Heat capacities and excess enthalpies of 1-3ethyl-3-methylimidazolium-based ionic liquids and water

Article abstract of DOI:10.1021/je800248w

Heat capacities and excess enthalpies were determined for three different binary water + ionic liquid systems, from (283.15 to 348.15) K, and covering the entire composition range. Specifically, the three completely water-miscible ionic liquids used were 1-ethyl-3-methylimidazolium ethylsulfate, 1-ethyl-3-methylimidazolium trifluoromethanesulfonate, and 1-ethyl-3- methylimidazolium trifluoroacetate. The influence of temperature and composition was assessed, and suitable equations were used to correlate the experimental data. In addition, it was found that 1-ethyl-3-methylimidazolium ethylsulfate decomposes in the presence of water to form 1-ethyl-3-methylimidazolium hydrogen sulfate and ethanol under ambient conditions.

Full text of DOI:10.1021/je800248w

2112
J. Chem. Eng. Data 2008, 53, 2112–2119
Heat Capacities and Excess Enthalpies of 1-Ethyl-3-methylimidazolium-Based
Ionic Liquids and Water
Lindsay E. Ficke, He´ctor Rodr´ıguez,† and Joan F. Brennecke*
Department of Chemical and Biomolecular Engineering, University of Notre Dame,
182 Fitzpatrick Hall, Notre Dame, Indiana 46556
Heat capacities and excess enthalpies were determined for three different binary water + ionic liquid systems,
from (283.15 to 348.15) K, and covering the entire composition range. Specifically, the three completely
water-miscible ionic liquids used were 1-ethyl-3-methylimidazolium ethylsulfate, 1-ethyl-3-methylimidazolium
trifluoromethanesulfonate, and 1-ethyl-3-methylimidazolium trifluoroacetate. The influence of temperature
and composition was assessed, and suitable equations were used to correlate the experimental data. In addition,
it was found that 1-ethyl-3-methylimidazolium ethylsulfate decomposes in the presence of water to form
1-ethyl-3-methylimidazolium hydrogen sulfate and ethanol under ambient conditions.
Table 1. IL Structures and Abbreviations
Introduction
Ionic liquids (ILs) are a family of low melting salts which
have generated burgeoning interest in industry and academia
over the past decade. The appealing set of properties exhibited
by these salts, which are liquid at or near room temperature,
suggests the possible use of ILs in a wide variety of applica-
tions.1 Since many ILs are entirely miscible with water, the use
of water + IL mixtures as working fluids in processes is
possible. Introducing an IL as a replacement may improve the
efficiency or decrease the environmental impact of conventional
processes involving aqueous solutions of salts or other soluble
substances. An example of particular interest is absorption
refrigeration.
The development of practical processes involving water +
exothermic enthalpies of mixing for water + [bmpy][BF4] over
IL systems will require basic thermodynamic data, which are
the entire composition range at (298.15 and 318.15) K.32 Rebelo
et al. have reported endothermic excess enthalpy values for water
generally lacking for water + IL systems. For thermal applica-
tions, one of the properties of interest is the heat capacity. To
+ [bmim][BF4] over the entire composition range from (278.15
date, several publications have reported experimental values of
to 333.15) K.6 Katayanagi et al. reported excess partial molar
heat capacity for pure ILs as a function of temperature.2–24
enthalpies for water + two ILs ([bmim][BF4] and 1-butyl-3-
However, there are only a few papers reporting heat capacities
of a binary mixture including water + IL (1-butyl-3-methylimi-
dazolium tetrafluoroborate, [bmim][BF4], 1-butyl-3-methylpy-
ridinium tetrafluoroborate, [bmpy][BF4], and 1-butyl-3-meth-
ylimidazolium tosylate, [bmim][Tos]).6,20,21
The excess enthalpy is another fundamental property for
which literature data for water + IL mixtures over the entire
composition range and at multiple temperatures are sparse.
