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1. (Liquid + liquid) phase equilibria of 1-alkyl-3-methylimidazolium
methylsulfate with alcohols, or ethers, or ketones
Urszula Domańska a,*, Aneta Pobudkowska a
, Frank Eckert b
a
Faculty of Chemistry, Physical Chemistry Division, Warsaw University of Technology, Noakowskiego 3, 00-664 Warsaw, Poland
b
COSMOlogic GmbH&Co. KG, Burscheider Str. 515, D-51381 Leverkusen, Germany
Received 15 June 2005; received in revised form 26 July 2005; accepted 27 July 2005
Available online 26 September 2005
Abstract
Solubilities of binary mixtures that contain a room-temperature ionic liquid and an organic solvent – namely, 1,3-dimethylim-
idazolium methylsulfate, [mmim][CH3SO4], or 1-butyl-3-methylimidazolium methylsulfate, [bmim][CH3SO4] with an alcohol
(hexan-1-ol, or octan-1-ol, or nonan-1-ol, or decan-1-ol), or an ether (dipropyl ether, or dibutyl ether, or methyl-1,1-dimethylethyl
ether, or methyl-1,1-dimethylpropyl ether), or a ketone (pentan-2-one, or pentan-3-one, or hexan-2-one, or heptan-4-one, or cyclo-
pentanone) – have been measured by a visual method from T = 270 K to the boiling temperature of the solvent. The (liquid + li-
quid) equilibria curves were predicted by the COSMO-RS method. For [bmim][CH3SO4], the COSMO-RS predictions correspond
better with experimental results than do the predictions for [mmim][CH3SO4].
Complete miscibility has been observed in the systems of [mmim][CH3SO4] with water and with alcohols ranging from methanol
to octan-1-ol and that of [bmim][CH3SO4] with water and with alcohols ranging from methanol to decan-1-ol at the temperature
T = 310 K.
Ó 2005 Elsevier Ltd. All rights reserved.
Keywords: Ionic liquids; Experimental (liquid + liquid) equilibria; Molecular interactions; COSMO-RS prediction
1. Introduction
This paper is a continuation of our systematic study
on solubilities of 1-alkyl-3-methylimidazolium salts in
different solvents [1–9]. The measurements of (liquid +
liquid) phase equilibria (LLE) may be helpful for the
modelling of separation processes [10]. Liquid phase
behaviour of 1,3-dialkylimidazolium salts in water and
organic solvents has been studied in many laboratories
[11–18]. The effect of different factors on the phase equi-
libria of imidazolium-based ionic liquids with alcohols
and the partitioning of alcohols between ionic liquids
(IL) and water have been studied [5,7,8,11–18]. In gen-
eral, (IL + organic solvent) binary mixtures show upper
critical solution temperatures (UCSTs). An increase in
the alkyl chain length of alkan-1-ol results in an increase
in the UCST. By increasing the alkyl chain length on the
imidazolium ring, the UCST decreases. ILs show higher
interaction A + B with alcohols than with aromatic
hydrocarbons [4,5].
The COSMO-RS [19], a predictive thermodynamic
model based on unimolecular quantum chemistry calcu-
lations was used to predict LLEs of (IL + alkan-1-ol)
binary mixtures [11,12]. Excellent agreement with the
experimental measurements was obtained. The method
has shown the correct trends for both the variation of
UCST and the alkyl chain length of the alcohol.
Recently, the liquid phase behaviour described by
0021-9614/$ - see front matter Ó 2005 Elsevier Ltd. All rights reserved.
doi:10.1016/j.jct.2005.07.024
*
Corresponding author. Tel.: +48 22 6213115; fax: +48 22 6282741.
E-mail address: ula@ch.pw.edu.pl (U. Domańska).
www.elsevier.com/locate/jct
J. Chem. Thermodynamics 38 (2006) 685–695

2. COSMO-RS of 1,3-dimethylimidazolium methylsulfate,
[mmim][CH3SO4], or 1-butyl-3-methylimidazolium
methylsulfate, [bmim][CH3SO4] in hydrocarbons has
been discussed [20]. Although the quantitative corre-
spondence of the hydrocarbon – IL LLEs with experi-
ment was less satisfying, the shape of the UCST curve
and the trends for the variation of UCST as well as
the size and chain length of the hydrocarbons were pre-
dicted correctly. In this work, the prediction of LLE of
ILs in alkan-1-ols, ethers, and ketones is presented.
