Does alkyl chain length really matter? Structure-property relationships in thermochemistry of ionic liquids

Article abstract of DOI:10.1016/j.tca.2013.04.003

DSC was used for determination of reaction enthalpies of synthesis of ionic liquids [C_nmim][Cl]. A combination of DSC with quantum chemical calculations presents an indirect way to study thermodynamics of ionic liquids. The indirect procedure for vaporization enthalpy was validated with the direct experimental measurements by using thermogravimetry. First-principles calculations of the enthalpy of formation in the gaseous phase have been performed for the ionic species using the CBS-QB3 and G3 (MP2) theory. Experimental DSC data for homologous series of alkyl substituted imidazolium, pyridinium, and pyrrolidinium based ionic liquids with anions [Cl] and [Br] were collected from the literature. We have shown that enthalpies of formation, enthalpies of vaporization, and lattice potential energies are linearly dependant on the alkyl chain length. The thermochemical properties of ILs generally obey the group additivity rules and the values of the additivity parameters for enthalpies of formation and vaporization seem to be very close to those for molecular compounds.

Full text of DOI:10.1016/j.tca.2013.04.003

Accepted Manuscript
Title: Does Alkyl Chain Length Really Matter?
Structure-Property Relationships in Thermochemistry of Ionic
Liquids
Author: Sergey P. Verevkin Dzmitry H. Zaitsau Vladimir N.
Emel’yanenko Ricardas V. Ralys Andrei V. Yermalayeu
Christoph Schick
PII:
S0040-6031(13)00186-X
DOI:
Reference:
http://dx.doi.org/doi:10.1016/j.tca.2013.04.003
TCA 76426
To appear in:
Thermochimica Acta
Received date:
Revised date:
Accepted date:
28-11-2012
16-3-2013
2-4-2013
Please cite this article as: S.P. Verevkin, D.H. Zaitsau, V.N. Emel’yanenko, R.V. Ralys,
A.V. Yermalayeu, C. Schick, Does Alkyl Chain Length Really Matter? Structure-
Property Relationships in Thermochemistry of Ionic Liquids, Thermochimica Acta
(2013), http://dx.doi.org/10.1016/j.tca.2013.04.003
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Thermochimica Acta, 2013
Does Alkyl Chain Length Really Matter?
Structure-Property Relationships in Thermochemistry of Ionic Liquids.
Sergey P. Verevkin,a,b Dzmitry H. Zaitsau, Vladimir N. Emel’yanenko, Ricardas V. Ralys,
Andrei V. Yermalayeu
Department of Physical Chemistry, University of Rostock, Dr-Lorenz-Weg. 1, 18059 Rostock,
Germany
Christoph Schickb
Department of Physics, University of Rostock, Rostock,Wismarsche Str. 43-45, 18057 Rostock
Germany
Abstract
DSC was used for determination of reaction enthalpies of synthesis of ionic liquids [Cnmim][Cl].
A combination of DSC with quantum chemical calculations presents an indirect way to study
thermodynamics of ionic liquids. The indirect procedure for vaporization enthalpy was validated
with the direct experimental measurements by using thermogravimetry. First-principles
calculations of the enthalpy of formation in the gaseous phase have been performed for the ionic
species using the CBS-QB3 and G3(MP2) theory. Experimental DSC data for homologous series
of alkyl substituted imidazolium, pyridinium, and pyrrolidinium based ionic liquids with anions
[Cl] and [Br] were collected from the literature. We have shown that enthalpies of formation,
enthalpies of vaporization, and lattice potential energies are linearly dependant on the alkyl chain
length. The thermochemical properties of ILs generally obey the group additivity rules and the
values of the additivity parameters for enthalpies of formation and vaporization seem to be very
close to those for molecular compounds.
Key Words: Ionic liquids, enthalpy of formation, enthalpy of vaporization, enthalpy of fusion,
lattice potential energy, differential scanning calorimetry, thermogravimetry, quantum chemical
calculations.
a To whom correspondence should be addressed, E-mail: sergey.verevkin@uni-rostock.de
Fax: +49 381 498 6524; Tel: +49 381 498 6508.
b
Faculty of Interdisciplinary Research, Department “Life, Light and Matter”, University of
Rostock
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1. INTRODUCTION
This century the intensive search for new environmentally benign nonaqueous solvents, which
can easily be recovered, is a long-standing priority for the development of sustainable chemical
processes. Four promising neoteric solvents are the main targets of current academic and
industrial research for chemical engineering: ionic liquids (ILs), supercritical fluids, fluorous
solvents, and liquid polymers. However, the most attention during the last decade have attracted
ILs due to their remarkable physical and chemical properties (thermal stability, conductivity,
non-flammability) and that offer a potential for industrial application.
ILs can be tailored in at least three-dimensional variations by combination of different cations,
different anions, as well as different alkyl chains connected to the cations or anions.
ILs can be additionally tailored to specific applications by fine-tuning the functional groups
connected to organic cation and the inorganic/organic anion. The use of new ILs based solvent
systems represents new avenues in the search of environmentally benign technologies. A good
overview of the advanced research on ILs is given in ref.[1] However, the physical-chemical
behaviour of ILs is still not clearly understood and needs further studies [2]. In order to
rationalise in a simple way the effect of changes in the IL property due to the structure variation,
studying homologous series of ILs seems to be a valuable methodology. It is already known that
the alkyl chain length of the cation influences both the physical and chemical properties of ILs.
The thermophysical properties such as density, viscosity, surface tension, melting point, and
thermal stability are significantly affected by changes in alkyl chain length of the cation or the
anion [3]. For example, when the alkyl chain length in the [Cnmim][NTf2] series (n = 1, 2, 4, 6,
8, 10, 12) is increased, a decrease in density is observed [4,5]. The surface tension initially also
decreases [4] but reaches a plateau for alkyl chain lengths greater than n = 8. For the same
homologues series an increase in the alkyl side chain length of the cation produces an increase in
viscosity of the IL [5]. An increase of heat capacities of [Cnmim][NTf2] with n = 2 and 18 was
observed by adiabatic calorimetry [6]. The X-ray scattering experiments and molecular dynamics
simulations
of
1-alkyl-1-methylpyrrolidinium
bis-(trifluoromethanesulfonyl)imides
([Cn1Py][NTf2], n = 3, 4, 6, 8, 10) revealed the alkyl chain length effects on the nanoscale
organization of pyrrolidinium-based ionic liquids. The latter fact is important because ILs have
been proven to be efficient and useful reaction media for the synthesis of inorganic nanomaterials
[7].
The use of mixtures of ILs to generate new composite materials provides another facile
means of readily tuning or tailoring the desired properties of ionic media. The variation of the
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alkyl chains lengths on the cations plays a significant role in determining both the phase
behaviour and the physicochemical properties of the mixtures [8].
Thus the very general chain-length tendencies and trends in thermophysical properties seems
to be quite apparent: a longer cation side chain is commonly accompanied by lower density,
lower solubility, slower diffusion, and higher viscosity [5], however tireless extended
experiments on this field are still of paramount importance for practical applications. We have
intensively contributed to the thermodynamics of ILs [9-17], but a recent focus of our ionic
liquids activities has been shifted more to the thermochemical than to thermophysical properties.
It has turned out that the thermochemistry of ILs is suffering of lack of experimental data and
more – the available data are in disarray [15]. The thermochemical properties of our interest are
generally combined in the common textbook equation:
gl Hm =  f Hm (g) -  f Hm (liq)
(1)
where gl Hm is the standard molar enthalpy of vaporization,
(g) is the standard molar
 f Hm
 f Hm
gaseous enthalpy of formation of an IL, and
(liq) is the standard molar enthalpy of
formation in the liquid state. We use a high-precise combustion calorimetry to measure energy of
combustion of an IL and to derive from this energy  f Hm (liq) [9]. We also use a commercially
available differential scanning calorimeter (DSC) to measure enthalpy of reaction, r Hm (liq), of
 f Hm
an ionic liquid synthesis from precursors [12]. The latter value provide
(liq) by using the
classic Hess’s Law. Enthalpy of vaporization of an IL is usually measured by any suitable