Several groups have reported the excess enthalpies at very low
composition ranges and room temperature for water + tetrafluo-
roborate ([BF4]-), + chloride ([Cl]-), + bis(trifluoromethyl-
sulfonyl)imide ([Tf2N]-), and + ammonioacetate ([Gly]-) based
ILs.11,25–30 Constantinescu et al. have reported exothermic excess
enthalpies for water + two ILs ([choline][lactate] and [cho-
line][glycolate]) over the entire composition range and temper-
atures of (303.15, 315.15, and 323.15) K.31 Ortega et al. reported
methylimidazolium iodide, [bmim][I]) over the entire composi-
tion range at 198.15 K.33 No excess enthalpy data have been
reported for the systems investigated in this study.
In this work, heat capacities and excess enthalpies are reported
for binary systems of water and three different hydrophilic and
noncorrosive ILs, with melting points below that of water:
1-ethyl-3-methylimidazolium ethylsulfate, [emim][EtSO4]; 1-ethyl-
3-methylimidazolium trifluoromethanesulfonate (or triflate),
[emim][OTf]; and 1-ethyl-3-methylimidazolium trifluoroacetate,
[emim][TFA]. The structures and abbreviations of the ILs are
shown in Table 1. In addition to the characteristics aforemen-
tioned, these ILs were selected because they were believed to
be more chemically stable than some other ILs,25,34 they have
relatively low viscosities,35 and they have a common cation.
The experimental determinations of heat capacity and excess
enthalpy for the systems water + [emim][EtSO4], water +
[emim][OTf], and water + [emim][TFA] covered the entire
composition range and a wide temperature range, from (283.15
to 348.15) K.
* Corresponding author. Telephone: (574)631-5847. Fax: (574)631-8366.
E-mail: jfb@nd.edu.
† Current Address: The QUILL Centre, School of Chemistry and Chemical
Engineering, The Queen’s University of Belfast, Belfast, BT9 5AG,
Northern Ireland, U.K.
10.1021/je800248w CCC: $40.75
2008 American Chemical Society
Published on Web 08/09/2008

Journal of Chemical & Engineering Data, Vol. 53, No. 9, 2008 2113
were determined by the sapphire method, extensively used in
the literature and particularly with ILs.3,5,7–10,14,22 This method
is based on the comparison of the DSC signal of the sample
with the DSC signal of a calibration sample of known heat
capacity (sapphire). Both curves must be blank corrected.
Therefore, for each measurement, a minimum of three runs were
carried out: the blank, with an empty crucible; the sapphire;
and the sample itself. In addition, a cleaning process at 600 °C
for 10 min, followed by one or two extra blank curves to
condition the measuring cell, was run prior to each of the
measurement sequences. The uncertainty attained by the sap-
phire method is better than ( 3 %; however, taking into account
the uncertainties in weighing involved, an overall uncertainty
of ( 4 % is estimated for the values reported in this work.
The temperature program used for the blank, sapphire, and
sample runs in the DSC included a 5 min 278.15 K isothermal
stage, followed by a temperature ramp from (278.15 to 348.15)
K at a rate of 10 K ·min-1 and a final 5 min isothermal segment
at 348.15 K. From this data, heat capacity values were calculated
at every 5 K between (283.15 and 343.15) K.
Experimental
Chemicals. The ILs used in the heat capacity experiments
were synthesized in our laboratory. The structures of the final
products were checked by nuclear magnetic resonance (NMR)
spectroscopy.
The IL [emim][EtSO4] was synthesized by direct reaction of
equimolar amounts of 1-methylimidazole (Aldrich, 99 %, freshly
redistilled over KOH) and diethylsulfate (Aldrich, 98 %),36
adding the latter dropwise to the 1-methylimidazole in a round-
bottomed flask, placed in an ice bath to keep the temperature
low. Removal of any unreacted starting material was carried
out by rotary evaporation and heating under high vacuum (<67
Pa) for two days. The water mass fraction of the final product
was 69 ·10-6, measured by Karl Fischer titration with a
Metrohm 831 KF Coulometer.
In a similar way, ethyltrifluoromethanesulfonate (Aldrich, 99
%) and 1-methylimidazole were reacted at low temperature in
a stirred round-bottomed flask with a reflux condenser, in the
absence of any solvent, to form [emim][OTf]. A similar
procedure can be found in the literature.37 Unreacted starting
materials were removed by heating under reduced pressure. The
The method was tested for water, and the results were found
to be in good agreement, within the experimental error, with
values in reference literature.38
water mass fraction of the final liquid was found to be 44 ·10-6
,
measured by Karl Fischer titration.
