Ionic liquids (ILs) are a fairly recent class of compounds that can be described as “salts with unusually low melting point temperatures”. A general rule-of-thumb used within the scientific community in this area of research defines an IL as any substance that has a melting point temperature lower than 100 °C (373 K) and is composed exclusively by ions or (in the case of protic ILs) ions in equilibrium with a very small proportion of their conjugated acid/base neutral species.
In terms of solution chemistry, ILs can be quite relevant simply because their liquid temperature range overlaps with that of common molecular species (MSs): unlike traditional molten salts, ILs can be mixed with other liquids like water or organic liquids, solubilize molecular gases like CO2or CFCs, and dissolve solid compounds like polymers or even (to a certain extent) traditional salts. The mixtures/solutions thus formed have unique properties conferred by the complex nature of the IL solvent.
ILs are composed by at least one molecular ion (generally the cation) characterized by charge delocalization between several of its atoms, asymmetrical shape, conformational flexibility, and (optionally) the presence of nonpolar moieties in the ion (generally dubbed alkyl side chains). The most commonly used cations are undoubtedly those belonging to the 1-alkyl-3-methyimidazolium family, [CnC1im]+,  but other cations such as tetra-alkyl ammonium  (including N,N-dialkylpyrrolidinium or cholinium derivatives), tetra-alkyl phosphonium, or N-alkylpyridinium are also commonly used. The range of possible anions is broader ranging from atomic anions such as the halides (from chloride to iodide) to fairly large molecular anions. One of the most interesting and popular choices has been the bis(trifluoromethylsulfonyl)imide anion, [Ntf2]–, that generally yields ILs with high thermal stability (above 250 °C) and relatively low viscosity (as compared with other IL analogues based on other anions).
Unlike conventional liquids, ILs are nanosegregated fluids that exhibit particular mesoscopic heterogeneities at a molecular level. This state of affairs should not be surprising since, after all, a substance that is composed solely of anions and cations must possess some kind of short-range organization in order to fulfill local electro-neutrality conditions that also maximize electrostatical interactions between ions of opposite sign (Fig. 1). In the case of ILs, the surprise comes from the fact that such imposed ordering does not lead to long-range ordered structures (crystals) at room temperature, a fact that can be associated with the characteristics of its ions (cf. above). For ILs with alkyl side chains, nanosegregation between polar and nonpolar domains can be regarded as the next logical step in this state of affairs: because the IL must necessarily order its high-charge density portions into local structures that obey electroneutrality and Coulomb-interaction maximization criteria, it follows that the low-charge density parts of the ions (the alkyl side chains) must be segregated elsewhere. It is this interplay between the two types of regions that eventually leads to the formation of medium-range structures and the recognition of ILs as a high-charge density network permeated by low-charge density regions (Fig. 2).
Molecular dynamics (MD) were carried out using the DLPOLY package . An all-atom force field, which is based on the OPLS-AA framework  but was to a large extent developed specifically for ILs. An OPLS-AA parametrization was employed for the solutes. When studying microstructured fluids, the size of the systems and the duration of the simulations are of particular importance: periodic boundary conditions can induce artificial finite size effects on the length scales of the observed nanostructures. Due to the slow dynamics of this type of system, special care was taken to ensure the attainment of equilibrium conditions, including the proper diffusion of the solutes in the IL media: (i) equilibrations started from low-density initial configurations; (ii) typical equilibrations were implemented for more than 1 ns at constant NpT; (iii) multiple re-equilibrations through the use of temperature annealing and switching off and on of the Coulomb interactions were performed; (iv) further simulation runs were used to produce equilibrated systems at the studied temperatures. Electrostatic interactions were treated using the Ewald summation method considering six reciprocal-space vectors, and repulsive-dispersive interactions were explicitly cut off at 1.6 nm (long-range corrections were applied assuming the system had a uniform density beyond this cut-off radius).
Ab initio calculations
Electrostatic charge distributions were calculated using Gaussian 03  at the MP2/cc-pVTZ(-f)//HF/6-31G(d) level of theory, thus using the same basis set as in the OPLS-AA model for ILs or for solutes. The cc-pVTZ(-f) basis set was used for single-point energy calculations in geometries optimized at the HF/6-31G(d) level, as is current practice in the development of force field parameters for molecular simulation [6, 7]. For the C and N atoms, the cc-pVTZ-(-f) basis set is created by removing the f functions from the definition of the triple-ccpVTZ basis set of Dunning . The combination of the levels of theory and basis sets used here has been tested on a large collection of molecules (Halgren test) and was demonstrated to yield accurate conformational energetics . It must be stressed that no constraints were placed on any of the molecules during geometry optimization or point charge calculation in order to reproduce a specific molecular moment.