method: Knudsen effusion [16], transpiration [9], quartz-crystal microbalance (QCM) [11], or
thermogravimetry (TGA) [13] available in our lab. Finally, the value of  f Hm (g) is usually
calculated as the sum of
experimental result for
(liq) and gl Hm . We used to perform a consistency test of this
 f Hm
 f Hm
(g) with a theoretical value calculated using one of the suitable
high-level first-principles methods (e.g. G3(MP2), CBS-QB3, etc.) [9].
From our experiences each experimental method is not free from certain restrictions and
disadvantages. For example, in the direct experimental measurement of vaporization enthalpies
by using TGA, QCM, or transpiration we are confronted with two main problems: due to the
extremely low vapor pressures of ILs, they are practically not measurable at ambient and even at
elevated temperatures, whereas at high temperatures over 500 K some of them may decompose.
Vaporization enthalpies, gl Hm of ILs obtained indirectly by using the calorimetry (combustion
or DSC) coupled with the modern high-level quantum chemical computations also always need
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to be checked for internal consistency. In our opinion, in order to obtain reliable thermodynamic
data sets, several experimental techniques have to be used to study the same IL sample
simultaneously. On the one hand, a close agreement of the results could give a good prove of the
data quality. On the other hand, a possible disagreement could serve for further development of
the experimental methods used. Following this guideline, in this work we have performed DSC
studies of the [Cnmim][Cl] (with n = 5, 6, 7, and 8) synthesis reactions according to the following
equation:
1-Me-imidazole (liq) + CnCl (liq) = [Cnmim][Cl] (liq)
(2)
Experimental enthalpies of reactions have been used to derive enthalpies of formation, enthalpies
of vaporization, and lattice energies of ionic liquids. Validity of vaporization enthalpies for some
[Cnmim][Cl] obtained indirectly over DSC measurements has been proven by direct TGA
measurements. Having established the reliable data set for [Cnmim][Cl] family, we collected
available literature data on thermochemical properties of imidazolium, pyridinium, and
pyrrolidinium based ILs and analysed their dependence on the chain length
2. EXPERIMENTAL SECTION
2.1. Materials and Chemicals. 1-Methyl-imidazole (MeIm) and alkyl chlorides (CnCl) were
commercially available with purity ≥ 99%. No impurities (greater than 0.1 mass per cent) could
be detected in the samples using a gas chromatograph equipped with a flame ionization detector.
All precursors were fresh distilled before use. Samples of 1-ethyl-3-methyl-imidazolium chloride
[C2mim][Cl], 1-hexyl-3-methyl-imidazolium chloride [C6mim][Cl], and 1-butyl-3-methyl-
imidazolium bis-(trifluoromethanesulfonyl)-imides [C4mim][NTf2] were commercially available
from Io-Li-Tec. Purities of IL samples were of 99% according to specification by supplier. 1H
and 13C NMR spectra were recorded but no impurities in the samples were detected. IL samples
were kept at 333 K in vacuum for more than 24 hours in order to remove possible traces of
solvents and moisture prior to TGA experiments. The water concentration (<100 ppm) was
determined by Karl Fischer titration. Handling with the samples occurred in a glove box.
2.2. DSC – measurement of reaction enthalpies. Experimental enthalpies of ILs synthesis
reactions (according to Eq. 1) measured by DSC open the direct way to derive enthalpies of
formation of ILs as well as the indirect way to estimate vaporization enthalpies. Enthalpies of
reaction were measured using the calibrated computer controlled Mettler-Toledo 822 DSC.
About 10-15 mg of starting materials were placed in a 40 µL aluminum pan. As a rule, the DSC-
pan was filled with the defined quantity of the first precursor 1-alkyl-chloride, then with the layer
of the solvent [C4mim][NTf2], and finally with the second precursor 1-methyl-imidazole. We
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5
used some excess of 1-methyl-imidazole (see Table 1) in order to reach complete conversion of
the alkyl chloride at the end of reaction. The layer of the solvent (about 75-80 wt. %;) was
necessary in order to separate precursors and to avoid beginning of the chemical reaction outside
of the DSC. The DSC pan was weighed with the resolution of ±0.000005 g and air buoyancy
corrections were taken into account. The pan was hermetically sealed with a cover. The pan was
weighed before and after each experiment to detect any mass loss during the measurements.
Measurements were performed in the range 273 - 500 K with scanning rates between 10 and 50
K·min-1. It is well known that ionic liquid synthesis reactions from an amine and an alkylhalide
occurs very rapidly already at 353 K. Reactions are also highly exotermic but very selective with
conversion over 99%. In the DSC pan reaction conditions the relatively volatile starting materials
(1-methyl-imidazole and alkylhalides) rapidly produced the very low volatile ionic liquids
already in the temperature range 343-373 K (see Fig. S7 in SI). Further heating of the pan only
facilitated the complete conversion of starting materials. Completeness of reactions under study
in the excess of 1-methyl-imidazole was over 99% and therefore no volatile materials were
present in the pan at elevated temperatures. The mass of the pan was controlled before and after
experiment. Any traces of unconverted starting materials or decomposition products would cause
an explosion of the pan, but it was not observed. During the development of the DSC method we
determined the degree of conversion in the following way: after completing of the DSC
experiment the pan was cooled and a pin hole was made in the cover. The pan was heated in
vacuum at 353 K within 12 h and weighed before and after the heating. The mass loss could be
ascribed to the volatility of precursors, because the ionic liquid in these conditions has no
measurable vapour pressure. In addition we checked the absence of unreacted alkylhalides or
decomposition products in the DSC pan by GLC. The degree of conversion determined on both
ways was always 99.8± 0.3%. The area of the DSC peak occurring during the chemical reaction
was a measure of the reaction enthalpy. In order to obtain a reaction heat effect in a dynamic
temperature scan mode the temperature difference between a sample pan and an empty reference
pan was measured and converted into a heat flux value. The peak corresponding to this
exothermic heat effect was then integrated over the time range of the effect and a total enthalpy
of reaction was determined from the area of the DSC peak. The heat released during the IL
synthesis reaction was related to the molar amount of a stoichiometrically deficient reactant in
the starting formulation. The experimental procedure has been elaborated in our lab just recently
and it was reported elsewhere [12]. Experimental results of the DSC study are collected in Table
1. This data was used to derive  f Hm (liq), gl Hm , and  f Hm (g).
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2.3. TGA measurements - vaporization enthalpy. We used the TGA procedure for validation
of the vaporization enthalpies, gl Hm , obtained from the indirect DSC results. We used a Perkin
Elmer Pyris 6 TGA to measure gl Hm from the temperature dependence of the mass loss rates
measurements. About 50-70 mg of the IL sample was placed in a plain platinum crucible inside
of the measuring head of the TGA. The sample was step-wise heated and a mass loss of 0.1-0.8
mg from the crucible was recorded at each isothermal step. Isothermal mass loss rate dm/dt was
monitored in the temperature range 403-462 K at a nitrogen flow rate of 140 ml·min-1. An
absence of decomposition of the IL in the experimental conditions was confirmed by ATR-IR
spectroscopy. No changes in the spectra taken from the initial and the residual IL in the crucible
were detected for the ILs under study. The detailed experimental procedure has been elaborated
in our lab and it was reported elsewhere [13]. Experimental results of the TGA studies are
collected in Tables S1 and S2 (ESI).
2.4. Quantum-chemical calculations. Standard first-principles calculations were performed
using the Gaussian 09 program package [18]. From 20 to 30 initial conformations of ionic liquids
were generated which were optimized on B3LYP/6-31e+(d,p) level of theory. After that the
lowest by energy conformation was calculated on CBS-QB3 as well as G3(MP2) level. CBS-
QB3 theory uses geometries from B3LYP/6-311G(2d,d,p) calculation, scaled zero-point energies
from B3LYP/6-311G(2d,d,p) calculation followed by a series of single-point energy calculations
