefficient coefficient (γG) through: m (2) aG = G γ G , m°

10

10

8

8

p / MPa


where m° = 1 mol kg–1. fG(T, p) is the fugacity of the (pure) gas in the vapour phase.

6

6 4

4

2

2

0

0

1

4 2 3 m(CO2) / (mol kg−1)

5

0

0

0.05 0.1 0.15 m(H2) / (mol kg−1)

Figure 2. Total pressure above solutions of (CO2 + [C4mpyr][NTf2]) and (H2 + [C4mpyr][NTf2]): (▲, 293.1 K (CO2), 293.2 K (H2); □, 333.1 K (CO2, H2); ○, 373.2 K (CO2), 373.1 K (H2); Δ, 413.2 K (CO2, H2)) experimental results; (–––––) correlation.

Henry’s constant of the gas (G) in an ionic liquid (IL) is expressed as:

∞ ⎛ Vm, ⎞ (0) ⎜ G, IL p ⎟ k H, G, IL (T , p ) = k H, G, IL (T ) exp⎜ ⎟ ⎝ RT ⎠


(3)

6 (0) ∞ In Equation (3), k H, G, IL (T ) is Henry’s constant at zero pressure, Vm, G, IL stands

for the partial molar volume of gas G at infinite dilution in the ionic liquid (IL), and R is the universal gas constant. Henry’s constant at zero pressure (0) k H, G, IL (T ) is determined by an extrapolation (at fixed temperature):


⎡ f G (T , p ) ⎤ (0) k H, ⎥ G, IL (T ) = lim ⎢ p → 0 (mG / m°) ⎣ ⎦

(4)

The fugacity is calculated from the solubility pressure p and the fugacity coefficient φG (which is calculated from an equation of state for the pure gas): fG(T, p) = p φG

(5)

Figure 3 shows the results from these evaluations for seven gases in the ionic liquid [C6mim][NTf2] together with Henry’s constant (transformed to the molality scale) from the literature (17–22).

According to Henry’s law mG p = (0) m° k H, G, IL

(6)

at low pressures (i.e., at low gas molalities), the solubility of a gas in a solvent is proportional to the reciprocal value of Henry’s constant. Figure 3 reveals the large differences in the solubility of gases in a particular ionic liquid as well as the differing influence of temperature on gas solubility.


7 100

T/K

373

413

293

333

H2

O2

CO CH4 10


(0)

kH,G(T) / MPa

CF4

Xe CO2

1 2.4

2.8

1000 / (T / K)

3.2

3.6

Figure 3. Henry’s constant of some gases in [C6mim][NTf2]. Solid lines: own data (cf. Table I). Other literature data: CO2: , (17); □, (18); ▲, (19); ●, (20); x, (21); CH4: ◊, (18); ■, (20); +, (21); CO: , (22); H2: ○, (17), Δ, (21). There is sound agreement between the Henry’s constants evaluated from our experimental data for [C6mim][NTf2] and the corresponding literature data for the gases CO2, CH4, and CO. The average relative deviation between our values and a corresponding series of reported values for Henry’s constant amounts to between 1.9 % (17) and 14.9 % (19) for CO2, 3.4 % (20) and 6.9 % (21) for CH4, and to 6.3 % for CO (22), respectively. For H2, however, the data from Costa Gomes (17) display an opposite temperature dependence, which is contrary to our results (7,9,11,16). The corresponding average relative deviation assigned to that data amounts to 24 %. At higher pressures, the influence of pressure on Henry’s constant has to be taken into account via the partial molar volume of the gas in the ionic liquid, cf. Equation (3). For very soluble gases (e.g., CO2 in [C6mim][NTf2] or [C4mpyr][NTf2], respectively), their partial molar volume was determined from the volumetric data that were obtained from the gas solubility measurements (23).


8

Volume Expansion The volume V of the liquid mixture in the equilibrated view cell (i.e., the volume of that view cell) is


V = nIL Vm,IL + nG Vm,G ,

(7)

where nIL and nG are the amounts of mass (number of moles) of the solvent and the gas in the cell, respectively, and Vm,IL and Vm,G are the partial molar volumes of the solvent and the gas, respectively. Dividing Equation (7) by the mass of ~ results in: the solvent m IL

V / (cm 3 mol −1 ) mG V / cm 3 1000 ~ = 1000 m, IL + Vm, G / (cm 3 mol −1 ) , mIL / g M IL m°

(8)

where MIL is the relative molar mass of the ionic liquid. Generally, Vm,IL and Vm,G depend on temperature and composition of the mixture. Therefore, the partial molar volume of the gas (G) at infinite dilution has to be determined by an extrapolation to zero gas molality:

