1. Introduction
In the last two decades, ionic liquids (ILs) have been considered as potential solvents to replace volatile organic solvents in chemical technologies, as they possess negligible vapour pressure and they cannot evaporate (even at elevated temperatures) and cause air pollution. Typical protic ILs are comprised of a quaternary cation (imidazolium, ammonium, phosphonium, and pyridinium, etc.) with a wide variety of common anions. Their physical and thermodynamic properties can be deliberately tuned by the choice of the cation, anion, and substituents implemented into the cation or anion structure. To develop ILs for the various practical applications, it is important to gain a fundamental knowledge of the factors that control the volatility and thermal stability of pure ILs, as well as the phase behaviour of ionic liquids with other common solvents and chemicals. Since 2001, our laboratory has contributed to this fundamental understanding and prediction of IL properties. We have launched our work with the development of the gas-chromatographic method for determination of infinite-dilution activity coefficients [
1]. Activity coefficients at infinite dilution of a solute
in ionic liquid are important for the optimal design of separation processes. They are especially important when the last traces of impurities must be removed. Knowledge on activity coefficients helps to avoid an oversizing of distillation columns or stripping processes. Moreover,
values provide information about the intermolecular interactions between solvent and solute; in particular, they are used for the selection of solvents for extraction and extractive distillation.
From the very beginning of the ILs era, there were two crucial advantages propagated for ILs: negligible vapour pressure and inflammability. Both predications have been challenging for our thermochemical lab specialised on vapor pressure measurements and combustion calorimetry. Thus, the long-standing concept of ILs involatility has been disproved with our first quantitative measurements on ILs vapor pressure [
2]. Additionally, the acknowledged postulate of ILs inflammability has been disproved by our very first combustion energy measurements of 1-butyl-3-methylimidazolium dicyanamide [
3]. Furthermore, we suggested the combination of calorimetric methods (combustion calorimetry and differential scanning calorimetry [
4]) with quantum-chemical and molecular dynamic calculations for the “indirect” appraisal of ILs vaporization enthalpies [
5]. At the same time, we have intensively developed “direct” experimental methods for ILs vapor pressures and vaporization enthalpy measurements. The main challenges for this method are that, at ambient temperatures, the extremely low vapor pressures of ILs are practically negligible and difficult to measure, whereas, at elevated temperatures, the vaporization process is very often aggravated by the thermal degradation. Nevertheless, we optimized experimental conditions of the conventional transpiration method for the reliable determination of the ILs’ vapor pressures temperature dependences [
3,
5]. Moreover, we demonstrated that the thermogravimetric analysis (TGA) can be successfully used not only for the common thermal stability studies, but also for the reliable vaporization enthalpy determinations at elevated temperatures [
6]. Indeed, it was found that thermal stability of many ILs was not always sufficient to exclude decomposition in the temperature range of the TGA study. In order to overcome the limitations due to the thermal instability of the ILs, we have developed an extremely sensitive quartz-crystal microbalance (QCM) method to determine the ILs’ vapor pressures [
7]. The enormous sensitivity of the quartz-crystal (placed over a cavity filled with IL) for the detection of a minute amount of ionic liquid, vaporised and deposited on the surface, has allowed a drastic drop in experimental temperatures (e.g., a starting temperature of the QCM-study of 353 K!) feasible for reliable determination of the mass loss and elimination of thermal decomposition as a consequence. Confronting the ever-questionable thermal stability of ILs at elevated temperatures, we have also invented a method for determining ILs’ volatility at the nanoscale by means of ultra-fast scanning calorimetry (UFSC) [
8]. The idea behind this method is to reduce the residence time of the IL sample at all experimental temperatures. It can be achieved by extremely high heating and cooling rates (up to 10
5 K s