at the MP2/6-311G(3df,2df,2p), MP4(SDQ)/6-31G(d(f),p) and CCSD(T)/6-31G† levels of theory
[19]. G3(MP2) theory uses geometries from second-order perturbation theory and scaled zero-
point energies from Hartree-Fock theory followed by a series of single-point energy calculations
at the MP2(Full)/6-31G(d), QCISD(T)/6-31G(d) and MP2/GTMP2Large levels of theory (for
details see reference [20]). However the G2MP2 method is not able to calculate the bromine
containing compounds. Calculated values of the enthalpies of reaction and enthalpies of
formation are based on the electronic energy calculations obtained using standard procedures of
statistical thermodynamics [21]. Details are given in ESI Tables S3-S8.
3. RESULTS AND DISCUSSION
The quaternization reactions broadly used for synthesis of ILs are highly exothermic [22].
Already first kinetic studies of the [C4mim][Cl] synthesis from 1-methylimidazole and 1-
chlorobutane in a batch reactor [22-23] according to Eq. (2) have shown very complex
phase behaviour of the reaction participants in the course of reaction [22]. With the
MeIm-conversion >8%, neat synthesis lead to two phases either rich in IL or 1-
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7
chlorobutane, whereby MeIm is soluble in both phases. Addition of ethanol (>20 vol%)
lead to a single-phase synthesis and to an increase of the effective rate constant by IL-
formation. The reaction was very slow and at 343 K about 50 h were required to reach
90% MeIm conversion. However, an acceleration of the rate in the neat conditions by
increasing the temperature was limited, because of decomposition of MeIm and
contamination the IL. Only dilution of the reacting system with ethanol as the solvent
ensured the single-phase synthesis and improved the quality of the IL [22]. Such a
sophisticated behaviour of the reacting system requires the reliable thermodynamic data for the
reaction engineering and process intensification. A careful heat management is also
indispensable for reactor design, preventing possible thermal runaway, and finally for the
production of high quality ILs. Thus the DSC studies performed in this work are important from
both theoretical and practical points of view.
3.1 Enthalpies for ILs synthesis reactions from DSC. Enthalpies of the ILs synthesis
reactions, rHm (liq.), were determined according to Eq. (1) in a series of three consequent DSC
runs under optimized conditions elaborated in [12]: in excess of MeIm, dilution with the solvent
[C4mim][NTf2], and temperature scanning with 50 K/min in the range 298-523 K. A sharp
reaction peak with well defined baseline was achieved in these conditions providing reproducible
results. Experimental results are given in Table 1. As can be seen from this table all reactions
were strong exothermic.
Strictly speaking, enthalpies of reaction, r Hm , derived by integration of the DSC-peak are
referred to the temperature of the peak maximum, Tmax (see Table 1). However, from our
experiences [12], a correction estimated for adjustment of the measured enthalpies to the
reference temperature 298 K (using heat capacities of the reaction participants) does not exceed
0.3-0.4 kJ·mol-1 and this correction is well negligible within the boundaries of the DSC
experimental uncertainties of 1-3 kJ·mol-1. We set in this work the experimental DSC-values of
r Hm (see Table 1) to be equal those at the reference temperature 298 K.
Recently [24] we reported the reaction enthalpy of the synthesis of the ionic liquid [C4mim][Cl]
to be (-78.6±0.8) kJ mol-1 from DSC. In this work we extended our DSC experiments with
reaction 2 using a series of longer chained alkyl chlorides CnCl (with n = 5, 6, 7, and 8) in order
to study the chain length dependence. It makes oneself conspicuous, that enthalpies of reaction
r Hm (liq) (see Table 1) seems to be independent on the chain length for n = 4, 5, 6, 7, and 8.
The similar trends were already observed for homologous series of [Cnmim][Br] [12], as well as
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for pyridinium based [CnPy][Cl], [CnPy][Br] and pyrrolidinium based [Cn,1Py][Cl], [Cn,1Py][Br]
ionic liquid families [25].
3.2. Molar enthalpies of formation of ILs from DSC-measurements. The combustion bomb
calorimetry with oxygen is the most acknowledged method to derive enthalpies of formation of
ILs [9]. But this method as a rule requires extra pure (99.9%) samples in the minimal quantity of
5-10 g. Moreover, it is not always easy to establish experimental conditions to ensure
completeness of the IL combustion process. Combustion calorimetry of the common ILs with
[Cl], [Br], and [I] anions is also not trivial and requires using of the rotation-bomb technique
[17]. Additionally, the experimental procedure for analyzing of the bomb solutions requires
careful elaboration of the ion-chromatography method. While working with ILs over the last
decade, we have recognised an additional advantage of using DSC to derive thermochemical
properties of ILs, namely a possibility to measure a broader range of thermochemical properties
indirectly. It has turned out [12,24] that measurements of r Hm (liq) using the DSC are less
demanding as compared to the combustion calorimetry. Moreover, purity of the fresh distilled
precursors at the level of 99% is already quite acceptable for DSC experiments. These
advantages are of crucial importance for analyzing ILs where research samples are expensive and
often available only in small quantities. DSC offers a feasible alternative to obtain reliable data
for enthalpies of formation of ILs. For example, instead of tedious combustion experiments with
[C5mim][Cl], enthalpy of e.g. reaction 2 was measured by DSC. According to the Hess’s Law,
reaction enthalpy is defined as the difference between enthalpies of formation,  f Hm of the
products and reactants (e. g. for Eq. 2):
r Hm (liq, 298 K) =
([C5mim][Cl], liq) -
(3)
 f Hm
 f Hm
 f Hm
(MeIm, liq) - (C5Cl, liq)
Using this equation and enthalpies of reaction listed in Table 1, molar enthalpies of formation for
the IL family [Cnmim][Cl] were calculated (see column 6 in Tables 2 and 3). Experimental molar
enthalpies of formation of the precursors of reactions (2) MeIm and CnCl given in Table S3 and
S4 in ESI. These enthalpies of formation,  f Hm , can open an excess to a lattice potential energy
and lattice enthalpy of an ionic liquids which are usually rationalized in different types of the
Born-Fajans-Haber cycles. [26-30] The enthalpies of formation of the ILs  f Hm ([Cnmim][Cl],
liq.) also can be used for estimation of vaporization enthalpies, gl Hm , of ILs in indirect way.
3.3. Molar vaporization enthalpies from DSC study. Assessment of solubility parameters,
safety calculations of thermal processes using ILs as thermofluid as well as validation of the
molecular dynamic simulations are based on vaporization enthalpy data. Recently we have
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developed an indirect method to derive gl Hm (298 K) using a combination of DSC with
quantum-chemical calculations [12,24,25]. In this work we applied this procedure for estimation
of vaporization enthalpies of the [Cnmim][Cl] family according to Eq. 2. There are two
possibilities [12] to derive enthalpies of vaporization by using Eq. 1. The first one is based on
Eq. (1) using the difference between theoretical gas phase enthalpies of formation,  f Hm (g, 298
K) calculated with the composite CBS-QB3 method (see Table 2, column 4) and the
experimental molar enthalpies of formation in the liquid state,  f Hm (liq, 298 K) measured by
DSC (see Table 2, column 6). Results are given in Table 2 column 7. The first way is not
optimal, because it suffers from ambiguity of value of the theoretical enthalpy of formation,
 f Hm (g) used in Eq. (1) [12,31]. Enthalpies of formation  f Hm (g, 298 K) for Br containing
compounds calculated by CBS-QB3 method by atomization procedure have to be corrected
-1
f Hm
.
according to the linear equation developed in ref [12]:
(g) (exp)/ kJ mol = 1.002 ×
f Hm
(g)(AT) + 28.2. Using this correlation the “corrected” enthalpies of formation of the
[Cnmim][Br] have been calculated (see Table 2, column 4).
The second way is to calculate with the first-principles methods directly the enthalpy of the
reaction,  H (g), shown in Eq. (4) ([Cnmim][Cl] family is taken only as example):
r
m
1-Me-imidazole (g) + CnCl (g) = [Cnmim][Cl] (g)
(4)
As a rule, the value of  H (g) is calculated from the enthalpies H298 at T = 298 K for all
r
m
reaction participants of reaction (4) by the CBS-QB3 (or any suitable ab initio method) approach
and using the Hess’ Law, e.g. for [C5mim][Cl]) the enthalpy of reaction was calculated as
follows:
 H (g) = H ([C mim][Cl], g) - H (MeIm, g) - H (C Cl, g) = -(15.2±4.0) kJ mol-1
(5)