⎡ ⎛ 1000 V / cm 3 ⎢∂⎜⎜ ~ /g m ⎢ ⎝ IL mG / m ° →0 ⎣

lim) (

⎞ ⎟ ⎟ ⎠

⎛m ∂⎜ G ⎝ m°

⎞⎤ ∞ 3 −1 ⎟⎥ = Vm, G, IL / (cm mol ) ⎠⎦⎥ T

(9)

At low gas molalities, the partial molar volumes of the solvent and the solute (0) can be approximated by the molar volume of the pure ionic liquid Vm, IL and the ∞ partial molar volume of the gas in that solvent at infinite dilution Vm, G, IL ,

respectively. In the concentration range where these approximations are valid, Equation (8) results in a linear relationship:

V (0) / (cm 3 mol −1 ) mG ∞ V / cm 3 1000 ~ Vm, G, IL / (cm 3 mol −1 ) = 1000 m, IL + mIL / g M IL m°


(8a)


9

Figure 4. The volume to solvent mass ratio versus the solute molality for CO2 in [C6mim][NTf2]: (▲, 293.15 K; □, 333.15 K; ●, 373.2 K; ◊, 413.2 K) experimental results; (–––) linear fit. ∞ The slope of the line is Vm, G, IL . Equation (8a) was employed to evaluate the

volumetric data. Figure 4 shows, as a typical example, the evaluation of the volumetric data for the system (CO2 + [C6mim][NTf2]). For sparsely soluble gases, where the gas molality is much smaller than shown in Figure 4, a reliable evaluation of the volumetric data for the determination of the partial molar volume of the gas in the ionic liquid is not possible. However, it can then be determined in a different procedure from the gas solubility data (17). Figure 5 shows the partial molar volumes of four gases in [C6mim][NTf2]: two very soluble gases (CO2 and CH4) and two sparsely soluble ones (CO and H2). The values of the partial molar volume for the four ∞ gases G range from about 10 to 45 cm3 mol–1. For a rough estimate, Vm, G, IL

can be replaced by the molar volume of the gas G in its liquid state at the triple point. The partial molar volumes can – together with the molar volumes of the ionic liquids – be used to determine the volume expansion which accompanies the dissolution of the gas. This volume expansion depends on the solubility (i.e., to a first estimate on Henry’s constant), on the partial molar volume of the dissolved gas at infinite dilution, and on the molar volume of the solvent. Table 2 shows some typical examples for the relative volume expansion, when a gas is dissolved at 373 K in [C6mim][NTf2] to reach a pressure of 10 MPa. Notably, there are large differences in the volume expansion. In those particular examples, the relative volume expansion stretches from only 0.4 % (for the least soluble gas H2 that also has the smallest partial molar volume) to 16 % for CO2


10 (which is one of the most soluble gases and also has the largest partial molar volume). The volume expansion is not a particular phenomenon for (gas + ionic liquid) systems, but is also common in other (gas + solvent) systems. The extent of those volume expansions is also within the commonly observed ranges, cf. Jessop and Subramaniam (24). Table 2. Volume Expansion Caused by the Pressurisation of the Ionic Liquid [C6mim][NTf2] with a Gas to 10 MPa at 373 K (23) ∞ Vm, G, IL

(0) k H, G, IL

(cm 3 mol −1 )

MPa

44 52 36 39 18

3.81 6.08 19.2 38.8 60.1

CO2 Xe CH4 CF4 H2

50

∞

Vm,G,IL / (cm3 mol–1)


Gas G

CO2

40

CH4

30

CO

20 10 290

Relative Volume Expansion (100 ΔV/V)/% 16 11 2.3 1.1 0.4

H2 310

330

350 T/K

370

390

410

Figure 5. Influence of temperature on the partial molar volume of some gases in the ionic liquid [C6mim][NTf2].

Prediction of Gas Solubility by Molecular Simulation Experimental data for the solubility of gases in ionic liquids are not only important from an application-oriented point of view, but they also provide a