The fluid phase behavior and solvation properties of liquid mixtures and solutions are largely influenced by the underlying intermolecular interactions between the system components. These interactions determine not only the macroscopic properties of the system but also define its local structure at the molecular level.
Pure ILs can be understood as mixtures at fixed composition (generally equimolar) of two ionic species. Moreover, sometimes the ions exhibit high-charge density regions (the so-called polar “heads”) and nonpolar moieties (the alkyl side chains, or tails). Before discussing the interactions and structure of solutions of ILs plus molecular species, let’s first consider this particular kind of pseudo-binary mixtures: pure ILs.
The different morphologies of the mesoscopic structure of pure ILs
The polar network – the flexible structure formed by the high-charge density parts of the ions that compose the system – is one of the defining characteristics of an IL or molten salt (cf. Fig. 1).
ILs formed by ions where the charge is delocalized among most of the atoms within the molecular ions and that do not have long alkyl side chains (e.g., [C1C1im][Ntf2]) exhibit just a large array of (apparently) disordered and diffusing ions that nevertheless manage to preserve local electroneutrality conditions (each ion is surrounded by a contact shell of counterions) and are also able to establish an intermediate-order network of ions with alternating charges (a succession of shells with opposing charge, cf. Fig. 1). This last feature – a periodic occurrence in direct space with a characteristic wavelength as given by the vertical gridlines in Fig. 1 – can be highlighted in reciprocal space by calculating the corresponding total structure factor functions, S(q) (Fourier transforms of the total correlation functions). These are shown in Fig. 3, where the intermediate peak at around q = 9 nm–1 corresponds to a wavelength of about 0.7 nm, i.e., the spacing of the vertical gridlines of Fig. 1. On the other hand, the peak at q = 13 nm–1 reflects the total correlation between contact ion pairs of opposing sign. This means that even an IL whose ions do not possess any nonpolar moieties will exhibit medium-range structural features that reflect the ionic nature of the fluid (e.g., the two peaks in Fig. 3 at 9 and 13 nm–1).
ILs with long alkyl side chains will exhibit an additional feature, the so-called pre-peak, in the corresponding S(q) functions. In the case of the S(q) functions of the three [CnC1im][Ntf2] ILs represented in Fig. 3, the pre-peak is absent for [C3C1im][Ntf2], starts to appear for [C6C1im][Ntf2], and is very intense for [C9C1im][Ntf2], in total agreement with X-ray diffraction data . These pre-peaks correspond to a further structuration of the ionic fluid, i.e., the segregation of the alkyl chains into progressively larger nonpolar domains and the eventual formation of a second continuous subphase (in addition to the already existent and continuous polar network). This phenomenon – a kind of percolation limit of the bicontinuous subphases – starts to occur after the length of the chains reach an IL-specific threshold value (around [C5C1im][Ntf2] in the case of the [CnC1im][Ntf2] IL family). Before that threshold is reached, the small nonpolar domains are just dispersed in the continuous polar network and no pre-peak is formed.
In other words, the relative size (volume occupied) by the polar and nonpolar moieties of the ions that compose an IL will determine its morphology (dispersed vs. bicontinuous).
Traditionally, most structural studies [11–13] have been based on 1-alkyl-3-methylimidazolium cations. However, many other ILs are based on other cations with quite different polar/nonpolar ratios and different distributions of the polar and nonpolar parts inside the ion: in 1-alkyl-3-methylimidazolium cations, the long alkyl-side chains end up in a polar “head”; in tetra-alkyl ammonium or phosphonium cations that polar “head” is completely surrounded by chains of different size, occupying in this way a more “central” position in the ion. Figure 4 illustrates such differences.