−1) of the sample placed on the sensor developed for the heat capacity measurements. Thus, the residence time of the IL at high temperature required for measurable mass uptake was optimised to some minima. One consequence was that the chemical processes of thermal degradation hardly have time to begin, as the sample is quickly cooled to ambient temperature. With the UFSC method (based on extremely high heating/cooling rates), much higher experimental temperatures (as compared to common techniques) can be reached without significant decomposition. For example, it was demonstrated that evaporation of 1-ethyl-3-methylimidazolium bis(trifluoromethyl-sulfonyl)imide ([C
2mim][NTf
2]) at temperatures of up to 750 K is still the dominating process of mass loss, even at such highly elevated temperatures [
8]. Despite the success of the suppression of decomposition during the TGA, QCM, and UFSC measurements, the idea of decreasing the temperatures of IL studies seems to be more attractive and practically relevant. Complementary to the development of the low-temperature QCM-method, we have come up with an idea to use alternating current chip calorimetry (AC) for vapor pressure measurements of ILs at low temperatures [
9]. A small droplet of an IL is vaporized isothermally from the chip sensor in a vacuum chamber. The highly sensitive chip sensor (SiN
x-membrane) allows for mass loss determination at temperatures also starting from 350 K. Finally, it is worth mentioning that the data set on vapor pressures, created for [C
2mim][NTf
2] from experiments with AC and UFSC methods for the extremely broad temperature range from 358 K to 780 K, has allowed the estimation of the boiling temperature of this IL. The value (1120 ± 50) K should be considered as the first reliable boiling point of the archetypical ionic liquid obtained from experimental vapor pressures measured in the closest possible proximity to the normal boiling temperature [
9].
The systematic development and mutual validation of experimental methods for vapor pressure and vaporisation enthalpies,
, of the extremely low-volatile ionic liquids delineated above, have been a focus of our scientific interests for the past two decades. As a result, consistent sets of experimental vapor pressures and vaporization enthalpies for imidazolium- [
1,
2,
3,
4,
5,
6,
7,
8,
9,
10,
11,
12,
13], pyridinium- [
14,
15,
16,
17], pyrrolidinium- [
17,
18], and phosphonium- [
19] based ionic liquids have been reported in the current literature. There is a broad demand for these data, as they are essential for the effective utilization of ILs such as thermofluids, or in separation processes, where precise thermodynamic data on solvents and working fluids are required. Absolute vapor pressures of ILs are indispensable for their modern catalytic applications such as Solid Catalysts with Ionic Liquid Layer (SCILL) [
20] or a supported ionic liquid phase (SILP) [
21]. The vital question in both cases is: is the IL vapor pressure in the catalytic temperature range low enough in order to avoid the uptake from the catalyst layer? The correct answer is important for the large-scale applications of the generally expensive ionic liquids.
Apart from the practical importance, vaporization enthalpy is one of the key values required for parameterization and optimization of the ILs’ force field for molecular dynamics simulations [
22]. Moreover, the systematic information on vaporization energetics within well-defined ionic liquid families opens the way for fundamental understanding of the structure–property relationships in ILs, essential for a reliable prediction of their thermodynamic properties [
23]. The latter values are crucial for the optimisation of IL synthesis reactions [
24] or IL use for the synthesis of nanoparticles [
25].
Admittedly, all standard thermodynamic properties used in practical thermochemical calculations are referenced to any common temperature (most frequently
T = 298.15 K). As is apparent from the survey of the experimental method for vaporization enthalpy determination outlined above, none of these methods provide the
(298.15 K) values. As a rule, the measured vaporization enthalpies are referenced to the average temperature
Tav of the experimental interval. These
(
Tav) values have to be adjusted to the reference temperature
T = 298.15 K according to the Kirchhoff’s Rule [
2,
3,
10] by using appropriate
values. Formally, the value
=
(g) −
(l) is the difference of the molar heat capacities of the gaseous
(g) and the liquid phase
(l), respectively. However, the ambiguity and prospects of the
values required for the adjustment of experimental vaporization enthalpies
(
Tav) to the reference temperature
T = 298.15 K have already been discussed in detail [
10]. Nevertheless, any new idea for independent assessment of
(298.15 K) values is valuable and desired.