The current study focuses on the solubilities of 1,3-
dimethylimidazolium methylsulfate, [mmim][CH3SO4],
or 1-butyl-3-methylimidazolium methylsulfate, [bmim]-
[CH3SO4] in water, in alcohols (hexan-1-ol, octan-1-ol,
nonan-1-ol, and decan-1-ol), in ethers (dipropyl ether,
dibutyl ether, methyl-1,1-dimethylethyl ether, and
methyl-1,1-dimethylpropyl ether), and in ketones (pen-
tan-2-one, pentan-3-one, hexan-2-one, heptan-4-one,
and cyclopentanone).
We would like to show how solubility phenomena in
systems formed by an ionic liquid with organic molecu-
lar solvents can be profitably and surprisingly used in
reaction/extraction processes.
2. Experimental
2.1. Materials
The origin of the chemicals (in parenthesis Chemical
Abstracts registry numbers, and mass fraction purity)
were as follows: [mmim][CH3SO4] (97345-90-9, Solvent
Innovation GmbH, 0.98); [bmim][CH3SO4] (401788-98-
5, Solvent Innovation GmbH, 0.98); hexan-1-ol (111-
27-3, Reachim, 0.99), octan-1-ol (111-87-5, Sigma–Al-
drich, 0.99), nonal-1-ol (143-08-8, Sigma–Aldrich,
0.99), decan-1-ol (112-53-8, Fluka AG, 0.99), dipropyl
ether (111-43-3, Sigma–Aldrich, 0.99), dibutyl ether
(142-96-1, Sigma–Aldrich, 0.99), methyl-1,1-dimethyl-
ethyl ether (1634-04-46, Sigma–Aldrich, 0.995),
methyl-1,1-dimethylpropyl ether (994-05-8, Sigma–Al-
drich, 0.99), pentan-2-one (107-87-9, BDH Chem.
LTD, 0.99), pentan-3-one (96-22-0, Sigma–Aldrich,
0.98), hexan-2-one (591-78-6, Sigma–Aldrich, 0.98),
heptan-4-one (123-19-3, Sigma–Aldrich, 0.98), and
cyclopentanone (120-92-3, Prolabo, 0.99). The ILs were
dried 24 h at T = 330 K in a vacuum before use. All sol-
vents were fractionally distilled over different drying re-
agents to mass fraction purity better than 0.998 and
were stored over freshly activated molecular sieves of
type 4A (Union Carbide). All compounds were checked
by GLC analysis and no significant impurities were
found. Analysis, using the Karl-Fischer technique,
showed that the water content in solvents was less than
0.02 mass%. Physical properties of pure ILs are collected
in table 1.
2.2. Procedure
(Liquid + liquid) phase equilibria (LLE) tempera-
tures have been determined using a dynamic method
that has been described in detail previously [6–8].
Appropriate mixtures of IL and solvent placed under
the nitrogen in dry box into a Pyrex glass cell were
heated very slowly (less than 2 K Æ h1
near the equilib-
rium temperature) with continuous stirring inside a cell.
The sample was placed in a glass thermostat filled with
silicone oil, or water. The temperature of the liquid bath
was varied slowly until one phase was obtained. The
two-phase disappearance temperatures in the liquid
phase were detected visually during an increasing tem-
perature regime. The observation of the regime ‘‘cloud
point’’ with decreasing temperature was very difficult
in these mixtures. The ILs under study often remain in
a metastable state and it is hard to detect the real phase
separation. The observation of the ‘‘cloud point’’ with
decreasing temperature was not repeatable during the
experiment and they have not been recorded here. The
effect of pre-cooling and kinetics of the phenomenon
of binary-phase creation were the reasons that the tem-
perature of ‘‘cloud point’’ was not repeatable. The differ-
ences in observed results were as high as DT = 3 K. The
temperature was measured with an electronic thermom-
eter P 550 (DOSTMANN electronic GmbH) with the
probe totally immersed in the thermostating liquid.
The thermometer was calibrated on the basis of ITS-
90. The accuracy of the temperature measurements
was judged to be ±0.01 K. Mixtures were prepared by
mass, and the errors did not exceed dx1 = 0.0002 and
dT1/K = 0.1 in the mole fraction and temperature,
respectively. The LLE measurements were limited at
the upper temperature by the boiling point of the
solvent.