In the case of the tetra-alkyl ammonium IL, it is not very easy to make out the contours of the polar network since its thread-like appearance is finely enmeshed within the nonpolar regions. The 1-tetradecyl-3-methylimidazolium IL exhibits a quite different situation with ‘‘globules” of polar and nonpolar regions quite conspicuous at the surface of the simulation box. The different connectivity of the polar network – sequential in the case of the thread-like structures of the ammonium and phosphonium cations, highly branched in the case of the globular structures of the imidazolium cations – reveals how the interactions between the high-charge density parts of the anions and cations affect the morphologies of their polar networks. Apart from the obvious (and rather isotropic) electrostatic interactions, anions and cations can also interact in some cases through other types of quite specific (and strong) interactions. A case in point is the hydrogen-bond-type interactions that the imidazolium cation can establish with different anions through its C2 (and in a lesser degree C4 and C5) hydrogen atoms. These interactions allow each imidazolium to interact with more than two anions at the same time and thus form “branched” polar networks that can develop into globular structures. On the other hand, the “central” tetra-alkyl phosphonium or ammonium cations do not establish such interactions and the hindered “central” atoms of phosphorus or nitrogen will be able to interact simultaneously with a limited number of counter-ions. The result will be nonbranched sequences of ions and the development of filamentous networks.
Finally, the alkyl side chains can be also functionalized – PEG-like chains or chains with a terminal hydroxyl group are just two examples [10, 14]. The functionalization of the chains will also influence the morphology of the corresponding ILs, either by the disappearance of the pre-peaks for long PEG-like alkyl side chains (instead of forming segregated domains, the PEG-like chains will tend to sheath the polar network)  or, in the case of the hydroxyl-functionalized chains, by producing an additional strong interaction center that is no longer centered in the charged part of the ion [14, 15].
Molecular solutes dissolved into the mesoscopic structure of ILs
The concept that ILs are highly structured fluids at a mesoscopic level changes the way one can interpret the solvation properties of molecular species in those media. Different molecular solutes, according to their polarity or tendency to form associative interactions, will interact selectively with different parts of the individual ions and may also be solvated in distinct (nanosegregated) local environments. In this context, ILs can be regarded as ‘‘two-in-one” solvents, with each domain acting as a specific solvent for a given class of solutes: molecules like n-alkanes will tend to dissolve in the nonpolar domains, whereas dipolar or associating solutes will interact with the polar network, Fig. 5 [16, 17].
A general conclusion is that solutes can be classified into three groups: those that prefer the nonpolar regions of the IL – typically nonpolar solutes like alkanes ; those that interact strongly with the polar network – water is the obvious example of such an associative fluid that can promote hydrogen bonding, specially with specific atoms of the anion; and finally dipolar solutes that can orient themselves at the interface between the polar and nonpolar regions of the IL–acetonitrile, acetone, small halogenated hydrocarbons fall into this category (Fig. 5) . It is in this last group (dipolar/quadrupolar non-associative fluids) that one can find good solutes/solvents for a large variety of ILs, a behavior that can be rationalized in terms of the positioning of the molecules at the interface between the polar and nonpolar regions and their ability to orient their dipoles in order to interact simultaneously with the anions and cations of the polar network. In some cases, the permeation of the polar network can lead to its disruption and eventual complete solvation of the ions that compose the IL – there is complete miscibility between many classical molecular solvents (acetonitrile, acetone) and a large proportion of ILs.
The 2-in-1 solvent: fluorinated gases into phosphonium-based ILs
Sometimes it is surprisingly difficult to rationalize the solubility of two related solutes in the same IL . The reason is that the two solutes are experiencing different local environments of the IL, in other words, it is like the two solutes are being dissolved in two distinct solvents. In this context, the interactions between trihexyl(tetradecyl)phophonium bis(trifluoromethylsulfonyl)amide, [P6 6 6 14][Ntf2], and totally fluorinated alkane gases (perfluoro-methane, -ethane, and -propane) have been investigated using gas solubility measurements and MD simulations.
Figure 4a vividly shows the nanosegregated nature of [P6 6 6 14][PF6]. The [P6 6 6 14][Ntf2] solvent under discussion exhibits a very similar filamentous polar network embedded in the extensive nonpolar domain. If this IL feature is taken into account, it can be assumed that different solutes can be solvated in different regions of the IL and so be exposed to completely different molecular environments. In the present case, it was possible to show that the C3F8 molecules will be restricted to the nonpolar, aliphatic regions of the IL, whereas the CF4 molecules will also be found nearer the polar, high-charged density regions (Fig. 6). This means that the enthalpy and entropy contribution to the solvation energies of the two molecules will be extremely difficult to compare – as if the two solutes were being dissolved in two different solvents – and explains the different slopes of the ΔsolvH(T) and ΔsolvS(T) functions obtained experimentally for the dissolution of CF4 and C3F8 in [P6 6 6 14][Ntf2].