Fortunately, in a series of our recent IL studies [
26,
27,
28] we have been able to show that the
(298.15 K) values can be derived from the collection of activity coefficients at infinite dilution
of a solute
1, in an ionic liquid (solvent
2) measured by using gas–liquid chromatography (GLC). These studies [
26,
27,
28] were performed on imidazolium based ILs with the trifluoroacetate [CF
3CO
3], trifluoromethanesulfonate [CF
3SO
3], and methanesulfonate [CH
3SO
3] anions. Is this observation also generally valid for imidazolium based ILs containing other types of anions? If yes, the method based on the experimental
values can be suggested as an independent source for the
(298.15 K) values, as well as for the valuable validation of
values applied for the temperature adjustment of experimental
(
Tav) values. Our pioneering paper on GLC determination of activity coefficients of aliphatic and aromatic compounds in 4-methyl-
N-butyl-pyridinium tetrafluoroborate [
1] has inspired the ionic liquid community. Indeed, according to the most recent compilation [
29], a comprehensive experimental data base covering 233 ILs and 150 molecular solutes was extracted from 182 references, dealing with activity coefficients at infinite dilution. This huge data pool could be considered as an independent source for the
(298.15 K) values for numerous differently structured ILs, provided that the procedure for the recovery of vaporisation enthalpies from solute
values is successful. In order to keep this task in size, in this study we selected only imidazolium-based ILs where the
(298.15 K) values were measured and evaluated in our recent papers.
3. Theoretical Background
In this work we follow the Flory–Huggins theory, which is the main basis of solution and blend thermodynamics [
33]. The Flory–Huggins equation handles molecules that are similar chemically, but differ greatly in size. A key value of this theory is a parameter
χ12 quantifying the enthalpic interactions between the components
1 and
2. The activity coefficient at infinite dilution
is linked to the Flory−Huggins interaction parameter
χ12 (at infinite dilution) according to the equation [
34]:
where
M1 and
M2 are the molecular weight of solute and solvent, respectively, and
V1⁎ and
V2⁎ and
ρ1 and
ρ2 are the molar volume and density of solute and solvent, respectively. The Flory–Huggins interaction parameters,
χ12 is related to the Hildebrandt solubility parameters
δ [
34]:
where
δ2 is the solubility parameter of the IL (solvent) and
δ1 is the solubility parameter of the solute,
R is the universal gas constant,
T is the arbitrary temperature, and
is the molar volume of the solute at the selected temperature. Solubility parameters are the numerical values that are responsible for the strength of the intermolecular interactions between solute and solvent molecules. The solubility parameters have been widely used as they help to assess the solvation powers of solvents.
The Hildebrand or total solubility parameter (
δi) is defined as follows [
35]:
where
Vm is the molar volume,
is the standard molar enthalpy of vaporization,
R is the ideal gas constant, and
T is the temperature. Vaporization enthalpies,
, of ionic liquids required for calculations
δ2 at
T = 298.15 K have been systematically evaluated in our recent work. Vaporization enthalpies,
, of molecular solutes required for calculations
δ1 at
T = 298.15 K were taken from the literature [
36,
37]. Values of the solubility parameters
δ1 and
δ2 have been calculated according to Equation (4) with help of experimental data on vaporization enthalpies for the solutes [
36]. Density values for the solutes were taken from the compilation by Lide [
37], and for ILs from the NIST Database [
38].
Finally, the algebraic rearrangement of Equation (3) gives:
As it was shown in a typical case (see
Figure 1), when the left side of Equation (5) is plotted against
δ1, then the mathematical term 2
δ2/(
RT) is the slope of the line and the module −
/(
RT) is its intercept. Using linear regression of the experimental data, the slope or intercept can be used to determine
δ2.
We fitted Equation (5) with the solubility parameters
δ2 derived from primary
values for the [C
nmim][Anion] available in the literature. For each data set, the solubility parameters
δ2 obtained using the slope and intercept were in agreement with one another within 3%. Thus, the
δ2 values can be estimated as the average from the slope and the intercept. Thus, the vaporization enthalpy,
(298.15 K), of an IL under study was calculated using the averaged
δ2 value as follows:
where all values, including
Vm, are referenced to an arbitrary temperature
T, which is 298.15 K in this work. Now these
(298.15 K) results, “indirectly” derived with help of the primary
-values, can be used for comparison with the “direct” experimental results on vaporization enthalpies obtained by the conventional methods.