Excess Enthalpy Measurements. Excess enthalpies of water
+ IL were obtained by calorimetry, using a Setaram C80
calorimeter. A DSC measures the heat flow to or from
the sample vessel compared to an empty reference vessel,
as the result of a chemical reaction, change of state, or mixing.
The temperatures of the vessels are maintained at the same
temperature of the calorimetric block by Peltier elements. The
calorimeter used here has high accuracy due to three-
dimensional heat flow transducers completely surrounding the
experimental vessels.
The stainless steel C80 mixing vessels were dried in an oven
at 388.15 K prior to sample preparation. Water was placed in
the bottom portion of the vessel and IL in the top. The mixing
occurred using continuous reversal which consisted of rotating
the entire vessel 180°. The bottom section of the vessel had a
stainless steel lid that fell off with the 180° turn and served as
a “stirrer”. In addition, mercury was used to completely separate
the two liquids prior to mixing.
The ILs were saturated with mercury prior to the experiments
by mixing for 22 h at 323.15 K. The mass fraction of mercury
which dissolved or remained suspended in the IL was less than
50 · 10-6, which was confirmed by analysis with inductively
coupled plasma optical emission spectroscopy (ICP-OES). In
addition, the difference in the excess enthalpies of the mercury
saturated IL and a clean IL was small and within the experi-
mental error.
The measurements were performed at atmospheric pressure
including temperatures from (313.15 to 348.15) K and over the
entire range of mole fraction (0.25, 0.50, and 0.75) as determined
by mass. The vessels were loaded in a dry nitrogen environment
using a Mettler Toledo XS205DU analytical balance accurate
to ( 1·10-5 g. The water content of the ILs was measured by
Karl Fischer titration.
The procedure for the synthesis of [emim][TFA] was similar
to that reported elsewhere.37 Equimolar amounts of 1-ethyl-3-
methylimidazolium bromide (Fluka, 97.0 %), further purified
in our laboratory prior to use, and silver trifluoroacetate (Aldrich,
98 %), used as received, were dissolved in water. Both solutions
were mixed and allowed to react for 24 h in a setup similar to
that described for the synthesis of [emim][OTf]. The resulting
solution was filtered to remove the AgBr precipitate, and water
was partially removed in a rotary evaporator. The presence of
bromide and silver ions in the product was tested by addition
of AgNO3 and NaCl solutions to probe vials containing aliquots
of the reaction product. According to the precipitates observed
in the tests, an additional amount of one of the starting materials
was added to neutralize the excess of the other. The reaction
and subsequent steps were repeated until no precipitate was
observed in either of the tests, indicating bromide and silver
mass fractions less than 1 ·10-3. The purification was completed
by drying the IL thoroughly under high vacuum at moderate
temperature for several days. A water mass fraction of 678 ·10-6
was assessed by Karl Fischer titration.
The ILs used for the excess enthalpy measurements were
purchased from Solvent Innovations ([emim][EtSO4] and [emim]-
[OTf]) and Merck KGaA ([emim][TFA]), all of which had a
purity of > 99 %, and the samples were used as received except
for drying. The average water mass fraction for the ILs was
(152, 157, and 64) ·10-6, respectively. High-purity water,
deionized through a Milli-Q Water System, was used for all
the experiments.
Heat Capacity Measurements. Heat capacity measurements
were performed for all three water + IL mixtures over a range
of temperatures and compositions. The mixtures were prepared
by mass using a Mettler Toledo AE 160 balance, precise to
within ( 1 · 10-4 g. Good mixing was ensured by magnetic
stirring. A sample of ca. (20 to 40) mg was placed in an
aluminum pan, covered with a lid of the same material, and
hermetically sealed. All the samples were prepared immediately
prior to performing the measurements to avoid variations in
composition due to evaporation or sorption of water.
The method included three segments: a dynamic segment
ramping up to the desired temperature; a three hour isotherm
prior to mixing followed by a one hour isotherm after mixing;
and a dynamic segment ramping down to room temperature.
The mixing itself occurred for 69 s.