298
5
298
298
5
r
m
It is important to underline that the value rHm (g) was obtained using quantum chemical
methods directly from the calculated H298 bypassing the conventional calculation of the
enthalpies of formation with help of atomization or isodesmic procedures [31]. This rHm (g)
value is not affected by the choice of atomization or bond separation procedure required for the
first way. Values of rHm (g) calculated for all reactions under study according to Eq. (5) are
listed in Table 2 column 3.
Having established  H (liq)exp from the DSC experiment and rHm (g) from quantum
r
m
chemistry, we have derived the vaporization enthalpy of e.g. for [C5mim][Cl] as follows:
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g H (IL, 298 K) = - H (liq)exp+ H (g) + g H (MeIm) + g H (C5Cl)] =
l
m
r
m
r
m
l
m
l
m
= - (-76.4) + (-15.2) + (55.0) + (38.2) = 154.5±4.1 kJ mol-1
(6)

Using the theoretical value rHm (g) calculated with CBS-QB3 in the gas phase and enthalpies of
vaporization at 298 K of precursors gl Hm (MeIm) = 55.0 ± 0.3 kJmol-1 (measured in our lab
recently [32]), g H (C5H11Cl)] = 38.2 ± 0.1 kJmol-1 (available from the literature [33]) and the
l
m
enthalpy of reaction  H (liq)exp measured using the DSC in the liquid phase, the enthalpy of
r
m
vaporization of [C5mim][Cl] has been calculated. In the same way, enthalpies of vaporization of
other ILs were calculated with help of Eq. (5) and listed in the last column of Table 2.
Comparison of vaporization enthalpies derived using two different procedures (i.e. Eq. 1 and Eq.
3) shows that both methods provide values in close agreement. However we should recommend
Eq. (3), because the  H (g)-values for this procedure were calculated directly from H298
r
m
avoiding any procedure for converting quantum-chemical results into enthalpy of formation.
We additionally tested another composite method - G3(MP2), which is significantly less
time consuming in comparison with the CBS-QB3. The G3(MP2) method seems to be also
suitable for calculations of f Hm (g, 298 K) of the chlorine containing ILs Results obtained from
-1
.
the CBS-QB3 and the G3(MP2) methods are in good agreement within a few kJ mol (see Table
3, columns 7 and 8). Thus the G3(MP2) method could be successfully used for calculations of
the chlorine containing ILs, but unfortunately this method method has not parameterization for
Br containing molecules.
3.4. Molar vaporization enthalpies from TGA. TGA procedure for determination of
vaporization enthalpies of extremely low volatile liquid compounds have been developed just
recently [13,15]. It has been shown that the vaporization enthalpies for the ILs could be obtained
with an accuracy better than ±3 kJ·mol-1. The TGA approach has few important advantages in
comparison to the well established methods: small amounts of sample (some 10 mg), short
experimental times (ca. 8 h), and the simplicity and accessibility of the experimental setup. In
this work we used the TGA method to measure enthalpies of vaporization of [C2mim][Cl] and
[C6mim][Cl] in order to get comparison with the indirect DSC results. The relationship between
the mass loss dm/dt and the vaporization enthalpy was derived according to the Clausius-
Clapeyron equation but by using the mass loss rate dm/dt measured by the TGA (instead of the
absolute pressure)13:
o
 glCp T
m
o
gl Hmo