4. Inspection of Activity Coefficients Placed at the Disposal
Experimental
values for different solutes in different types of ILs are regularly appearing in the literature, beginning in 2001 with the aprotic ionic liquids [
1] and, ten years later, the protic ionic liquids [
40,
41]. As a rule, the set of
values reported for a particular ionic liquid consists of 15 ÷ 40 solutes of different polarity. In the majority of cases, solutes are subdivided into groups:
n-alkanes with the chain-length C
5–C
10, cycloalkanes with the chain-length C
5–C
8, alkenes with the chain-length C
5–C
8, alkynes with the chain-length C
5–C
8, benzene and alkylbenzenes with the chain-length C
1–C
3, and linear and branched aliphatic alcohols with the chain-length C
1–C
5, as well as a number of solutes of different polarity: acetone, acetonitrile, thiophene, tetrahydrofuran, aliphatic ethers, and esters. In addition, the water is also included if the studies are performed with the thermal conductivity detector. A typical example of the regression of solubility parameters
δ298.15 derived from the experimental
values is given in
Figure 1. For demonstration, we have deliberately chosen the data set containing water (see
Figure 1Left). It is apparent that the solubility parameter of water due to the exceedingly high polarity is totally out of correlation with all other types of solutes. Therefore, it is senseless to keep this molecule in the set for regression of solubility parameters. The same data set but without water looks more appropriate (see
Figure 1Right), but in this iteration it is obvious that subsets of linear and cyclic aliphatics (alkanes, alkenes, and alkynes) definitely demonstrate individual behavior, and the
δ298.15 points represent a cloud rather than a decent linear attitude. It is noticeable that the cyclical molecules deviate even more drastically from the general trend. We could suggest at least two possible reasons for the variations observed. The first reason is an objective one—the individual behavior of alkane, alkene, and alkyne series can be considered as evidence of specific intermolecular interactions between the different types of aliphatics and the IL under study. Compounds with the double or triple bond are expected to interact with the polar framework of the IL more intensively than the similarly shaped alkanes. Additionally, the cyclic molecules interact with the IL differently than linear molecules, containing the same number of C-atoms. This explanation is in accordance with the general thermodynamic interpretation of the
values, which are meaningful for interpretation of intermolecular interactions between solutes and solvents.
The second reason for the significant spread of the
δ298.15 points on
Figure 1 is a rather subjective one. From our own experiences, the
measurements are thwarted with considerable technical complications. The main troubles are due to the very short retention times of the highly volatile aliphatics. Even at the low temperatures of the GC experiment, and by using the long packed columns, the retention times of the C
5–C
7 hydrocarbons are close to the “dead time”, and the
values can be significantly affected by the inaccuracy of the time registration. Moreover, in the GC experiment, the C
5–C
7 hydrocarbons usually eluate in the immediate vicinity after the solvent peak (e.g., CH
2Cl
2). Admittedly, the tail of the solvent peak still contains sufficient residual amounts of solvent molecules dissolved in the IL layer covering the solid support. The solute of interest (one of C
5–C
7 hydrocarbons) not only interacts with the IL, but also with the residual CH
2Cl
2 molecules and, as a consequence, the true retention time is counterfeit. Thus, the unavoidable solvent peak tailing also contributes to the inaccuracy of the time registration. The practical conclusion for the further
values acquisition is that the experimental points for C
5–C
6 hydrocarbons (alkanes, alkenes, and alkynes) should be omitted.
In contrast to the very volatile C
5–C
6 hydrocarbons, the retention times measured for a series of different polar compounds (e.g., acetone, benzene, alkylbenzenes, and ethers) are not affected by the experimental perturbations. Additionally, a reasonable correlation of solubility parameters
δ298.15 of different solutes with the
Y-module has been virtually observed (see
Figure 2Left). However, such a good correlation is rather due to the fact that the series of polar compounds involved in the examination belongs to a relatively narrow range of 0.1 ÷ 0.3 units, according to the normalized solvent polarity scale [
42]. For this reason, the moderate intensity of intermolecular interactions between each solute with the IL-solvent is localized at a comparable level, reflecting the individual straight line for the polar compounds in
Figure 2Left.