3. Results and discussion
The solubilities of [mmim][CH3SO4] and [bmim][CH3-
SO4] in alcohols, ethers, and ketones are listed in tables
2–6 and in figures 1–8. The tables include the direct
experimental results of the LLE temperatures, (T/K)
TABLE 1
Physical constants of pure ionic liquids: melting temperature, Tfus,1,
enthalpy of fusion, DfusH1, temperature of glass transition, Ttr,1(g), heat
capacity at half Cp(g) extrapolated temperature, as determined from
DSC data
Ionic liquid Tfus,1/
K
DfusH1/
kJ Æ mol1
Ttr,1(g)/
K
DCp(g)/
J Æ mol1
Æ K1
[mmim][CH3SO4]a
308.90 16.58 185.85 73.72
[bmim][CH3SO4]b
193.78 90.36
a
Reference [20].
b
Reference [6].
686 U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695

3. versus x1, (the mole fraction of the IL at the equilibrium
temperature) for the investigated systems. The maxi-
mum of the curves was recorded for two systems of
[mmim][CH3SO4] in alcohols but for all other systems
was not recorded because in these cases the boiling tem-
perature of the solvent was lower.
The ability of the ionic liquid to form hydrogen
bonds or other possible interactions with potential sol-
vents is an important feature of its behaviour. Pure
[mmim][CH3SO4] and [bmim][CH3SO4] ionic liquids
can act both as hydrogen-bond acceptor ([CH3SO4]
)
and donor ([m- or bmim]+
) and is expected to interact
TABLE 2
Experimental (solid + liquid) and (liquid + liquid) phase equilibria
temperatures for {x1 [mmim][CH3SO4] + (1 x1) an alcohol} systems
x1 T/K
Hexan-1-ola
0.2855 290.35
0.3275 292.19
0.3845 294.53
0.4212 295.36
0.4502 296.00
0.4769 296.80
0.4963 297.24
0.5140 297.73
0.5552 298.76
0.6161 299.89
0.6677 300.76
0.7673 303.46
0.8453 305.05
0.9105 306.84
0.9551 307.98
1.0000 308.90
Heptan-1-ola
0.2414 289.03
0.2731 291.20
0.3371 293.53
0.3783 295.47
0.4497 297.28
0.5021 298.82
0.5630 300.20
0.6348 301.67
0.7247 303.64
0.7997 305.12
0.8967 307.60
0.9599 308.40
1.0000 308.90
Octan-1-ol
0.2780 322.33
0.2960 325.69
0.3140 329.33
0.3296 332.45
0.3555 337.03
0.3873 339.47
0.4159 341.56
0.4409 342.87
0.4425 342.97
0.4439 343.22
0.4465 342.97
0.4537 343.20
0.4704 344.18
0.4716 344.62
0.4858 344.82
0.5047 345.85
0.5792 345.83
0.5827 345.47
0.6276 342.59
0.6680 341.56
0.6948 340.35
0.7440 335.96
0.7550 335.47
0.7872 333.47
0.8072 332.04
0.8469 330.49
0.8675 328.15
0.8949 325.37
TABLE 2 (continued)
x1 T/K
0.9233 323.08
0.9491 319.26
Nonan-1-ol
0.3755 351.23
0.4158 352.83
0.4355 354.26
0.4586 354.49
0.4770 355.42
0.5120 354.50
0.5369 353.76
0.5799 353.48
0.6150 353.82
0.6546 352.38
0.7042 351.71
0.7628 350.57
0.7975 348.92
0.8319 349.63
0.8745 347.57
0.9216 341.35
0.9519 327.80
Decan-1-ol
0.1589 311.36
0.2042 314.24
0.2354 319.14
0.2599 322.25
0.2828 325.63
0.3023 329.25
0.3411 332.76
0.3710 338.10
0.4079 343.28
0.4500 349.18
0.5019 353.45
0.5288 355.85
0.5981 356.55
0.6836 355.45
0.7181 354.65
0.7559 354.65
0.7787 353.33
0.7994 353.98
0.8256 352.75
0.8592 351.37
0.8964 349.19
0.9304 346.65
0.9679 331.13
a
SLE.
U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695 687

4. with solvents which have both accepting and donating
sites. On the other hand, water and alcohols are known
to form hydrogen bonds and ethers and ketones to form
a n + p or p + p interactions with other compounds.