T0

glCp




dm
dt
1
1
T




0 






m
(7)
ln
T
 A'


ln

R
T
T0
R
T0






Page 10 of 44

11
with a constant A´ which is essentially unknown and which includes parameters specific for our
setup but is independent from the substance studied. T0 appearing in Eq. (7) is an arbitrarily
chosen reference temperature (which here has been chosen to be 298 K) and
glCpom = Cpom
is the difference of the molar heat capacities of the gaseous Cpom
g

– Cpom
l

(8)
and the liquid phase Cpom
g

l
respectively. Enthalpies of vaporization of [C2mim][Cl] and [C6mim][Cl] derived at the average
temperature Tav of the temperature interval of the TGA studies are as follows: (Tav= 432.1)
gl Hmo
(Tav= 432.4) = 143.9 ± 2.4 kJmol-1 for
= 126.9 ± 1.8 kJmol-1 for [C2mim][Cl] and
gl Hmo
[C6mim][Cl]. In order to compare these results with the results of DSC studies values of
glCpom (Tav) have to be adjusted to the reference temperature T = 298 K with help of suitable
values of glCpom
.
3.5. Adjustment of vaporization enthalpies from TGA to 298 K. For the temperature
adjustments of vaporization enthalpies to 298 K a rough estimate of glCpom = - 100 J·K-1·mol-1 is
commonly used in the recent literature [9,16]. Our recent studies have definitely revealed that use
the constant value of glCpom = - 100 J·K-1·mol-1 regardless on structure of ILs is definitely wrong
[10,15]. Experimental studies of [Cnmim][NTf2] series have shown [15] that the value of gCo
l
p m
is not constant and it is definitely chain length dependent. The value of glCpom should decrease
with the increasing n in the alkyl chain [15]. In order to avoid the ambiguity due to the
acknowledged constant glCpom -value as well as due to the large temperature adjustment (over
100-300 K) of the vaporization enthalpy to the reference temperature 298 K, the reliable glCpom
values are required. In this work we evaluated both contributions Cpom and Cpom
in order get reasonable estimate for glCpom for each member of the [Cnmim][halide] series under
study. The experimental constant pressure heat capacities in the liquid phase Cpom
-
l


g of Eq. (8)


l
for some

representatives of the [Cnmim][Cl] and [Cnmim][Br] families were collected from the recent
review article [34]. Heat capacities of the missing homologous were interpolated with the CH2-
conribution to Cpom
Cpom

l
of 32.1 J·K-1·mol-1 developed in the same work [34]. Compilation of the

l

-values for [Cnmim][Cl] and [Cnmim][Br] is given in Table 4.
Experimental methods to measure heat capacity in the ideal gas phase, Cpom

g
, are absent.

However, it can be estimated using the modern quantum chemical methods [35]. The heat
Page 11 of 44

12
capacity of an IL in the ideal gas state could be considered as a sum of contributions from heat
capacity of a cation and a contribution of an anion. Heat capacities in the ideal gas state for
[Cnmim]+ cations at 298 K are known [35]. The contribution from the halide anion should
increase the number of vibration frequencies of [Cnmim][halide] in comparison to the spectra of
the pure cation. According to the quantum-chemical calculations of the [Cnmim][halide] the
additional bands were observed with the frequencies of 22, 56, 185 cm-1 specific for the [Cl]- and
of 22, 49, 107 cm-1 specific for the [Br]- anion. Assuming that these vibration frequencies will be
saturated at 298 K the contributions into the heat capacity of an IL in the ideal gas phase is
expected to be close to R. With this assumption and the values for [Cnmim]+ reported in the
literature, heat capacities of [Cnmim][halide] has been calculated as the sum Cpom 
gIL =
Cpom
g

[Cnmim]+ + 3R. Results are listed in the Table 5.
Having assessed both contributions Cpom and Cpom
l


g
we have used Eq. (8) to estimate

glCpom for each member of the [Cnmim][halide] series (see Table 6). As can be seen from Table
6 the glCpom -values are not far away from the commonly acknowledged glCpom = - 100 J·K-1·mol-
1
but now they fulfil our expectations – they are not constant and they are also chain length
dependent. A valuable check for validity of the procedure applied here for calculation of glCpom
provide comparison of the DSC results for [C4mim][Cl] and [C4mim][Br] reported in our
previous work [12,24] with the results from TGA and QCM studies (see Table 7). As a matter of
fact, the vaporization enthalpies derived by DSC are hardly affected by temperature adjustments
(see 3.1). We used now the new glCpom -values for [C4mim][Cl] and [C4mim][Br] given in Table
6 to adjust the measured by TGA and QCM vaporization enthalpies from Tav to 298 K. For both
gl Hmo
tested ionic liquids the agreement between direct and indirect
(298 K)-values is well within
the boundaries of the experimental uncertainties. This test convincingly demonstrates the validity
of the glCpom -values derived in Table 6 for reliable temperature adjustment of vaporization
enthalpies of ILs. We used glCpom -values estimated for [C2mim][Cl] and [C6mim][Cl] (see Table
7) and adjusted the vaporization enthalpies measured in this work by TGA from at Tav to 298 K
(see Table 7). It is quite obvious from Table 7 that the experimental vaporization enthalpies
derived from TGA are in agreement with the DSC values within the boundaries of their
experimental uncertainties.
3.6. Chain length dependence for thermodynamic properties of ILs. Properties of a
compound are a function of its molecular structure. Structures of ILs presume at least two initial
Page 12 of 44