The energetics of interactions of alcohol molecules with the IL is expected to be more profound, as their polarities range from 0.75 (methanol) to 0.60 (1-propanol) in units of the normalized solvent polarity scale [
42]. Subsequently, the ROH series (methanol, ethanol, 1-propanol, and 1-butanol) also represents the individual straight line (see
Figure 2Left).
5. Development of Activity Coefficients Data Acquisition and Processing
We have deliberately launched the discussion of the
measurements with the most consistent data set reported for the for [C
6mim][SCN] [
39]. However, among the available literature compilations on the
values for imidazolium based ILs of the general formula [C
nmim][Anion], we have dealt with many less consistent data sets. For example, the
values for series including benzene and alkylbenzenes, alcohols (1-propanol, 2-propanol, and 2-methyl-1-propanol), polar solutes (acetone, acetonitrile, ethyl acetate, 1,4-dioxane, and tetrahydrofuran), and halogen containing solutes (dichloromethane, trichloromethane, 1,2-dichloromethane, chlorobenzene, and bromobenzene) were reported for [C
6mim][CF
3CO
2] [
43].
The regression of solubility parameters
δ298.15 of different solutes with the
Y-module for [C
6mim][CF
3CO
2] is presented in
Figure 3. Numerical values used for this correlation are collected in
Table S2. The significant spread of the experimental data points is apparent on
Figure 3, especially for the polar compounds such as acetone, acetonitrile, 1,4-dioxane, and trichloromethane. Most probably, this scatter is an indication of more pronounced intermolecular interactions between these solutes with the [C
6mim][CF
3CO
2] in comparison to [C
6mim][SCN]. Certainly, one or the other outlier can be cancelled in order to minimize scattering. However, each removed point inevitably leads step by step to the significant change of the slope and intercept of the regression presented on
Figure 3. As a consequence, the
δ2-result (and finally the
(298.15 K)-result, calculated according to Equation (6)) becomes questionable and virtually dependant on manipulations with the data set on
values. Moreover, the resulting
(298.15 K) value for the IL under study becomes crucially dependent on the collection of
values selected for the evaluation. This conclusion is easy to support with numerical results obtained from
Figure 2Right. The value
(298.15 K) = 115.1 kJ⋅mol
−1 is calculated using the selection of “polar” solutes. The value
(298.15 K) = 152.8 kJ⋅mol
−1 is calculated using only the set of alcohols. Seemingly, such a significant disagreement leaves no hope for practical application of
values for the
(298.15 K) appraisal. Nevertheless, let us recall and rationalise the background of the disagreement once more. In general, each available
data set comprises a series of non-polar solutes (e.g.,
n-alkanes), weakly polar solutes (e.g., alkylbenzenes, ethers, and esters, etc.), and strongly polar solutes (usually alcohols). Each subset exhibits specific intensity of intermolecular interactions with the IL-solvent. Apparently, the slopes of the
δ1(298.15 K)–
Y regressions (see
Figure 1,
Figure 2 and
Figure 3) constructed for each subset of
-values could be helpful to quantify the specific energetics of these forces. Unfortunately, the scales of the
δ1(298.15 K) values and
Y values for non-polar and weakly polar solutes are impractically narrow. For example, the variations of
δ1(298.15 K) values for aliphatics are only from 15.2 to 17.4 MPa
0.5 (see
Table S1), and the variations for
Y values for the same series are from 0.0583 to 0.0989 unities (see
Table S1) [
26]. As a rule, slopes derived in restricted ranges of parameters are very sensitive and suffer from the individual uncertainties of experimental data points. The variation of
δ1(298.15 K) values and
Y values for series of weakly polar solutes is somewhat broader (see
Table S1), but still insufficient for establishment of robust correlations. Additionally, only the scale of
δ1(298.15 K) values and
Y values for alcohols ROH with R = methyl, ethyl,
n-propyl, and
n-butyl (see
Table S1 and
Figure 1Right) is adequate for the reliable regression (e.g., the variations of
δ1(298.15 K) values is from 29.4 to 23.3 MPa
0.5, as well as the variations for
Y values is from 0.3450 to 0.2155 unities (see
Table S1) [
26]. It is only regrettable in that papers where at least four C
1–C
4 alcohols are measured are quite restricted, precluding any meaningful interpretation of results for this series.