For this discussion, the solubilities of [mmim][CH3-
SO4] and [bmim][CH3SO4] in alcohols ranging from
methanol to decan-1-ol were measured and compared
with other ionic liquids, published by different authors
[5,11–15]. The ability of a polar anion to create the
hydrogen bond with an alcohol significantly increases
the solubilities of IL in solvent. The interaction of the
methylsulfate [CH3SO4]
anion with water and alcohols
is so strong that complete miscibility was observed for
TABLE 3
Experimental (liquid + liquid) phase equilibria temperatures for {x1
[mmim][CH3SO4] + (1 x1) an ether} systems
x1 T/K
Dipropyl ether
0.5698 359.36
0.5865 356.97
0.6000 355.48
0.6081 354.39
0.6209 353.28
0.6409 352.20
0.6520 350.02
0.6680 347.74
0.6720 348.06
0.6969 345.13
0.7167 342.66
0.7394 340.49
0.7509 338.02
0.7653 335.80
0.7820 333.88
0.7977 331.80
0.8228 328.86
0.8381 325.97
0.8454 324.62
0.8507 324.01
0.8726 321.07
0.8968 317.68
0.9187 315.67
0.9417 310.92
0.9631 304.48
Dibutyl ether
0.7662 407.12
0.7795 401.51
0.7888 394.64
0.8115 380.97
0.8435 368.38
0.8571 365.51
0.8725 361.36
0.8851 358.40
0.9056 354.69
0.9286 350.16
0.9362 349.17
0.9614 344.77
0.9800 341.67
MTBE
0.6872 324.13
0.7155 322.95
0.7423 321.87
0.7560 321.33
0.7759 320.38
0.7927 319.76
0.8094 318.56
0.8130 318.17
0.8407 315.68
0.8654 312.43
0.8774 309.71
0.9247 304.12
0.9334 301.88
0.9404 299.52
0.9720 296.71
MTAE
0.9654 302.77
0.9421 316.82
TABLE 3 (continued)
x1 T/K
0.9232 321.23
0.8988 325.57
0.8717 331.16
0.8650 333.13
0.8444 336.71
0.8078 341.62
0.7696 344.82
0.7440 347.98
0.7134 350.66
0.6913 353.17
0.6521 356.76
285
305
325
345
0.0 0.2 0.4 0.6 0.8 1.0
x1
T/K
FIGURE 1. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {[mmim][CH3SO4] (1) + an
alcohol (2)}: m, hexan-1-ol; *, heptane-1-ol; d, octan-1-ol; j, nonan-
1-ol; r, decan-1-ol. Dotted line presents the melting temperature of
[mmim][CH3SO4].
688 U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695

5. [mmim][CH3SO4] in water and alcohols ranging from
methanol to octan-1-ol and for [bmim][CH3SO4] in
water and alcohols ranging from methanol to und-
ecan-1-ol at temperature T = 310 K. Only the binary
systems of [mmim][CH3SO4] with alcohols ranging from
octan-1-ol to undecan-1-ol exhibit UCSTs. The typical
patterns were observed for new anion, that an increase
in the alkyl chain length on the cation increases solubil-
ity in alcohols. The influence of the cations alkyl chain
length was observed previously for [BF4]
and [Tf2N]
[13–15] and [PF6]
anions [5,12].
Experimental phase diagrams of LLE investigated in
this work are characterized mainly by the following:
(1) The solubility of [mmim][CH3SO4] decreases with an
increase of the alkyl chain length of the alcohol from
C7 to C10 (see figure 1); for the longer chain alcohol,
the upper critical solution temperature is shifted
to the higher temperatures and higher IL mole frac-
tion; the (solid + liquid) phase equilibria, SLE was
observed for the hexan-1-ol and heptan-1-ol; the liq-
uidus curves are very flat but the binary liquid phases
were not observed.
(2) An increase in the alkyl chain length of the ether
resulted in an increase in the upper critical solution
temperature (UCST) (shown in figures 2 and 3 for
[mmim][CH3SO4] and [bmim][CH3SO4], respec-
tively); the solubility is lower in dibutyl ether than
in dipropyl ether and is lower in MTAE than in
MTBE for both salts; the best solubility was
observed in the branched chain ether, MTBE; the
difference between the solubility of the two salts is
presented in figure 4 – a smaller binary liquid area
(better solubility) is observed for [bmim][CH3SO4];
it has to be the result of the packing effect and the
interstitial accommodation which is better for the
branched chain ethers (for the longer alkane chain,
e.g. tert-amyl, the solubility is worst); additionally
the inductive effect of the alkane chains of the
branched chain ethers causes the stronger interaction
of the oxygen with the ionic liquid.