First hints on the charge-template topic: the unusual phase behavior of aromatic solutes
Aromatic compounds are much more soluble in ILs than their aliphatic counterparts. Basically, one would expect that an hydrocarbon as benzene would dissolve preferentially in the nonpolar domains of the IL, and in fact the solubility of benzene in [C10C1im][Ntf2] (larger, continuous nonpolar domains) is much greater than in [C2C1im][Ntf2] (smaller, dispersed nonpolar domains). Nevertheless, at 298 K it is possible to dissolve almost four molecules of benzene per [C2C1im][Ntf2] ion pair before liquid–liquid demixing occurs. Interestingly, it is almost impossible to dissolve any IL in pure benzene.
The unexpected fluid phase behavior of aromatic compounds in ILs has been rationalized taking into account not only the liquid–liquid and solid–liquid phase diagrams of benzene plus [C2C1im][Ntf2] , but also the solubility data of all 12 fluorinated benzene derivatives . It was shown that the distinct dipole and quadrupole moments of benzene and its derivatives can be correlated to their solubility in [C2C1im][Ntf2] [22, 23].
Those that have large dipole moments (when the di-, tri-, or tetra-fluorination is on the same side of the molecule) are completely miscible with [C2C1im][Ntf2]. On the other hand, the five most symmetrical elements of this family (that have null dipole moment but exhibit non-null quadrupole moments) are only partially miscible with [C2C1im][Ntf2] (however, all are much more soluble similar non-aromatic hydrocarbons). The interesting thing is that the relative solubility of all 13 solutes in [C2C1im][Ntf2] can be correlated, taking into account the way their dipole and quadrupole moments affect (or are affected by) the polar network of the IL (Fig. 7). Albeit empirical, the correlation was built taking into account molecular insights gained from ab initio calculations of the isolated aromatic solute molecules and MD simulations of all 13 aromatic solutes plus [C2C1im][Ntf2] mixtures. This type of molecular-assisted rationalization unveiled a simple correlation between the dipole and quadrupole moments of the solutes and the IL solvent. It also revealed the complex nature of the interactions between aromatic compounds and ILs, with the charge density functions of the former acting as a sort of molecular charge template that promotes the partial re-organization of the ions of the latter and defines the fluid phase behavior (liquid–liquid demixing) of the corresponding binary mixtures.
The charge template concept can be illustrated if we take into account Fig. 7. The top row shows the electrostatic potential functions (EPFs) around benzene, hexafluorobenzene, and 1,2,3-trifluorobenzene. Due to their symmetry, the first two molecules have null dipole moment and quadrupole moments (of opposing sign) in the direction normal to the aromatic plane; the latter molecule has a strong dipole moment superimposed on its quadrupole moment. Regions of high electron density are shown in red/orange tones; those with electron deficiency are shown in blue/cyan/green colors. The bottom row shows spatial distribution functions (SDFs) of the ions of [C2C1im][Ntf2] (the cation in red, the cation in blue) around the same aromatic solutes. One can easily see that benzene permeates the polar network in a way that the cations will preferentially occupy the polar positions and anions the equatorial regions around the aromatic ring; the situation is reversed for hexafluorobenzene and is shifted away from the aromatic plane for the case of 1,2,3-trifluorobenzene. Moreover, the reds and the blues are reversed both in the benzene/hexafluorbenzene case but also if one compares the EPFs and SDFs of the same molecule: the SDFs are the “negative” of the EPFs. Such complementarity is the so-called charge-template effect: the ions of the IL organize themselves around the aromatic molecules (or the latter permeate the polar network) in a way that Coulomb interactions are maximized – blue areas in the top row of Fig. 7 are matched by red areas in the bottom row and vice versa.
Subtle templates and exhaustive diagrams: haloalkane solutes
The charge-template idea was then extended to other solutes, namely, those where multiple dipole moments can be combined to yield different fluid phase behavior outcomes.