The temporary conclusion for the continuation and development of
values evaluation is that the none of individual subsets of solutes can be used for a reliable assessment of vaporisation enthalpies
(298.15 K). However, what speaks against combination of two subsets (see
Figure 2Right)? Indeed, by the merging the
n-alkanes subset with the alcohol subset, we are capturing both edge cases: explicitly weak intermolecular interactions and explicitly strong intermolecular interactions. This idea facilitates an insight into the relationships between activity coefficients
of solutes in solvent (IL) and vaporisation enthalpy
(298.15 K) of the solvent (IL). It is the basic knowledge, that activity coefficient
is the measure for strength of interactions between solute and solvent. Additionally, the vaporization enthalpy is the measure of intermolecular forces in pure solvent. Capturing the
values only for the edge cases (
n-alkanes + alcohols), we establish an unbiased average level of intermolecular energetics, which is fixed to the well-defined, restricted set of solutes. Provided that this set will be common for any ionic liquid, the regression of
values (see
Figure 2Right) becomes an unbiased thermodynamic tool for recovering vaporisation enthalpy
(298.15 K). Such a strict limitation only to the set (
n-alkanes + alcohols), allows crucial simplification of the
values evaluation, as the regression
δ1(298.15 K)–
Y is represented now by practically impeccable straight line (see
Figure 2Right). Merging the
n-alkanes and alcohols subsets is felicitous from the mathematical point of view, as the best possible scales of
δ1(298.15 K) values (from 15.2 to 29.4 MPa
0.5) and of
Y values (from 0.0583 to 0.3450 unities) are now encompassed for the
values evaluation (see
Figure 2Right). What is more, the conscious acquisition of only
n-alkanes and alcohol subsets releases oneself from the troublesome analysis of
δ1(298.15 K)–
Y regressions (see
Figure 1,
Figure 2 and
Figure 3) and the desperate retrieval of outliers justification. For the sake of brevity, the procedure of
values evaluation with help of merged alkanes and alcohols subsets along this paper is designate as “heptane–methanol” or
HM approach (as both solutes are most frequently present in the published data sets and they configure the frame of data acquisition). In our opinion the
HM approach could be considered as the virtually unbiased tool for examination of
and
(298.15 K) interrelations.
6. Examination of the HM Approach
For a preliminary validation of the
HM approach it is senseless to examine all 184 available papers [
29] on
values. Imidazolium based ionic liquids have been the most intensively studied in the current literature, and it is reasonable to test
HM approach with data on the [C
nmim][Anion] series. Among them, the most popular series of ILs was that associated with the NTf
2-anion. Enthalpies of vaporisation
(298.15 K) for this series of ionic liquids have been published recently [
10]. They were used for the reconciliation with results derived from
values reported in the literature [
44,
45,
46,
47,
48,
49,
50,
51,
52,
53,
54,
55]. Most of the data were measured by the GLC technique. Some data sets were also measured by the dilutor technique, but results from both methods have been shown to be indistinguishable. For the sake of transparency, we have preferred the comprehensive studies of activity coefficients, and we have omitted some papers dealing only with separating industrially relevant binary systems (e.g., hexane/hexane, etc.). The
values available in the literature at temperatures other than
T = 298.15 K have been adjusted to this reference temperature by using the linear extrapolation ln(
) =
f (1/
T). We fitted Equation (5) with the solubility parameters
δ2 derived from primary
values with help of Equations (2)–(4) for the [C
nmim][Anion] available in the literature, and estimated the
(298.15 K) values according to Equation (6). Typical results of the data treatment for the [C
nmim][NTf
2] series are given in
Table 1.