13
patterns to study structure-property relationships: to keep the cation constant and change the
anion systematically and vice versa. Due to very restricted data only some general qualitative and
quantitative regularity possible to assess between structure of ILs and their properties. In our
recent paper we studied a series of ILs of the general formula [C2mim][X], where the
imidazolium based cation [C2mim]+ was fixed and due to a variation of the anion it was possible
gl Hmo
to follow the influence of the anion on the value of the vaporization enthalpy
and to
understand the structure-property relationships in these imidazolium based ILs [14]. Structure-
property relationships are more easy to understand for homologous series of organic compounds
and in the current work we have collected the available data for ILs of the general formula
[CnX][Cl] and [CnX][Br] with the fixed chlorine or bromine anions respectively, as well as with
imidazolium [12,24], pyridinium [25,36], and pyrrolidinium [36] cations having alkyl chains
with different length.
3.6.1. Enthalpy of IL synthesis reaction. Experimental values for reaction enthalpies,
 H (liq, 298 K), for homologous IL series [Cnmim][halide], [CnPy][halide], and
r
m
[Cn,1Pyrr][halide] are listed in Tables 8-9, and S6. It has turned out that all six sets of reaction
enthalpies were practically independent on the chain length for the alkyl substituents with
number of the C-atoms n = 4-8. In order to comprise the general regularities, we calculated the
average values for each IL family (see Table 10). It is safe to assume now from numbers
presented in Table 10 that differences in reaction enthalpies between [Cl] and [Br] containing ILs
seems to be constant at the level of about 10-13 kJmol-1 and independent on the cation. This
assessment is important for further practical application of the DSC method for similar reactions.
As a matter of fact the DSC-experiments with alkyl halides having n= 1-3 were not possible
because the precursors methyl, ethyl, or propyl halides were too volatile for handling in the scale
of milligram quantities. However, we were able to calculate thermochemical data for all “short-
chained” ILs using quantum chemical methods (see Tables 2,8,9, and S6, as well as Fig. 1).
Indeed, it is apparent from Figure 1, that gaseous reaction enthalpies,  H (g), of [Cnmim][Cl]
r
m
calculated by using the CBS-QB3 method for reactions according to Eq. (4) show a quite non-
linear dependence for n 3. Similar trend was also observed for [C mim][Br] in our previous

n
work [12]. This observation has encouraged us to assess (at least the level) of thermodynamic
properties for [C1mim][halide], [C2mim][halide], and [C3mim][halide], where the DSC
experiments were not possible. To do this, we assumed the similarity of tendencies [12] for
Page 13 of 44

14
r Hm (g, 298 K) and rHm (liq, 298 K) for n 3 (see Fig. 1 and Fig. S1). The assessed values of

r Hm (liq, 298 K) are given in Tables 2 and 3 in brackets.
In contrast to imidazolium based halides the gaseous reaction enthalpies, r Hm (g, 298 K), of the
pyridinium and pyrrolidinium based halides are hardly chain-length dependent even for for n
3

(see Fig. 2). This observation has prompted to set constant values for the “short chained” ILs
 H (liq, 298 K) = -87.0 kJ mol-1 for [C Py][Br] and  H (liq, 298 K) = -78.9 kJ mol-1 for


n
r
m
r
m
[Cn,1Pyrr][Br] with n= 1-3. Also the constant values for the “short chained” ILs  H (liq, 298
r
m
K) = -65.4 kJ mol-1 for [C Py][Cl]  H (liq, 298 K) = -76.5 kJ mol-1 for [C Pyrr][Cl] with n=


n
n,1
r
m
1-3 were taken from our previous DSC study [36]. We are conscious of a tentative quality of
these guesstimations for the “short-chained ILs. Anyway, in our opinion these reaction enthalpies
are reliable within ± 5 kJ mol-1 and can be applied for further assessment of enthalpies of

formation and enthalpies of vaporization as well as for optimization of chemical engineering
processes with ILs.
3.6.2. Enthalpy of formation in the liquid state. Correlation of enthalpies of formation with the
number of C-atoms (n) within homologues series is a valuable test to check an internal
consistency of experimental results. As a rule, for molecular organic compounds values of liquid
 f Hm
and gaseous enthalpies of formation,
, [40] depend linearly on the n (see Fig. S2, ESI)
except for the first representative of the series, e.g. for alkyl chlorides, the enthalpy of formation
of methyl chloride lies slightly out of the general linear correlation (see Fig. S3, ESI). Correlation
of the  f Hm (liq) with n for ILs under study is presented on Figure 3. The apparent linear
dependences for all IL-series revealed that  f Hm (liq) for ionic and molecular liquids show very
similar behaviour. This linearity seems to be a good prove for the data consistency, as well as
provides an additional confidence for our guesstimations for the “short-chained” ILs (see 3.6.2).
The dependences of liquid phase enthalpies of formation on the total number of C-atoms for n

2 are expressed by the following equations:
-1

.
[Cnmim][Cl]:  f Hm (liq, 298 K)/ kJ mol = -80.7 - 27.4 n
(r= 0.997)
(r= 0.993)
(r= 0.999)
(r= 0.998)
(r= 0.999)
(9),
-1

.
[Cnmim][Br]:  f Hm (liq, 298 K)/ kJ mol = -52.2 - 26.8 n
(10),
(11),
(12),
(13),
-1
 f Hm
 f Hm
 f Hm
.
[CnPy][Cl]:
[CnPy][Br]:
[Cn1Pyrr][Cl]:
(liq, 298 K)/ kJ mol = -57.0 - 26.4 n
-1
.
(liq, 298 K)/ kJ mol = -25.5 - 25.9 n
-1
.
(liq, 298 K)/ kJ mol = -181.9 - -26.4 n
Page 14 of 44

15
-1

.
[Cn1Pyrr][Br]:  f Hm (liq, 298 K)/ kJ mol = -153.9- 25.8 n
(r= 0.997)
(14),
from which enthalpy of formation  f Hm (g, 298.15 K) other representatives of these two series
with n > 2 can be calculated.
It is not completely clear yet whether enthalpies of formation of ionic liquids obey the group
additivity rules [37,38]. The slopes of the linear dependencies given in Eqs. (9-14) represent the
contribution of the CH2 group to  f Hm (liq). Comparison of slopes for all six families reveals
-1
.
remarkable similarity around of -26 kJ mol and this value seems to be essentially the same as
-1
.
acknowledged for the molecular compounds (-25.5 kJ mol ) [39]. This observation would be
 f Hm
very helpful for development of any group-contribution scheme for prediction of
(liq) of
ILs.
gl Hmo
3.6.3. Enthalpy of vaporization. Do we also expect the linear dependence of
on the chain
length for ILs? Before we discuss this question regarding the ionic liquids, let us consider the
general tendencies in the molecular compounds – e.g. the alkyl halides, which are reaction
gl Hmo
participants. For the molecular liquids, it is well established, that vaporization enthalpies
appear to be a linear function of the number of carbon atoms of the alkyl chain in alkanes [41],
gl Hmo
alcohols [42], aliphatic esters [43] and aliphatic nitriles [44]. The plot of
(298 K) against
the number of C-atoms in the linear aliphatic halides CH3-(CH2)n-Hal with n = (3 to 15) [45] is
presented in Figures S4 and S5. However, the first two representatives – methyl and ethyl halides
are slightly out of the linear correlation. Such a peculiarity was also observed for aliphatic esters
[43] and aliphatic nitriles [44] and this fact might be caused by the high dipole moment of these
species.
For ionic liquids homologous series the vaporization enthalpy is expected to increase
monotonically in accordance with the increase in the number CH2 groups. However, the
vaporization behaviour of an IL is apparently affected due to interplay of the van der Waals and
Coulombic interactions. The expected increase of the van der Waals interactions with the chain
length (as observed for pure n-alkanes) for a certain extend have to be impacted by electrostatic
forces and the energy required for the transfer of the contact ionic pair to the gaseous state (in
general – vaporization enthalpy) could become unique for each representative of the IL series.
Contrary to this expectations a remarkable linear dependence of vaporization enthalpy on n for
[Cnmim][NTf2] with n = 2, 4, 6, 8, 10, 12, 14, 16, and 18, when measured by using the quartz-
crystal microbalance (QCM) [10]. Also the linear chain length dependence was observed for the
[Cnmim][Br] series, when derived indirectly by using DSC [12]. It has turned out that only the
Page 15 of 44