An important example is a systematic study performed with mono- and di-haloalkanes (bromo- or chloro-) in ethylsulfate- or ethylsulfonate-based ILs, [C2C1im][C2SO4] and [C2C1im][C2SO3], respectively.  The study considered systems with all 1-haloalkanes from 1-halopropane to 1-haloheptane (5 species) and with all 1,n-dihalopropanes (2 isomers) or 1,n-dihalobutanes (3 isomers). This means that the phase diagrams corresponding to the liquid–liquid equilibria (LLE) of 40 [(5 + 3 + 2) × 2 × 2] different systems were experimentally determined and the underlying fluid phase behavior rationalized using MD simulation results.
The correlation between fluid-phase behavior, the electrostatic potential functions of the solutes (determined by the existence of strong dipole moments), and the charge-templating effect of the IL ions enabled us to established a few rules that were able to explain the overall phase behavior of all 40 systems:
The difference between the sulfonate- and sulfate-based IL solvents can be, obviously, attributed to the presence of the extra oxygen bridge-atom in the sulfate anion (Fig. 8). This atom, placed between the charged SO3 group and the nonpolar alkyl side chain, causes shifts in the charge distribution of the anion especially in the first carbon atom of the alkyl side chain (cf. table 1 of  and Fig. 8). Despite their structural similarities, the ethylsulfate and ethylsulfonate anions are in fact quite different, with the former exhibiting a larger polar moiety and the possibility of forming more extensive polar networks when combined with different cations. This means that dipolar aprotic solutes like the haloalkanes studied in this work will have more difficulty in dissolving in the tightly bound [C2C1im][C2SO4] than in the less bound [C2C1im][C2SO3].
Both ILs have very short alkyl side chains (the ends of the two ethyl groups of the anion and the cation) attached to their polar head groups (the sulfate and sulfonate groups of the anions and the imidazolium ring and adjacent atoms of the cation), that is, their nonpolar regions are very small. This means that haloalkanes with progressively larger alkyl groups will have more difficulty in being solvated by the (almost exclusively polar) IL.
The effects of chlorine vs. bromine substitution are subtler than what one could anticipate based on just the size and electronegativity differences between the chlorine and bromine atoms. The substituted end of the haloalkane (where the dipolar charge separation is concentrated, Fig. 9) is responsible for most of the interactions with the high-charge density parts of the ILs ions, with the rest of the haloalkane (the alkyl side chain) acting as a residue that has to be “tolerated” by the IL solvent (which in this case does not have any nanosegregated nonpolar regions in its midst). This explains the already discussed lower solubility of the longer 1-haloalkanes and also the fact that those (low) solubilities are almost the same irrespective of the type of halogen substitution, although the 1-bromoalkanes tend to be slightly more soluble.
Only in the case of 1-bromopropane and 1-chloropropane there are significant differences in the solubility of these solutes in the two studied ILs with the former 1-haloalkane exhibiting larger solubility values, especially in the [C2C1im][C2SO3] solvent. Given the similarity of the dipole moment of the two molecules, it is hard to interpret such large differences. However, it is known that solutes that are able to maintain simultaneously strong interactions with both ions of an IL exhibit enhanced solubility. This is the case of small strongly dipolar molecules like acetone or acetonitrile, but also strongly quadrupolar molecules such as carbon dioxide, benzene, or perfluorinated benzene [16, 20, 22]. In the present case, 1-bromopropane and 1-chloropropane do not have to accommodate a long alkyl side chain. This means that, released from that extra burden, the solute capable of interacting simultaneously with the cation (interactions between the halogen and the hydrogen atoms of the imidazolium ring) and the anion (interactions between the substituted carbon atom of the haloalkane and the oxygen atoms of the anions) exhibits a larger solubility in the IL. Apparently, 1-bromopropane accomplishes that task more efficiently than 1-chloropropane (especially in the sulfonate-based IL that has the ions not so tightly connected in its polar network), and, therefore, allows for the simultaneous interaction of the solute with parts of that network. In this context, the larger polarizability of the bromine atom certainly plays a role in promoting more efficient, flexible, and dual interactions with both ions of the IL.
In the case of dihaloalkanes, the solubility of the dichloroalkanes is generally better than their dibromo- counterparts. In summary, one can say that 2 C-Cl dipoles are more versatile than 2 C-Br dipoles,  that the internal rotation barriers that are larger in the dibromoalkanes play a role in the existence of large upper critical solution temperatures (UCSTs) in the phase diagrams of the corresponding (IL + dibromoalkane) mixtures, and that those hindered rotations (that define the magnitude of the interacting dipoles and the ability of the solutes to permeate/template the IL polar networks) are a monotonous function of the position of the double halogenation (1,n-dibromoalkanes are more hindered in the order 1.2 > 1.3 > 1.4).