Experimental vaporization enthalpies are listed in column 2. In the first iteration, we processed all
values presented in the particular paper. The single data point for water was removed from regression, but all other outliers remained at this step. Vaporisation enthalpies calculated in this way are given in column 3. Without going too deeply into the details and peculiarities of each data set, as well as in reason of agreement or disagreement for each individual IL from
Table 1, it is quite apparent that the blind application of the traditional
δ1(298.15 K)–
Y regression is guesswork. Of course, the quality of each individual experimental data point is definitely responsible for the final quality of correlation. However, explicitly for the [C
6mim][NTf
2], the special IUPAC project [
44] was designed for testing different methods of
values. According to the final IUPAC conclusions for the majority of measurements where different techniques were used, the agreement is generally within the expected uncertainties for the measurement methods [
44]. However, this optimistic conclusion is not able to explain the dramatic scattering of
(298.15 K) values from 131.4 to 154.0 kJ·mol
−1 calculated for the [C
6mim][NTf
2] (see
Table 1, column 3). It is worth mentioning that four contributors [
45,
46,
47,
48] of
values listed for this IL in
Table 1 were participants of this IUPAC project. Additionally, what about the
HM approach? Results derived from this approach are given in
Table 1, column 4. Let us examine the
HM approach applied to [C
6mim][NTf
2] data sets, where the
-values are seemingly of certified and impeccable quality. The most comprehensive
-data set was reported by Heintz et al. [
45] and it comprised twelve alcohols (from methanol to
n-hexanol, iso-propyl-, iso-butyl-, sec-butyl-, tert-butyl-, and tert-pentyl-alcohols) and six
n-alkanes (from
n-heptane to
n-dodecane).
The value
(298.15 K) = 138.5 kJ·mol
−1 for [C
6mim][NTf
2] derived from the
data set by Heintz et al. [
45] is in very good agreement with the experimental value 139.9 ± 1.8 (see
Table 1, column 2). The less extensive collection of
values was reported by Kato and Gmehling [
46]. It included only methanol, ethanol,
n-propanol, iso-propanol, and only two
n-alkanes (
n-heptane and
n-octane). The value
(298.15 K) = 140.3 kJ·mol
−1 for [C
6mim][NTf
2] derived from their
data set is also in perfect agreement with the experimental value (see
Table 1, column 2). The modest
data set was reported by Letcher et al. [
47] and it consisted from methanol,
n-heptane, and
n-octane. The value
(298.15 K) = 144.4 kJ·mol
−1 for [C
6mim][NTf
2] derived from this very limited
data set is in acceptable agreement with the experimental value (see
Table 1, column 2). The
data set reported by Dobryakov et al. [
48] consists only of eight linear and branched alcohols, thus the data treatment according to the
HP approach is not possible. However, we could combine the alcohols set from Dobryakov et al. [
48] with the reliable
n-alkanes set from Heintz et al. [
45] and the joined treatment of these
data provides the value
(298.15 K) = 141.1 kJ·mol
−1 for [C
6mim][NTf
2], which agrees well with the experiment. Summing up the experiences with the
HM approach application towards [C
6mim][NTf
2] data, it is obvious, that the “traditional” treatment (see
Table 1, column 3) of the whole set of
values measured for the non-polar and polar solvent is not able to reproduce the reliable vaporization enthalpies measured by conventional methods (see
Table 1, column 2). In contrast, the application of the
HM approach allows estimated values that can be compared well with the experiment. This conclusion is supported by the comparison of the “theoretical”
(298.15 K) values derived from the
HM approach for the [C
nmim][NTf
2] series with
n = 2, 4, 6, 8, 10, and 12 (see
Table 1), as well as for the [C
nmim][CF
3CO
2] series with
n = 2 and 6 (see
Table 2).
Such a good correspondence between “theoretical” and experimental
(298.15 K) values has motivated further systematic comparisons. The results for [C
nmim][Anion] series with [CF
3SO
3], [CH
3SO
3], and [Cl] anions are collected in
Table 3.
To our surprise, column 2 and column 4 of
Table 3 reveal that the “theoretical” values are systematically underestimated in comparison to the “experimental” values. In an attempt to be strict and stringent, we have determined a factor F
im =
(exp)/
(
HM) (see
Table 1,
Table 2,
Table 3,
Table 4,
Table 5 and
Table 6 column 5) in order to quantify the degree of underestimation. We found that for the IL series collected in
Table 3, the factor F
im is close to unity (1.13 to 1.16) for each type of the IL. Additionally, it is important that F
im is hardly dependent on the chain-length within each series.