(3) The solubility of ILs in cyclic molecules is usually
greater than in long chain molecules, this was
observed previously for salts under investigation in
cycloalkanes in comparison with n-alkanes [20]; the
solubility of [mmim][CH3SO4] in cyclopentanone is
much greater than in other ketones: pentan-2-one,
pentan-3-one, hexan-2-one, and heptane-2-one (see
figure 5 and table 5); the cyclic molecule of the imid-
azole is more soluble in the cyclic solvent; the posi-
tion on the carbonyl group does not make any
difference in the solubility; the other equilibrium
280
300
320
340
360
380
400
420
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 2. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {[mmim][CH3SO4] (1) + an
ether (2)}: s, dipropyl ether; h, dibutyl ether; n, MTBE; e, MTAE;
dotted lines, boiling temperatures of the solvents; lines are from
prediction by the COSMO-RS calculation: ( ), dipropyl ether;
(––), dibutyl ether; (–––), MTBE; and (3), MTAE.
250
300
350
400
450
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 3. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {[bmim][CH3SO4] (1) + an
ether (2)}: d, dipropyl ether [6]; j, dibutyl ether; m, MTBE; r,
MTAE; dotted lines, boiling temperatures of the solvents; lines are
from prediction by the COSMO-RS calculation: ( ), dipropyl
ether; (––), dibutyl ether; (–––), MTBE; and (3), MTAE.
U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695 689

6. curves in ketones exhibit similar shapes; the solubil-
ity decreases with an increase of the molar mass of
the ketone as it is shown in figures 5 and 6; the results
comparing two salts in pentan-3-one and hexan-2-
one are presented in figures 7 and 8, respectively;
as in the case of ethers, the experimental equilibrium
ILs mole fractions is lower for [bmim][CH3SO4] than
for [mmim][CH3SO4] in each of the ketones.
(4) The mutual (liquid + liquid) solubility of
[mmim][CH3SO4] and [bmim][CH3SO4] increases in
the order an alcohol an ether ketone.
(5) The effect of the cations, [m- or bmim]
on the solu-
bility of these salts in the tested solvents is the result
of the (A + B) hydrogen bonding and (n + p) inter-
action between the imidazolium ring and the solvent.
The stronger interaction on one side and the
increased influence of the inductive effect of the
longer alkyl chain on the imidazole ring on the other
side have a significant influence on the solubility.
This effect was explained by the increasing van der
Waals interactions between the alkyl chain of the
cation and the alkyl chain of the alcohol [14]. Much
greater solubility of [bmim][CH3SO4] than of
[mmim][CH3SO4] was also observed in cycloalkanes
and aromatic hydrocarbons [20].
For the investigated mixtures, it was impossible to de-
tect by the visual method the mutual solubility of ILs in
the solvent-rich phase. These data can be approximated
by the COSMO-RS predictions.
4. Prediction of the (liquid + liquid) phase equilibria
The Conductor-like Screening Model for Real Sol-
vents (COSMO-RS) is a unique method for predicting
the thermodynamic properties of mixtures on the basis
of unimolecular quantum chemical calculations for the
individual molecules [20–23]. COSMO-RS combines
the electrostatic advantages and the computational effi-
ciency of the quantum chemical dielectric continuum
solvation model COSMO [24] with a statistical thermo-
dynamics method for local interaction of surfaces, which
takes into account local deviations from dielectric
behaviour as well as hydrogen bonding. In this ap-
proach, all information about solutes and solvents is ex-
tracted from initial quantum chemical COSMO
calculations, and only very few parameters have been fit-
ted to experimental values of partition coefficients and
290
310
330
350
370
390
410
430
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 5. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {[mmim][CH3SO4] (1) + ketone
(2)}: s, pentan-2-one; h, pentan-3-one; e, hexan-2-one; n, heptan-4-
one; , cyclopentanone; dotted line, boiling temperature of a solvent;
lines are from prediction by the COSMO-RS calculation: (3), pentan-
2-one; (–––), pentan-3-one; ( ), hexan-2-one; (––), heptan-4-one;
and (– –), cyclopentanone.