Functionalized ILs: reversed UCST trends
The next issue to be addressed is the possibility of enhancing the complexity of the IL–MS interactions by functionalizing the alkyl chain of the former component. As we have mentioned already, a PEG-like alkyl side chain leads to a much more subdued segregation of the nonpolar domains and the suppression of the corresponding pre-peaks. In the present paper we will focus the remainder of the discussion on the introduction of a terminal hydroxyl group at the end of a short chain, as in the case of cholinium, a functionalized tetra-alkylammonium cation.
The LLE of mixtures of cholinum-based ILs (N-alkyl-N,N-dimethylhydroxyethylammonium bis(trifluoromethane)sulfonylimide, [N1 1n 2OH][Ntf2] (n = 1, 2, 3, 4, and 5), plus water or 1-octanol exhibit phase diagrams dominated by the existence of UCSTs. The solubility of [N1 1n 2OH][Ntf2] in water is lower for cations with longer alkyl side chains (larger n values), whereas the corresponding trend in the octanol mixtures is reversed. Furthermore, the ([N1 1 1 2OH][Ntf2] + water + octanol) ternary system shows triple liquid−liquid immiscibility at room temperature and atmospheric pressure.
In this context, [N1 1n 2OH][Ntf2] ILs are quite different from nonfunctionalized ones. For instance, whereas aqueous mixtures of [CnC1im][Ntf2] are generally dominated by LLE conditions with no measurable UCSTs, aqueous mixtures of [N1 1 1 2OH][Ntf2] and [N1 1 2 2OH][Ntf2] exhibit UCSTs at 345 and 367 K, respectively. Moreover, the imidazolium-based systems also exhibit pronounced asymmetries in terms of the mutual solubilities of the two compounds in the aqueous mixtures, a point that is not true for the cholinium-based systems. The difference between the [N1 1 1 2OH][Ntf2] and [CnC1im][Ntf2] aqueous systems can be inferred from the RDFs obtained from simulations of a few water molecules dissolved in the choline-based ILs (Fig. 10). Previous works [26, 27] have shown that the solubility of water in ILs is controlled by the different affinity of the water molecule toward the IL ions: the tendency to bind selectively to just one type of ion can have a detrimental effect. This is the case of [CnC1im][Ntf2] ILs where water strongly competes with the cation for strong interactions with the [Ntf2]− anion. In the case of [N1 1 1 2OH][Ntf2] the cation has an extra interaction center that can promote its binding to the water molecules, and this means that the interaction of the water molecules with the two ions is more balanced, allowing for a better dissolution of water molecules among the ionic network of the IL (IL-rich mixtures) and the solvation of ionic aggregates in the midst of water (water-rich mixtures).
Other factors besides the hydrogen-bond-type interactions between the hydroxyl group of the cholinium-based cations and the water molecules determine the position of the UCSTs and their dependence with the alkyl chain length of the cations. In fact, if hydrogen bonding between water and the OH group of the cholinium-based cations were dominant at around room temperature, it would become less so at higher temperatures, which would mean that as temperature went up the systems would tend to remain demixed (an entropically driven effect that generally promotes LCST and not UCSTs). The demixing at lower temperatures in the aqueous [N1 1 1 2OH][Ntf2] is therefore defined by the hydrophobic effect caused by the presence of larger or smaller nonpolar domains in the IL. This effect also explains the reverse UCST trend in the ([N1 1n 2OH][Ntf2] plus 1-octanol) mixtures: the radial distribution functions between the terminal carbon atoms, CTO and CT, of the octyl chain of the 1-octanol molecule and the alkyl chain of the cation, respectively, show that the octyl residue of the 1-octanol molecule is accommodated inside a alkane-like domains whenever they are large enough . This means that the solubility of 1-octanol in [N1 1 5 2OH][Ntf2] will be enhanced relative to that of [N1 1 1 2OH][Ntf2] and explains the new UCST trend. It must also be noted that other relevant RDFs also show  that the interaction between the hydroxyl group of the octanol molecule and the polar network remains active even when the rest of the octyl chain is plunged inside the nonpolar domains of the ILs.