For the development of the
HP approach it is essential to trace whether the factor F
im is accidental, or if it is possibly anion-specific for each [C
nmim][Anion] series. In
Table 4, we evaluated results of regressions for the [C
nmim][Anion] series with fluorine-containing anions [BF
4], [PF
6], and [FAP].
As can be seen from
Table 4, the factor F
im = 1.28 ± 0.13 for [C
nmim][BF
4] and F
im = 1.41 ± 0.03 for [C
nmim][PF
6] are closely within their combined uncertainties. Additionally, the chain-length dependence of the F
im factors is absent in both series. Moreover, the increase in the number of F atoms in the anion from 4 in [BF
4] to 6 in [PF
6] does not seem to increase the F
im factor. In order to reinforce this conclusion, we extended our study with the perfluorinated anion [FAP]. Astonishingly, the factor F
im = 0.94 ± 0.02 (see
Table 4 for [C
nmim][FAP]) does not fulfil our expectations. Hence, it is clear that the factor F
im is virtually anion-dependent. However, F
im does seem to be chain-length independent (it is explicitly verified in
Table 4, column 5). These two conclusions have been additionally proven with the available
data sets for [C
nmim][NO
3] (see F
im= 1.44 ± 0.06 in
Table 5), as well as with the data sets for cyano-containing ionic liquids [C
nmim][Anion] with [SCN], [DCA], [TCM], and [TCB] anions (see
Table 6). As shown in this table, the increase in the number of CN-groups from one in thiocyanate, to two in di-cyano-amide, and three in tri-cyano-methane, does not seem to increase the F
im factor (F
im ≈ 1.3 to 1.4 within the combined experimental uncertainties). However, an accumulation of four CN-groups in the [TCB] reduces the factor significantly to the value of F
im = 1.19 ± 0.08 (see
Table 6 for [C
nmim][TCB]). Such a trend is similar to the accumulation of F atoms in the [C
nmim] [FAP] series (see
Table 4) and should be investigated more closely.
As a final conclusion, it has now been proven that the
HM approach can generally be used to estimate of
(298.15 K) of [C
nmim][Anion] by using the anion specific correction factors F
im to reconcile “theoretical” and “experimental” values. The variation of the F
im factors in the narrow range from 1 to 1.4 localized in this work enables a quick assessment of the general
(298.15 K) level, even with the averaged factor F
im = 1.2. However, these rough estimates could be aggravated by the uncertainties of 5–6 kJ·mol
−1, provided that good quality
data were used in the evaluation. Nevertheless, such agreement between “conventional”
(298.15 K) values and those “theoretical” from the
HM approach can be considered as acceptable, taking into account the experimental uncertainties at the level 2 to 3 kJ·mol
−1. The application of the anion-specific F
im factors (see
Table 1,
Table 2,
Table 3,
Table 4,
Table 5 and
Table 6) can improve agreement between “theoretical” and “experimental” vaporization enthalpies to 2–4 kJ·mol
−1. Taking into account the difficulties described in the introduction when measuring the enthalpies of vaporization of the extremely low-volatile ionic liquids, the
HM approach is considered a valuable indirect complementary method. It opens a wide window of opportunity to collect a large amount of reliable data for studies on structure–property relationships for ionic liquids. Moreover, the
(298.15 K) values derived from the
HM approach could help to resolve contradictions with the adjustments of vaporization enthalpies from experimental elevated temperatures to the reference temperature
T = 298.15 K.
During this study, we have frequently noticed that methanol on the regression plot [δ
1(298.15 K)—Y] was slightly above the general trend (e.g.,
Figure 1Right and
Figure 2Right). At the same time, the experimental point of ethanol met the line better. We use this observation in order to modify the (Heptane-Methanol) approach to the (Heptane-Ethanol) approach (or
HE approach). To do this, we excluded methanol from the primary
data sets and recalculated the anion-specific F
im factors (see
Tables S2–S8). The
HE approach with the specific F
im factors can be helpful in the evaluation of the primary
data sets, for which methanol was not included or the activity coefficient of methanol is of questionable quality. The reasonable combination of the
HM and
HE approaches increases the unbiased flexibility of the
data sets evaluation and estimation of reliable
(298.15 K) values.