Measurements were performed in a differential scanning
calorimeter (DSC), manufactured by Mettler-Toledo, model
DSC822e, with STARe software data collection. Heat capacities
The experimental uncertainty associated with this method is
estimated to be 2 % considering the propagation of error and
the ability to reproduce measurements with different samples

2114 Journal of Chemical & Engineering Data, Vol. 53, No. 9, 2008
Table 2. Fit Parameters a and b, Equation 1, and Standard
Deviation, σ, for the Empirical Correlation of Molar Heat Capacity
in Binary Systems Water (1) + IL (2) as a Function of
Temperature, in the Range from (283.15 to 343.15) K
x1
a/102 J·mol-1 ·K-1
b/104 J·mol-1
σ/J·mol-1 ·K-1
Water + [emim][EtSO4]
0.0000
0.2120
0.4079
0.5925
0.7660
0.8495
0.9282
0.9684
0.9916
1.0000
5.827
-6.161
-4.899
-4.087
-3.862
-2.740
-2.178
-1.368
-0.888
-0.303
-0.096
0.5
0.8
0.4
0.4
0.5
0.5
0.3
0.4
0.3
0.2
4.656
3.792
3.164
2.390
1.929
1.433
1.150
0.886
0.783
Water + [emim][OTf]
-5.166
0.0000
0.2259
0.4351
0.6152
0.7835
0.8614
0.9353
0.9712
0.9924
1.0000
5.351
4.454
3.553
2.745
2.146
1.771
1.274
1.086
0.926
0.783
0.8
0.5
0.5
0.5
0.3
0.3
0.3
0.2
0.3
0.2
-4.622
-3.651
-2.636
-2.270
-1.789
-0.931
-0.796
-0.458
-0.096
Water + [emim][TFA]
-4.583
0.0000
0.2025
0.3973
0.5817
0.7562
0.8420
0.9256
0.9668
0.9912
1.0000
4.699
4.021
3.846
3.089
2.396
1.818
1.326
1.114
0.914
0.783
0.5
0.9
0.7
0.3
0.7
0.3
0.3
0.8
0.5
0.2
-3.907
-4.480
-3.600
-2.819
-1.816
-0.935
-0.636
-0.346
-0.096
for [emim][OTf] and [emim][EtSO4] agree with those reported
by Diedrichs and Gmehling14 and Zhang et al.18 within
experimental uncertainty. Ferna´ndez et al.19 reported a heat
capacity of [emim][EtSO4] at 298.15 K (421.1 J ·mol-1 ·K-1
)
Figure 1. Molar heat capacity, Cp, for the binary systems water (1) + IL
(2) as a function of temperature, at different approximate mass fractions of
water: b, 0 %; O, 2 %; 1, 5 %; 3, 10 %; 9, 20 %; 0, 30 %; 2, 50 %; ∆,
70 %; [, 90 %; ], 100 %. Empirical correlations are plotted as solid lines.
ILs: (a) [emim][EtSO4]; (b) [emim][OTf]; (c) [emim][TFA].
which is higher than the values obtained in both this study and
by Zhang et al.18 [(377 and 378) J ·mol-1 ·K-1, respectively].
This can be attributed to the purity (Sigma-Aldrich g95 %)
and high water mass fraction (2·10-3) of the IL used by
Ferna´ndez et al.19
Most correlations of experimental heat capacity data as a
function of temperature in the literature consist of empirical
polynomial equations (often just straight lines).2,3,6–10,13,22 For
the three binary systems in this work, we found excellent fit to
the experimental data with the following equation
over time. The method was verified using organic + organic
systems (i.e., benzene + cyclohexane, 1-propanol + propyl
acetate, and pyridine + acetic acid) at the extreme temperatures,
both (313.15 and 348.15) K, which produced results in agree-
ment with the literature and within the experimental uncer-
tainty.39
b
Cp ) a +
(1)
T
Results and Discussion
Heat Capacity: Effect of Temperature. Molar heat capacities
for water + IL systems, from pure IL to pure water, are plotted
in Figure 1 as a function of temperature. The experimental data
can be found in Tables S1 to S3 of the Supporting Information.