16
first representative of [Cnmim][Br] series – methyl bromide was slightly out of the linear
correlation (see Figure S6, ESI) in accord with the expected counterplay of the van der Waals
and Coulomb interactions in the IL. In this work we have systematically derived vaporization
enthalpies for [Cnmim][halide], [CnPy][halide], and [Cn,1Pyrr][halide] families and plotted them
with n on Figure 4. As can be seen from this figure, the linearity of vaporization enthalpy with
the growing chain length on the cation seems to be a general feature for ionic liquids,
independent of the cation and anion structures. The dependence of vaporization enthalpy on the
total number of C-atoms for n  2 is expressed by the following equation:
-1
gl Hmo
gl Hmo
gl Hmo
gl Hmo
gl Hmo
gl Hmo
.
[Cnmim][Cl]:
[Cnmim][Br]:
[CnPy][Cl]:
(298 K)/ kJ mol = 133.1 + 4.2 n
(r= 0.980)
(r= 0.988)
(r= 0.941)
(r= 0.985)
(r= 0.991)
(r= 0.983)
(15),
(16),
(17),
(18),
(19),
(20),
-1
.
(298 K)/ kJ mol = 136.6 + 4.1 n
-1
.
(298 K)/ kJ mol = 157.1 + 3.6 n
-1
.
[CnPy][Br]:
(298 K)/ kJ mol = 155.7 + 4.4 n
-1
.
[Cn1Pyrr][Cl]:
[Cn1Pyrr][Br]:
(298 K)/ kJ mol = 96.5 + 4.2 n
-1
.
(298 K)/ kJ mol = 101.3 + 4.1 n
gl Hmo
from which enthalpy of vaporization
these series can be calculated.
(298 K) for other longer chained representatives of
It also worse mentioning, that Eqs. (15-20) demonstrate that the contribution to vaporization
-1
.
enthalpy of about 4.0 kJ mol per each CH2 group is apparently constant for all ILs under study.
-1
.
But this contribution is slightly lower in comparison with those of 4.95 kJ mol per each CH2
group typical for the n-alkane family [41]. An additional important prove for consistency of the
gl Hmo
new
(298 K) data sets for chlorine and bromine based ILs could be a direct comparison of
the differences in vaporization enthalpies for [Cl] and [Br] containing species. Such a
comparison was performed for imidazolium (Table 11), pyridinium and pyrrolidinium (Table 12)
based ILs. As can be seen from these tables, the differences for the molecular liquids
gl Hmo
(CnBr) -
(CnCl) are at the average level of about 3.0 kJ.mol-1. For the ionic liquids
gl Hmo
the differences between [Cl] and [Br] also similarly spread close to the average level of 3-4
-1
.
kJ mol . Somewhat large spread of the differences for ILs in comparison to the molecular
-1
.
compounds is quite explainable taking into account the uncertainties of ±1-3 kJ mol typical for
the DSC measured reaction enthalpies (see Table 1).
3.6.4. Enthalpy of formation in the gas state. Having established the linear correlations of
liquid phase enthalpies of formation as well as vaporization of ILs with the number of C-atoms n
within the homologues series it was logical to get the correlation for gas enthalpy of formation as
well. The dependence of the gas enthalpy of formation on the total number of C-atoms in the
alkyl chain for n 2 is expressed by equations:

Page 16 of 44

17
-1

.
[Cnmim][Cl]:  f Hm (g, 298 K)/ kJ mol = 46.1 - 22.1 n
(r= 0.999)
(r= 0.999)
(r= 0.999)
(r= 0.999)
(r= 0.999)
(r= 0.997)
(21),
(22),
(23),
(24),
(25),
(26),
-1

.
[Cnmim][Br]:  f Hm (g, 298 K)/ kJ mol = 76.5 - 21.3 n
-1

.
[CnPy][Cl]:  f Hm (g, 298 K)/ kJ mol = 96.1 - 22.2 n
-1

.
[CnPy][Br]:  f Hm (g, 298 K)/ kJ mol = 25.9 - 20.8 n
-1

.
[Cn1Pyrr][Cl]:  f Hm (g, 298 K)/ kJ mol = -89.3 - 21.6 n
-1

.
[Cn1Pyrr][Br]:  f Hm (g, 298 K)/ kJ mol = -57.4 - 21.0 n
These equations are useful to predict values for other representatives of these series with n > 8.
Analysis of slopes in Eqs. (21-26) reveals that the contribution from CH2 group in the alkyl chain
-1
.
of ionic liquids is close to -21 kJ mol and it seems to be also essentially the same as for
-1
.
molecular compounds (-20.9 kJ mol ) [39]. We could conclude now that group additivity rules
are valid for the gaseous enthalpies of formation of the imidazolium, pyridinium, and
pyrrolidinium series of ionic liquids with contributions for alkyl chain on the cation similar to
those in molecular compounds. However, more experimental information is required for
detecting similarities or disparities of additivity parameters of molecular and ionic compounds.
3.6.5. Enthalpy of formation in the crystalline state. Using Eqs. (9-14) it was possible to
estimate enthalpies of formation  f Hm (liq, 298 K) in the liquid state. However, the most ionic
liquids under study are not liquids at room temperature. They are solids with melting points
above 298 K up to 600 K (a high melting temperature is typical for the first representatives of ILs
homologous series). From the basic thermodynamic knowledge, enthalpy of formation in the
crystalline state related to the liquid state over the enthalpy of fusion:
 f Hm (cr) =  f Hm (liq) - lcr Hm .
(27)
l
Unfortunately the data on the molar fusion enthalpies cr Hm at Tfus are scarce [46-54]. But it
could be calculated using the modified Walden’s rule [52]:
-1
-1
l
.
cr Hm (Tfus) = 54.4 (J K mol ) ·Tfus (K)
(28)
We checked this rule for ILs with help of available temperatures and enthalpies of fusion for
alkylimidazolium based ILs with [Cl], [Br], and [NTf2]. Thus, the molar enthalpy of fusion of an
IL of interest could be roughly assessed within ±3 kJ mol-1 with help of Eq. (28), provided that

the fusion temperature is known. Systematic search in the open and the patent literature have
revealed very limited and often ambiguous data on the melting points [50-54]. We collected all
available melting temperatures in Table S9. For some ILs we had to assess missing data
according to general trends (Table S9, column 2, data given in parentheses). Collection of the
Page 17 of 44