11 data base for the development of correlation methods and predictive tools. Molecular simulation is such a tool, and it has attracted a steadily increasing attention in recent years. It allows for predicting thermodynamic properties from information on intermolecular pair potentials between the components of a mixture (25–30). Figure 6 shows – as typical examples – predictions from Gibbs Ensemble Monte Carlo simulations for the solubility of CO2 in [C4mim][CH3SO4] – upper diagram (Urukova and Maurer (31)) – and CO in [C4mim][PF6] – lower diagram (Urukova et al. (32)). For the simulations, the pair potentials of the gases were taken from the literature (CO: Straub and Karplus (33); CO2: Harris and Yung (34)). Both pair potentials are the sum of Coulombic interactions and Lennard-Jones contributions. The pair potentials of both ionic liquids were similarly approximated by multicentre electrical point charges in combination with multicentre Lennard-Jones contributions. For [C4mim][PF6], the pair potential was taken from Shah et al. (35); for [C4mim][CH3SO4], it was estimated from quantum chemical calculations (for the geometry and the electrical point charges) and the United Atoms approach for the Lennard-Jones contributions (Urukova and Maurer (31)). Classical mixing rules without adjustable parameters were used to describe interactions between different Lennard-Jones sites. As shown in Figure 6, the prediction results for the solubility of CO2 in [C4mim][CH3SO4] at 333 K almost perfectly agree with the experimental data. However, there are large deviations between the predictions results and the experimental data for the solubility of CO in [C4mim][PF6], where the molecular simulation results overestimate the gas solubility by a factor of about three. Thus, dependent on the input, the results from molecular simulation vary from merely capturing trends to an almost quantitative agreement. According to Maginn (27), an upgrade of the performance is to be expected from not only force field refinement, but also from more sophisticated sampling procedures.


12

10

p / MPa

8 6 4

0

0

0.5

1.0 1.5 m(CO2) / (mol kg−1)

2.0

2.5

0

0.1

0.2 0.3 m(CO) / (mol kg−1)

0.4

0.5

10 8

p / MPa


2

6 4 2 0

Figure 6. Gas solubility in ionic liquids. Comparison between experimental data (full symbols) and data from molecular simulation (open symbols). Top: Solubility of CO2 in [C4mim][CH3SO4]; ■, 333.1 K; Bottom: Solubility of CO in [C4mim][PF6]; ■, 293 K; ●, 334 K; ▲, 373 K.


13

Salting-Out by a Near-critical Gas in Systems with Ionic Liquids

p2 > p°

p1 = p° organic solvent

0.5

0.5

L2

lv.)

gas

T = const.

. so

x(

L

LV

1 V

organic solvent

1

org

gas

)

0

x(


When a gas is dissolved under pressure in a solvent mixture (for example, an aqueous solution of a low molecular-weight alcohol) at a temperature near its critical temperature, a liquid-liquid phase split might be induced (36). Then, both liquid phases (L1 and L2) coexist with a vapour phase V. This phenomenon is called “Salting-Out by a Near-critical Gas” (SONG), and that particular term directly dates from the paper by Elgin and Weinstock, who were the first to report this, in 1959 (37). An example is shown in Figure 7.

L2

0

water

V

L1L2 L1

L1

V

gas

water

Figure 7. Typical phase behaviour in a system (“near-critical” gas + water + organic solvent). Here, p° = 0.1 MPa.


14

CO L

L

G P C


T

AT SL

SL

LP SV

SV

SV

GC

HPLC

DM

Figure 8. Scheme of the apparatus for measuring phase equilibria at elevated pressures: (AT) air thermostat; (C) pivot-mounted cylindrical high-pressure view cell with two sapphire windows and a magnetic stirrer; (CO) compressor; (DM) densimeter; (G) gas container; (GC) gas chromatograph; (HPLC) highpressure liquid chromatograph; (L) high-pressure spindle press containing liquid solvents; (LP) liquid pump; (P) pressure transducer; (SL) one of two sampling loops; (SV) sampling valve; (T) platinum resistance thermometer.



15 Such liquid-liquid phase splits are also observed when a completely miscible binary mixture of an ionic liquid (e.g., [C4mim][PF6]) and an organic solvent (e.g., methanol) is pressurised by CO2 at temperatures below, as well as above, the critical temperature of CO2 (Tc = 304.13 K) (38–41). A scheme of the apparatus to investigate the SONG phenomenon is given in Figure 8 (42). The equipment is similar to that employed for measuring gas solubilities (cf. Figure 1), but it operates according to a static analytical method, and is thus equipped with two sampling loops. When the mixture to be analysed is equilibrated, samples of the coexisting (liquid) phases are taken and analysed by gas chromatography (the volatile compounds, e.g., CO2 and methanol) and by high-pressure liquid chromatography (the ionic liquid), respectively. Two phases can be analysed simultaneously, and the density is also recorded. Figure 9 shows a scheme of such phase equilibria at a temperature above the critical temperature of CO2, for example, at 313 K. On the left-hand side of Figure 9, triangular diagrams – like those shown in Figure 7 – are arranged into a prism, where the pressure increases from bottom to top. The SONG phenomenon exists only in a limited pressure range, i.e., at pressures between the “lower critical endpoint” (LCEP) line and the “upper critical endpoint” (UCEP) line. On the LCEP line, both liquid phases are critical and coexist with a noncritical vapour phase. At the UCEP line, one liquid phase is critical with the vapour phase, and both critical phases coexist with another (noncritical) liquid phase. The right-hand side of Figure 9 shows a pressure versus temperature diagram with both critical endpoint lines, as determined by Aki et al. (39), Zhang et al. (40), and Chobanov (43). The UCEP line lies slightly above the critical line for the vapour-liquid equilibrium of the binary system (CO2 + methanol). This was expected, as the solubility of [C4mim][PF6] in the vapour phase is extremely small (for a better understanding, that solubility is largely out of scale on the left-hand side of Figure 9). The liquid-liquid phase split extends over about 2 MPa at low temperatures, and over about 1 MPa at high temperatures. That pressure range contracts with rising temperature, as both critical endpoint lines converge and end in a tricritical point. There, the three phases (L1, L2, and V) all become critical. For the (CO2 + methanol + [C4mim][PF6]) system, the temperature at that tricritical point seems to lie far beyond the experimentally investigated temperature region, presumably at a temperature where the ionic liquid is no longer stable.