A slight increase of Cp with increasing temperature is observed
for IL-rich compositions, which diminishes for water-rich
samples. From Figure 1, it is clear that the Cp of the pure ILs
is much greater than that of water. This is not surprising, since
Cp is related to the number of translational, vibrational, and
rotational energy storage modes in the molecule,7 and the
molecular weight of the ILs is much greater than water. The
heat capacity of the mixtures decreases with increasing water
mole fraction. Our measurements of the pure IL heat capacities
where T is the absolute temperature and a and b are the fit
parameters. This equation has been used previously to
correlate this type of data for pure ILs.14 Least-squares fits
led to the values summarized in Table 2. The low values of
the corresponding standard deviations, σ, also shown in Table
2, confirm the ability of the proposed equation to provide a
good correlation of the experimental data, with only two
adjusting parameters. Although the lowest standard deviations
are for pure water, this is because the absolute value of Cp
is the lowest for pure water. The empirical correlation
presented here is not intended to provide an accurate
description of Cp of pure water itself but just a suitable
correlation for the binary systems water + IL throughout the

Journal of Chemical & Engineering Data, Vol. 53, No. 9, 2008 2115
Table 3. Specific Heat Capacities (cp) of the Pure ILs at Selected
Temperatures
cp/J·g-1 ·K-1
T/K
[emim][EtSO4]
[emim][OTf]
[emim][TFA]
283.15
293.15
303.15
313.15
323.15
333.15
343.15
1.54
1.57
1.61
1.63
1.66
1.68
1.71
1.36
1.37
1.41
1.42
1.44
1.46
1.48
1.37
1.40
1.43
1.44
1.46
1.48
1.50
whole compositional spectrum. The resulting correlations are
plotted with solid lines in Figure 1.
In Table 3, the specific heat capacities, cp, are shown for the
three pure ILs over the temperature range studied. The cp values
for ILs are much lower than those for water, 4.18 J ·g-1 ·K-1
at 293.15 K.40 This means less energy will be required to
produce a certain temperature increase in a given mass of IL
than in the same mass of water.
The relative value of the heat capacities of the IL changes
depending on whether a mole or mass basis is considered. Of
the three ILs studied, the Cp values are highest for [emim]-
[EtSO4], followed closely by [emim][OTf], but [emim][TFA]
is significantly lower. [emim][EtSO4] also has the highest cp
values, but [emim][OTf] and [emim][TFA] have very similar
cp values over the temperature range studied.
Heat Capacity: Effect of Composition. The variation of Cp
with composition is shown in Figure 2 for the highest and lowest
isotherms of the three water + IL binary systems. As expected,
Cp decreases with increasing mole fraction of water in the
E
mixture. The excess molar heat capacity, Cp , is the difference
between the mixture heat capacity and the pure components
CpE ) C -
x · C
p,i
(2)
∑
p
i
where Cp is the molar heat capacity of the mixture; Cp,i is
the molar heat capacity of the pure compound; and xi is its
mole fraction. The values of Cp for the three systems studied
E
Figure 2. Molar heat capacity, Cp, for the binary systems water (1) + IL
(2) as a function of mole fraction of water, at different temperatures: b,
283.15 K; O, 343.15 K. Linear fits are plotted as solid lines. ILs: (a)
[emim][EtSO4]; (b) [emim][OTf]; (c) [emim][TFA].
here are very small, on the same order of magnitude as the
experimental uncertainty, so no significant analysis can be
carried out. This means that an assumption of a linear decrease
E
of Cp with the concentration of water (i.e., Cp ≈ 0) is a very
The water + [emim][EtSO4] mixture (Figure 3a) is exother-
mic. These negative values of the excess enthalpy suggest that
the water/IL interactions are stronger than the corresponding
IL/IL and water/water interactions. It is important to note that
at the lower composition the excess enthalpy decreases with
increasing temperature and at the higher composition the excess
enthalpy increases with increasing temperature. The experi-
mental uncertainty of the measurements is 2 %; therefore, the
change in the temperature trend is real and is not within
experimental uncertainty.
The consistency of the excess enthalpies with the excess heat
capacities cannot be checked because, as mentioned previously,
the experimental uncertainty for the heat capacity measurements
is of the same order of magnitude as the excess heat capacities.
The excess enthalpy is endothermic for the water + [emim]-
[OTf] mixture (Figure 3b), and the excess enthalpy increases
with increasing temperature. The water + [emim][TFA] mixture
(Figure 3c) is exothermic, and again, the temperature trend
shows that the excess enthalpy increases with increasing
temperature.
good estimate. In Table 4, the parameters corresponding to the
linear fit of the isotherms of Cp as a function of mole fraction
of water are shown for the three binary systems studied. The
linear correlations for the highest and the lowest experimental
temperature are plotted as the solid lines in Figure 2. The
experimental points do not lay far from the straight lines,
consistent with the low values of standard deviation listed in
Table 4.