310
330
350
370
390
410
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 4. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {ILs (1) + dibutyl ether (2)}: h,
[mmim][CH3SO4]; j, [bmim][CH3SO4]; lines are from prediction by
the COSMO-RS calculation: (3), [mmim][CH3SO4]; and (–––),
[bmim][CH3SO4].
690 U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695

7. vapour pressures for a wide range of neutral organic
compounds. The COSMO-RS is capable of predicting
activity coefficients, partition coefficients, vapour pres-
sures, and solvation free energies of neutral compounds
with an error of 0.3 log-units (rms) and substantial expe-
rience has been gathered during the past years about its
surprising ability to predict mixture thermodynamics
[20–23].
In addition, successful applications of COSMO-RS
to the predictions of the thermodynamic properties of
ions in solution and ionic liquids have been reported
[12,25,26]. We applied the standard procedure for COS-
MO-RS calculations, which consists of two steps, viz: (1)
quantum chemical COSMO calculation for all molecu-
lar species involved and (2) COSMO-RS calculations.
4.1. Quantum chemical COSMO calculation for all
molecular species involved
In these calculations, the solute molecules are calcu-
lated in a virtual conductor environment. In such an
environment, the solute molecule induces a polarization
charge density on the interface between the molecule
and the conductor, that is, on the molecular surface.
These charges act back on the solute and generate a
more polarized electron density than in a vacuum.
250
270
290
310
330
350
370
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 7. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {ILs (1) + pentan-3-one (2)}:
h, [mmim][CH3SO4]; j, [bmim][CH3SO4]; lines are from prediction by
the COSMOtherm calculation: (3), [mmim][CH3SO4] and (–––),
[bmim][CH3SO4].
250
300
350
400
450
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 6. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {[bmim][CH3SO4] (1) + ketone
(2)}: d, pentan-2-one; j, pentan-3-one; r, hexan-2-one; m, heptan-4-
one; dotted line, boiling temperature of a solvent; lines are from
prediction by the COSMO-RS calculation: (–––), pentan-3-one;
( ), hexan-2-one; and (––), heptan-4-one.
250
290
330
370
410
0 0.2 0.4 0.6 0.8
x1
T/
K
1
FIGURE 8. Plot of temperature against mole fraction component 1
for the phase diagrams for the systems {ILs (1) + hexan-2-one (2)}: e,
[mmim][CH3SO4]; r, [bmim][CH3SO4]; lines are from prediction by
the COSMO-RS calculation: (3), [mmim][CH3SO4] and (–––),
[bmim][CH3SO4].
U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695 691

8. During the quantum chemical self-consistency cycle, the
solute molecule is thus converged to its energetically
optimal state in a conductor with respect to electron
density. The quantum chemical calculation has to be
performed once for each molecule of interest. The geom-
etry of all compounds was fully relaxed at the density
TABLE 4
Experimental (liquid + liquid) phase equilibria temperatures for {x1
[bmim][CH3SO4] + (1 x1) an ether} systems
x1 T/K
Dipropyl ethera
0.7366 356.05
0.7772 340.95
0.8195 330.27
0.8684 315.75
0.8964 306.05
Dibutyl ether
0.6816 407.85
0.6904 398.18
0.7053 392.06
0.7175 385.78
0.7307 380.42
0.7439 374.37
0.7545 369.73
0.7658 363.62
0.7999 349.37
0.8280 337.66
0.8688 321.88
0.8821 317.59
0.8974 311.66
MTBE
0.6844 352.10
0.6947 346.95
0.7065 339.59
0.7164 329.55
0.7219 323.20
0.7295 313.76
0.7342 308.72
0.7386 303.17
0.7414 297.78
0.7449 292.44
MTAE
0.4200 467.14
0.4481 465.33
0.4731 461.71
0.4892 460.39
0.5073 458.10
0.5515 448.57
0.5707 443.80
0.6525 421.87
0.6782 415.29
0.7168 396.50
0.7506 381.71
0.8057 363.59
0.8546 342.51
0.8865 324.08
0.9251 297.06
a
From reference [6].