18
experimental and estimated by eq. 28 data on the molar enthalpies of fusion given in Table S9,
column 3). However, because of the differences between Tfus and the reference temperature, the
experimental enthalpies of fusion (referred to Tfus) have to be adjusted to T = 298 K (see details
in ESI) and they are given in Table S9, column 4. Analysis of the lcr Hm (298 K) results given in
Table S9, column 4 has revealed that fusion enthalpies of all ionic liquids under study adjusted to
-1
.
the reference temperature 298 K are generally close to the average valued 15±3 kJ mol
regardless on the structure and the melting temperature. This rough estimate could be very useful
in the future for a quick appraisal of lattice enthalpy of ILs not studied yet. When adjusted to 298
K, the values lcr Hm were used to derive molar enthalpies of formation of ILs in the crystalline
state,  f Hm (cr, 298 K), according to Eq. (27) (see Table S9, column 6).
 f Hm
(cr, 298 K),
3.7 Lattice potential energy of ILs from DSC-measurements. With help of
it is possible now to assess a lattice potential energy (UPOT) and lattice enthalpy ( H ) of an ionic

L
solid material. UPOT is a measure of the strength of bonds in that ionic compound. Lattice
potential energy is usually defined as the enthalpy of formation of the ionic compound from
gaseous ions [26-30] (for example for [C4mim][Br]):
+
-



UPOT([C mim][Br])=
([C mim] ,g) +
([Br] ,g) – 2·RT –
([C4mim][Br],cr)(29)
 f Hm
 f Hm
 f Hm
4
4
Lattice potential energy is a dominant term in the thermodynamic analysis of the existence and
stability of ionic solids. Direct experimental determination is generally not possible since, in
practice, the crystalline solid dissociates into atoms and not into gaseous ions, as is required in
the lattice energy evaluation. Therefore, its indirect experimental determination, computation, or
estimation is of considerable interest in modern materials science. Indeed, whenever the
energetics of condensed-state materials are studied, the chemical processes under consideration
may be rationalized if the appropriate lattice energy steps can be incorporated into the
thermochemical cycle [27]. By using experimental data, empirical correlations, and quantum
chemical calculations, Born-Fajans-Haber cycles can be constructed to gain a deeper insight into
reaction or phase transition thermodynamics [28,29]. A simple “volume-based thermodynamics”
(VBT) procedure was developed recently for the evaluation of UPOT, which relies only on the
chemical formula, ionic charges and molecular volume (or density) of the ionic compounds [26-
30]. However, for a proper validation of the VBT-equation the extended data set of the
 f Hm
experimental
(cr) data is needed.
Page 18 of 44

19
In this work we calculated gaseous enthalpies of formation of cations  f Hm ([cation]+,g) and
anions  f Hm ([anion]-,g) by using the CBS-QB3 method (see Table 13). Lattice potential
energies were calculated according to Eq. (29) and given in Table S9, last column. It has turned
out (see Fig. 5) that the values of UPOT also linearly dependent on the chain length n, similar to
other thermodynamic properties discussed in this work (see Fig. 3 and 4). The simple structure-
property relations for UPOT observed in this work is important as a prove for consistency of the
experimental procedures and assumptions used in this work.
Conclusions
Alliance of experimental and computational methods has been used to derive
thermochemical values for homologous series of imidazolium, pyridinium, and pyrrolidinium
based ionic liquids. These values are key properties required for temperature management and
optimization of the industrial synthesis of ILs. We have shown that enthalpies of formation,
enthalpies of vaporization, and lattice potential energies are linearly dependant on the alkyl chain
length. These thermodynamic properties generally obey the group additivity rules and the values
of the additivity parameters for ionic liquids seem to be very close to those for molecular
compounds.
Supporting Information. Table S1-S2. The results of the temperature dependence of mass loss
for [C2mim][Cl] and [C6mim][Cl] in TGA determination; Table S3. Thermochemical data for
precursors discussed in this work; Table S4-S5. Results of ab initio calculation of  f Hm and
experimentally values of Δgl Hm for alkyl halides; Table S6. Thermodynamic data for
[Cnmim][Br] synthesis reaction calculated using the CBS-QB3 composite method; Table S7-S8.
Results of Calculation of the Enthalpy of Formation of Br-containing ILs and compounds in the
Gaseous Phase at 298 K; Table S9. Data on melting temperatures and enthalpies of fusion of ILs;
o
m
Figure S1. Enthalpies of reaction, H , of [C mim][Br] synthesis reaction; Figure S2-S5.
r
n
Enthalpies of vaporization and formation of alkylhalides of various alkyl lengths; Figure S6.
Enthalpies of vaporization (298 K), of [Cnmim][Br]; This material is available free of
Δgl Hm
charge via the Internet at http://pubs.acs.org.
Acknowledgements
This work has been supported by the German Science Foundation (DFG) in the frame of the
priority program SPP 1191 “Ionic Liquids”.
Page 19 of 44

20
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25
10
0
[Cnmim][Br]
[Cnmim][Cl]
-10
-20
-30
-40
0
2
4
6
8
10
n
Fig. 1. Enthalpies of reaction of 1-methylimidazole with of alkylhalides of various alkyl lengths
calculated using CBS-QB3.
Page 25 of 44

26
40
30
20
10
[CnPy][Br]
0
[CnPy][Cl]
[Cn1Pyrr][Br]
[Cn1Pyrr][Cl]
-10
-20
-30
-40
0
2
4
6
8
10
n
Fig. 2. Enthalpies of reactions of pyridine and N-methyl-pyrrolidine with alkylhalides of various
alkyl chain lengths (data from Table 8 and 9).
Page 26 of 44

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210
190
170
150
130
110
90
[Cnmim][Cl]
[Cnmim][Br]
[CnPy][Cl]
[CnPy][Br]
[Cn1Pyrr][Cl]
[C1nPyrr][Br]
0
1
2
3
4
5
6
7
8
9
n
Fig. 3. Correlation of the vaporization enthalpies of ILs under study on the chain lengths n in the
cation.
Page 27 of 44

28
50
-50
[Cnmim][Cl]
[Cnmim][Br]
[CnPy][Cl]
[CnPy][Br]
[Cn1Pyr][Cl]
[Cn1Pyrr][Br]
-150
-250
-350
-450
0
2
4
6
8
10
n
f Hm
(liq), of ILs under study on the chain
Fig. 4 Correlation of the enthalpies of formation,
lengths n in the cation.
Page 28 of 44

29
600
580
560
540
520
500
480
460
[Cnmim][Cl]
[Cnmim][Br]
[CnPy][Cl]
[CnPy][Br]
[Cn1Pyr][Cl]
[Cn1Pyr][Br]
1
2
3
4
5
6
7
8
9
10
n
Fig. 5. Lattice potential energies of ILs under study as the chain length n dependence (data from
Table S9).
Page 29 of 44