Excess Enthalpy. The excess molar enthalpies, HE, of the
three water + IL mixtures, [emim][EtSO4], [emim][OTf], and
[emim][TFA], are shown in Figure 3a, b, and c, respectively,
at temperatures (313.15, 323.15, 333.15, and 348.15) K. The
experimental data were fit to a Redlich-Kister-type equation
as represented by the solid lines
HE ) x1 · x2(A + B(x1 - x2) + C(x1 - x2)2)
(3)
where xi is the mole fraction of component i and A, B, and C
are the fitting parameters. The parameters A, B, and C, along
with the standard deviations, σ, are shown in Table 5. The
experimental data can be found in Tables S4 to S6 in the
Supporting Information.
For the system of water + [emim][TFA], Domanska et al.
calculated an infinite dilution partial molar enthalpy of water

2116 Journal of Chemical & Engineering Data, Vol. 53, No. 9, 2008
Table 4. Slope, a, and Intercept, b, Along with the Standard
Deviation, σ, for the Correlation of Molar Heat Capacity in Binary
Systems Water + IL as a Function of Mole Fraction of Water
T/K
a/102 J·mol-1 ·K-1 b/102 J·mol-1 ·K-1 σ/J·mol-1 ·K-1
Water + [emim][EtSO4]
283.15
288.15
293.15
298.15
303.15
308.15
313.15
318.15
323.15
328.15
333.15
338.15
343.15
-2.836
-2.862
-2.883
-2.927
-2.960
-2.987
-3.015
-3.041
-3.068
-3.094
-3.121
-3.149
-3.180
3.566
3.599
3.625
3.680
3.717
3.747
3.778
3.808
3.838
3.868
3.899
3.929
3.961
5.1
5.0
5.1
5.1
5.0
5.0
5.0
5.0
5.1
5.0
5.0
4.9
4.8
Water + [emim][OTf]
3.489
283.15
288.15
293.15
298.15
303.15
308.15
313.15
318.15
323.15
328.15
333.15
338.15
343.15
-2.745
-2.763
-2.789
-2.816
-2.862
-2.885
-2.908
-2.932
-2.953
-2.975
-2.994
-3.021
-3.038
2.7
2.7
2.5
2.3
2.7
2.6
2.7
2.8
2.9
2.9
2.9
2.8
3.0
3.512
3.541
3.579
3.628
3.653
3.679
3.705
3.729
3.754
3.776
3.805
3.826
Water + [emim][TFA]
283.15
288.15
293.15
298.15
303.15
308.15
313.15
318.15
323.15
328.15
333.15
338.15
343.15
-2.336
-2.357
-2.372
-2.396
-2.429
-2.448
-2.468
-2.487
-2.509
-2.541
-2.561
-2.592
-2.618
3.133
3.165
3.186
3.223
3.259
3.282
3.305
3.327
3.352
3.386
3.408
3.440
3.468
4.1
4.3
4.8
5.3
5.7
6.1
6.4
6.8
7.0
7.7
7.9
8.0
8.1
Figure 3. Excess enthalpy, HE, for the binary systems water (1) + IL (2)
as a function of mole fraction of water, at different temperatures: b, 313.15
K; 9, 323.15 K; 2, 333.15 K; [, 348.15 K. Empirical correlations are
plotted as solid lines. ILs: (a) [emim][EtSO4]; (b) [emim][OTf]; (c)
[emim][TFA].
Table 5. Constants A, B, and C, along with the Standard Deviation,
σ, for the Correlation of Excess Enthalpy in Binary Systems Water
+ IL as a Function of Mole Fraction of Water, Equation 3
to be -5607 J ·mol-1 which was obtained from infinite dilution
activity coefficients at temperatures of (348.15, 358.15, and
368.15) K.41 Using the experimental data from this study at
348.15 K (the lower end of temperature for Domanska et al.),
the calculated partial molar enthalpy of water at a mole fraction
of 0.001 is -6020 J ·mol-1. This is in very good agreement as
the partial molar enthalpy for this system increases with
increasing temperature.