TABLE 5
Experimental (liquid + liquid) phase equilibria temperatures for {x1
[mmim][CH3SO4] + (1 x1) ketone} systems
x1 T/K
Pentan-2-one
0.5846 368.81
0.6498 365.90
0.6636 364.63
0.6922 361.96
0.7213 356.61
0.7457 352.02
0.7819 344.07
0.8118 330.60
0.8455 320.92
0.8909 311.04
0.9465 305.40
Pentan-3-one
0.6186 367.51
0.6556 364.49
0.6888 361.77
0.7128 359.18
0.7350 355.80
0.7615 348.34
0.7789 343.36
0.7910 339.51
0.8066 334.83
0.8295 327.53
0.8526 322.09
0.8871 312.85
0.9402 306.96
Hexan-2-one
0.6758 394.01
0.7020 389.37
0.7386 381.51
0.7828 374.57
0.8079 367.71
0.8336 360.97
0.8386 359.60
0.8569 354.59
0.8717 346.17
0.8859 338.91
0.9146 325.93
0.9554 296.99
Heptan-4-one
0.7473 405.46
0.7748 399.01
0.7915 393.39
0.8208 385.38
0.8413 380.40
0.8646 373.28
0.8799 367.62
0.8977 358.97
0.9153 349.64
0.9408 338.76
0.9746 322.38
Cyclopentanone
0.3944 396.80
0.4337 387.43
0.4522 382.27
0.4722 376.39
0.4925 371.32
0.5164 364.98
692 U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695

9. functional (DFT) level of theory [27]. This was done
with the Turbomole programme package [27,28] using
the B-P density functional theory with a TZVP quality
basis set [27] and the RI approximation [29]. During
these calculations, the COSMO continuum solvation
model was applied in the conductor limit (e = 1). Ele-
ment-specific default radii from the COSMO-RS param-
eterizations have been used for the COSMO cavity
construction [19,21]. Such calculations end up with the
self-consistent state of the solute in the presence of a vir-
tual conductor that surrounds the solute outside the
cavity.
4.2. COSMO-RS calculations
These have been done using the COSMOtherm pro-
gramme [30]. In these calculations, the deviations of a
real solvent, in our case ionic liquids and organic sol-
vents, compared with an ideal conductor are taken into
account in a model of pair-wise interacting molecular
surfaces. For this purpose, electrostatic energy differ-
ences and hydrogen bonding energies are quantified as
functions of the local COSMO polarization charge den-
sities r and r0
of the two interacting surface pieces. The
chemical potential differences arising from these interac-
tions are evaluated using an exact statistical thermody-
namics algorithm for independently pair-wise
interacting surfaces, which is implemented in COS-
MOtherm. More detailed descriptions of the COSMO-
RS method are given elsewhere [19,21–23]. The
COSMO-RS method depends on a small number of
adjustable parameters, some of which are predetermined
from physics. The others are determined from selected
properties of mixtures, with none of the mixtures related
to the mixtures under study here. The parameters are
not specific to functional groups or type of molecule.
As the parameters are used as a basis for the calculations
only, the resulting parameterization [19] is completely
general. No special adjustments of radii or other param-
eters have been made for this application.
If more than one conformation were considered to be
potentially relevant for a compound, several conforma-
tions have been calculated in step 1 and a thermody-
namic Boltzmann average over the total Gibbs free
energies of the conformers was consistently calculated
by the COSMOtherm programme in step 2. The COS-
MO-RS0
statistical thermodynamics procedure results
in the chemical potential of all components of a given
mixture. Thermodynamic properties of the mixture,
such as activity coefficients and (vapour + liquid) equi-
libria can be derived from the chemical potentials. The
(liquid + liquid) equilibrium properties have been calcu-
lated from the equity
xI
i cI
i ¼ xII
i cII
i ; ð1Þ
where indices I and II denote the liquid phases, xi are the
mole fractions of the two solvents and ci are the activity
coefficients of the solvents as computed by COSMO-RS.
The predicted solubilities of [mmim][CH3SO4] and
[bmim][CH3SO4] in every investigated solvent are pre-
sented in figures 2–8. The required COSMO-RS param-
eters, used in the calculations and the methodology
details were presented together with the theory in the
previous paper [19].