16

CH3OH 12

p

p / MPa

10

8

4

CO2

300

310

320

330

T/K

[bmim][PF6]

[C4mim][PF6]

Figure 9. Phase behaviour of the system (CO2 + methanol + [C4mim][PF6]). Left-hand side: qualitative phase prism at T > Tc (e.g., 313.15 K). Right-hand side: ▲, UCEP; ●, LCEP; Δ, VLE (CO2 + methanol) (43); ◊, UCEP and LCEP (39); x, UCEP and LCEP (40).

LCEP (p = 6.91 MPa)

UCEP (p = 8.16 MPa) 0.8

0.8

L1

UCEP

L2

0.6

0.6

xi


6

0.4

0.4

0.2

0.2

0

8.0

7.8

7.6

7.4

7.2

7.0

0

7.0

7.2

7.4

7.6

7.8

8.0

p / MPa Figure 10. Three-phase liquid-liquid-vapour equilibrium in the system (CO2 + methanol + [C4mim][PF6]) at 313.15 K: Compositions of the coexisting liquid phases L1 and L2 (■, [C4mim][PF6]; ●, methanol; ▲, CO2).


Figure 10 shows the compositions of the coexisting liquid phases in the ternary system (CO2 + methanol + [C4mim][PF6]) at 313.15 K (43). The left-hand side of Figure 10 shows the composition of the ionic liquidrich phase (L1) plotted versus pressure. In that diagram, the pressure increases from right to left. The right-hand side of Figure 10 shows the composition of the liquid phase L2, which is depleted of the ionic liquid – also plotted versus pressure. But in this diagram, the pressure increases from left to right. The three-phase region extends between 6.91 MPa and 8.16 MPa. At the lower critical endpoint, the liquid phase contains approximately 0.08 mol mol–1 [C4mim][PF6] and 0.41 mol mol–1 CO2. With increasing pressure, the mole fraction of [C4mim][PF6] in the phase L1 increases to about 0.25 at the upper critical endpoint. Vice versa, the concentration of [C4mim][PF6] in the liquid phase L2 decreases rapidly with increasing pressure, as at the upper critical endpoint that liquid phase becomes critical with the vapour phase V (which has nearly no solvent power for an ionic liquid). At the upper critical endpoint, the liquid phase L2 contains also very little methanol as the temperature is close to the critical temperature of CO2, and the upper critical endpoint line nearly coincides with the critical line for vapour-liquid equilibrium in the binary system (CO2 + methanol).

TCP (L=V)CO CPCO 2 V CO

L

p


17

UEP 2-water

L1(L2=V)

2

ter wa 2O VC

L2 L1

(L1=L2)V

T Figure 11. Qualitative phase behaviour of the system (CO2 + water + hydrophilic ionic liquid): (– – – –) ternary critical endpoint line; UEP, upper endpoint; TCP, tricritical point (L1 = L2 = V).


18 A liquid-liquid phase split is always observed when a completely miscible mixture of (water + hydrophilic ionic liquid) is pressurised by CO2, as at temperatures below the critical temperature of CO2, the three-phase (L1L2V) line of the binary system (water + CO2) is the upper critical endpoint line of the ternary system (cf. Figure 11). However, for all such systems for which experimental data on this phenomenon are currently available {cf. Ventura et al. (44) and Chobanov (43)}, the SONG phenomenon exists only in a very small pressure range (less than about 0.1 MPa).


Acknowledgements G. M. gratefully acknowledges financial support by Deutsche Forschungsgemeinschaft (DFG), Bonn, Germany, grant no. Ma 713/40 and grant no. Ma 713/45 (within the priority programme DFG-SPP 1191 “Ionic Liquids”).

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