T/K
A/103 J·mol-1 B/103 J·mol-1 C/102 J·mol-1 σ/J·mol-1
Water + [emim][EtSO4]
313.14
323.12
333.11
348.12
-3.104
-3.169
-3.193
-3.203
-2.732
-2.407
-2.175
-1.635
-28.28
-23.99
-22.87
-16.41
3.29
9.69
5.48
1.44
The excess enthalpies for all three of these systems are very
small, with the maximums at ( (1 to 2) kJ·mol-1. Therefore,
as a first approximation, one may be able to neglect HE in
calculating enthalpies for these mixtures. For example, to
calculate the enthalpy of the water + IL liquid mixture at a
temperature T relative to a reference temperature by eq 4, one
finds that the first two terms are large in comparison to the
excess enthalpy.
Water + [emim][OTf]
313.13
323.12
333.13
348.12
4.512
4.625
4.764
5.000
1.609
2.003
2.400
3.095
2.030
4.508
7.270
11.79
0.80
0.76
0.55
0.71
Water + [emim][TFA]
-3.873
313.04
323.04
333.06
348.06
-8.561
-8.411
-8.284
-8.011
-16.45
-12.29
-8.225
-1.854
0.52
0.61
0.91
0.80
-3.387
-2.930
-2.172
n
n
nIL
H2O
total
T
T
H )
C dT
+
C dT
+ HE (4)
∫
∫
(
)
(
)
f
p
p
273
H O
273
IL
n
2
total
the acid dissociation constant, a measure of unlike species
interactions, and the glass transition temperature, a measure of
the strength of IL/IL interactions, to discuss the results.
The pKa, the acid dissociation constant, is an indicator of the
strength of the conjugate base of an acid. In our case, it is a
Nonetheless, the excess enthalpies can provide a starting point
for understanding species interactions. The excess enthalpy can
be best explained as the difference in the strength of interactions
between unlike species compared to like species. Here we use

Journal of Chemical & Engineering Data, Vol. 53, No. 9, 2008 2117
Table 6. Acid Dissociation Constants for the Conjugate Acids of the
IL Anions Used in This Work
Table 7. Glass Transition Temperatures of the Pure ILs
IL
Tg/K
[emim][OTf]
[emim][EtSO4]
[emim][TFA]
259a
192.8518
182
a Melting point.
the strength of IL/IL interactions, are also consistent with the
observed trends in excess enthalpies.
Decomposition of [emim][EtSO4]. It is of great importance
to note that the IL [emim][EtSO4] was found to decompose
over time by reacting with water to form 1-ethyl-3-meth-
ylimidazolium hydrogen sulfate, [emim][HSO4], and ethanol.
This was discovered when checking the batch to batch
consistency of the ILs. The batch used in this study was
compared to two other batches: a commercial sample from
2004 and a batch made in-house in 2006. A triple quadrupole
liquid chromatography-mass spectrometry (LCMS) ap-
paratus confirmed the decomposition products. Quantitatively,
NMR spectroscopy revealed that the 2004 and 2006 batches
were approximately 23 % and 28 % decomposed, respec-
tively, even though these samples had been dried before
storage. The decomposition of the batch of [emim][EtSO4]
used in this study was below the detection limit of the
instrument (< 1 % decomposed). The NMR results can be
found in Figure S1 of the Supporting Information. Clearly,
decomposition must be taken into account when trying to
reproduce any physical property data with this IL. Due to
the composition and property changes with time, [emim]-
[EtSO4] may not be a good IL choice for practical applica-
tions.
measure of the strength of the basicity of the anion. A stronger
base should have greater interactions with water; therefore, a
large pKa corresponds to a strong base and strong anion-water
interactions. The pKa’s of the acids of interest are shown in
Table 6. The pKa of trifluoromethanesulfonic acid is -11,42
indicating that it should have weak interactions with water,
which is consistent with the experimental endothermic excess
enthalpy (HE < 1373 J ·mol-1). The ethylsulfuric acid has a
larger pKa of -3.14,43 indicating stronger interactions with water
than [emim][OTf], and this is consistent with the slightly
exothermic excess enthalpy for water + [emim][EtSO4] (HE >
-979 J ·mol-1). Lastly, the trifluoroacetic acid has the largest
pKa of 0.52,44 making the [TFA]- anion the strongest base of
the three. This suggests that [emim][TFA] should have the
strongest interactions with water. Indeed, water + [emim][TFA]
has the largest exothermic excess enthalpy of the systems studied
(HE > -2258 J ·mol-1).