In this work, the COSMO-RS predictions correspond
better with experimental results for [bmim][CH3SO4]
TABLE 5 (continued)
x1 T/K
0.5397 358.17
0.5685 349.63
0.6031 336.33
TABLE 6
Experimental (liquid + liquid) phase equilibria temperatures for {x1
[bmim][CH3SO4] + (1 x1) ketone} systems
x1 T/K
Pentan-2-one
0.4323 377.09
0.4388 365.50
0.4499 350.92
0.4566 339.82
0.4646 327.30
0.4704 318.83
0.4747 313.81
0.4818 302.42
0.4884 293.35
0.4982 281.20
Pentan-3-one
0.4320 372.70
0.4435 360.52
0.4510 349.70
0.4652 326.60
0.4761 313.19
0.4782 306.70
0.4861 292.20
0.5022 276.50
Hexan-2-one
0.5182 436.17
0.5395 399.32
0.5544 373.48
0.5759 345.10
0.5881 330.79
0.6071 305.90
0.6184 294.23
Heptan-4-one
0.7196 451.30
0.7306 420.37
0.7649 385.06
0.7939 363.00
0.8259 322.61
0.8892 295.76
U. Domańska et al. / J. Chem. Thermodynamics 38 (2006) 685–695 693

10. than for [mmim][CH3SO4]. This can be explained partly
by the stronger polarity of [mmim][CH3SO4], which re-
sults in possibly larger errors in COSMO-RS predic-
tions. The COSMO-RS overestimates the polarity
difference of the compounds and thus the miscibility
gap prediction is too wide and the UCST is too high.
This is especially true for the alcohols in this work. As
one would expect, this overestimation is a result of
strong (A + B) interaction for the investigated mixtures.
The interaction is due to four atoms of oxygen in the an-
ion, which is much stronger than that observed for mix-
tures with tetrafluoroborate or hexafluorophosphate
salts. If the comparison is made for the curves below
the boiling point of the solvent, it looks hardly accept-
able, but the trend and character of the equilibrium
curves is predictable. Calculations are very helpful in
attempting to predict the solvent-rich side equilibrium
curves. In most of the figures, the order of the predicted
curves for certain solvents is correct with exception of
dipropyl ether and MTAE for both salts. However, the
COSMO-RS predictive curves show significant disagree-
ment with the experimental findings, but the character of
calculated lines is similar to the experimental one. The
calculated curves are especially useful in the solvent rich
mixtures (left side of the figures), where the experiment
was not successful (tables 3–6 and figures 2–8).
It can be noticed as well that for the ketones, the
COSMO-RS prediction was the best (see figures 7 and
8). Generally, we recently found that the [CH3SO4]
an-
ion is extremely polar and that prediction by COSMO-
RS is very difficult. Probably due to the increasing
impact of the long-range interactions which, to date,
cannot be described satisfactorily by COSMO-RS, or
the modified COSMO-RS(Ol) model [31,32]. These re-
sults of the prediction do not warrant discussion about
the standard deviation of the prediction, because for
the measured ILs only the qualitative agreement was
obtained.
5. Conclusion
Knowledge of the impact of different factors on the
liquid phase behaviour of IL with other liquids is useful
for developing IL as designer solvents. The size of the
area of the two liquid phases decreases with an increase
in the chain length of the alkyl substituent at the imidaz-
ole ring. The observations of upper critical solution tem-
peratures (UCSTs) were limited by the boiling
temperature of the solvent. The observation of the (li-
quid + liquid) de-mixing at the solvent-rich phase was
inhibited by the permanently foggy solution.
The anions play a major role in the solubility of ILs
and miscibility with polar solvents, as can be seen from
the complete miscibility of [bmim][CH3SO4] with water
and alcohols in comparison with [bmim][PF6] or [bmim]
[BF4]. The cations play a secondary role; the solubility
of methyl- versus butyl-methyl-imidazolium cation in
different solvents, results in only a 10% to 20% difference
in solubility. We can tentatively explain the difference in
the solubility by the suggestion that the longer alkyl
chain in the imidazolium ring creates more free volume
in the solution, thus allowing solvent molecules to reach
their preferred sites of interaction more easily.
These results can be easily used for the chemical reac-
tions with a specific ionic liquid as a reaction media. If
one uses an extractant solvent to separate reactants
and products, the phase behaviour of the extractant with
the ionic liquid must be known.
Acknowledgements
This research has been supported by the Warsaw
University of Technology. Authors thank Mrs. A. Wis-
niewska for some solubility measurements.
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