The pKa’s indicate the strength of the anion-water interac-
tions. The [emim][OTf] has the weakest anion-water interac-
tions, and [emim][TFA] has the strongest anion-water inter-
actions. Therefore, the pKa’s aptly explain the excess enthalpies
for these three systems. However, there are other examples in
the literature which do not follow this trend.
Conclusions
Heat capacities for the binary systems water + [emim]-
[EtSO4], water + [emim][OTf], and water + [emim][TFA] were
experimentally determined in the temperature range from
(283.15 to 343.15) K and over the whole range of composition.
The molar heat capacities (Cp) decreased as the water concen-
tration increased, with the Cp of the pure ILs being 4 to 5 times
greater than that of water. A slight increase of Cp with increasing
temperature was observed. Since the molecular weights of the
ILs are high, the specific heat capacity (cp) of water is much
higher than that of the pure ILs. The variation of Cp with
temperature was correlated with a two-parameter empirical
equation; for the correlation of Cp as a function of mole fraction,
a straight line was adequate.
The excess enthalpies were positive values (endothermic) for
water + [emim][OTf] and negative values (exothermic) for
water + [emim][EtSO4] and water + [emim][TFA]. The HE
values were fit with a Redlich-Kister-type equation using three
fitting parameters. The excess enthalpy measurements provide
insight into the interactions of the species in solution, which is
important for gaining a fundamental thermodynamic under-
standing, selecting ILs for processes, and predicting the effect
of water as an impurity.
Constantinescu et al. reported the excess enthalpies of both
[choline][lactate] and [choline][glycolate] with water.31 The struc-
tures of the respective acids are very similar, and the pKa’s are
virtually the same;44,45 however, the excess enthalpies are very
different. For a given isotherm, there is an order of magnitude
difference between the excess enthalpies despite the similar pKa’s.
Therefore, the interactions in water + IL systems cannot be
explained solely by the pKa’s of the conjugate acids of the anions.
The glass transition temperature is the temperature below which
an amorphous material acts as a solid and above which it acts as
a liquid. This temperature can indicate the strength of the
cation-anion interactions in a pure IL and, thus, is an indicator of
the strength of IL/IL interactions. An IL with strong cation-anion
interactions should have a high glass transition temperature. The
glass transition temperature can be measured on a DSC.
For similar unlike pair interactions (IL/water), the IL with
stronger IL/IL interactions should have a larger (larger positive
value or smaller negative value) excess enthalpy. The glass
transition temperature of [emim][EtSO4] has been reported as
192.85 K.18 On the basis of the excess enthalpies, we would then
expect the glass transition temperature of [emim][OTf] to be higher
and the [emim][TFA] to be lower than that of [emim][EtSO4]. DSC
measurements, using a ramp rate of 10 °C·min-1, of [emim][OTf]
and [emim][TFA] confirm this trend (Table 7). The glass transition
temperature of [emim][TFA] is 182 K, which indicates weaker
cation-anion interactions than in [emim][EtSO4]. [emim][OTf]
does not have a glass transition temperature but melts at 259 K.
This suggests strong interactions between the anion and cation for
[emim][OTf]. Thus, the glass transition temperatures, indicating
In addition, these experimental measurements showed that
[emim][EtSO4] decomposes in the presence of water under
ambient conditions.
Acknowledgment
We thank Dennis Birdsell and Michelle Joyce of the Center for
Environmental Science & Technology at the University of Notre
Dame for their expertise in optical emission and mass spectrometry,

2118 Journal of Chemical & Engineering Data, Vol. 53, No. 9, 2008
The heat capacity and standard enthalpy of formation of EMIES.
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his help with the NMR spectroscopy.
Supporting Information Available:
Experimental molar heat capacities of the systems water +
IL (Table S1 to S3); experimental excess enthalpies of the
systems water + IL (Tables S4 to S6); and NMR results of the
[emim][EtSO4] decomposition (Figure S1). This material is
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