More functionalized ILs: unusual LCSTs
Systems of cholinium-based ILs [N1 1n 2OH][Ntf2] (n = 2, 5, 8), with ether molecules can exhibit unusual liquid–liquid demixing behavior characterized by the existence of lower critical solution temperatures (LCSTs). The mutual solubilities of the ILs and ethers are larger (higher LCST values) for ILs with longer alkyl side chains attached to the cations. On the other hand, the solubility of the ethers is defined by the exposure of the corresponding C–O–C dipole and the extension/compactness of the alkyl residues surrounding the ether function. Auxiliary ab initio calculations and MD simulations were able to rationalize again the experimental results at a molecular level in terms of the charge-template concept . In summary, four main points defining the solubility in these systems can be identified:
The demixing-upon-heating phenomenon is related to the breaking up of the hydrogen-bonded network that can form between the functionalized cation of the IL and the oxygen atom of the ether molecule (intense first peaks in Fig. 11).
It is generally easy to interpret from a molecular point of view the increase in solubility of nonpolar molecules in ILs as the alkyl moieties attached to the ions of the latter get larger. Larger domains will allow a better solubilization of the hydrocarbon chains of the different ethers and justify the solubility order [N1 1 8 2OH]+ > [N1 1 5 2OH]+ > [N1 1 2 2OH]+.
If one assumes that the solubility is also conditioned by the formation of specific interactions between the hydroxyl group of the alkyl cholinium cations and the oxygen atom of the ethers, then the strength of such interactions will depend mainly on the proton acceptor ability and accessibility of the oxygen atom in the ether. Figure 11 also shows, for three selected ethers, the corresponding EPFs mapped onto an electron density isosurface and also the atomic point charges assigned to the oxygen atoms present in each molecule. From the point of view of the negative charge induced in the oxygen atom by its two substituents, one can see that for aliphatic noncyclic moieties the order tert-C > methylene > methyl is followed. This means that ethers with a methyl group should exhibit weaker hydrogen bonds and be less soluble in the cholinium-based ILs (lower LCST values). This is clearly not the case because there is an opposing trend defined by the accessibility of the oxygen atom to perform a stereospecific interaction such as a hydrogen bond. In this particular case, a methyl substituent confers to the ether a much less encumbered oxygen atom (compare, for instance, the relative sizes of the negative areas of the tert-butyl methyl ether and tert-butyl ethyl ether molecules (Fig. 11). The figure also highlights the same fact by showing the RDFs between those two ethers and [N1 1 8 2OH][Ntf2]. The simulation results show that the two opposing trends (more negative charge on the oxygen atoms induced by more substituted aliphatic carbon atoms vs. more accessible oxygen atoms due to smaller methyl groups) can be reconciled and optimized if one of the substituents in the oxygen atom is a methyl group and the other a tertiary carbon like in the case of a tert-butyl radical. It also explains why diethyl ether has a comparatively poor solubility when compared with that of tert-butyl methyl ether in spite of the fact that the calculated charge in the oxygen atoms is almost the same in both cases.
When comparing alkyl methyl ethers (where the accessibility of the oxygen atom is assured by the presence of the common methyl group), the inductive effect of the second alkyl chain can explain the order of solubility in the series. However, in this case the IL itself can also play an important role. As discussed above, the nonpolar domains of the IL are responsible for dissolving the alkyl moieties of the ether molecules. Ethers with more compact substituents (more ramified or cyclic alkyl groups) can be more readily accommodated in such nonpolar domains, especially when the latter tend to be larger and bulkier as in [N1 1 8 2OH][Ntf2].
2D or not 2D: from solution to surface templates
Finally, the charge-template concept can be extended from solute/solvent interactions in a bulk (3D) solution to an IL phase adsorbed at the (2D) surface of a solid substrate. In this contextn MD simulations of a 5-nm-thick layer of the functionalized IL 1-(2-hydroxyethyl)-3-methylimidazolium tetrafluoroborate, [(OH)C2C1im][BF4], over silica, alumina, and boro-silicate glass substrates have been performed. The structure of the IL at the solid/liquid interface has been interpreted. taking into account the corresponding normal density profiles, lateral interfacial structure, orientational ordering, and planar density contours . Comparisons with experimental data suggest that the adsorption and stratification processes that occur when ILs are placed in contact with solid substrates [31, 32] can be correctly modeled using a realistic rendition of a nonuniform amorphous substrate such as a glass material.
The MD results show that an IL that interacts strongly with a given amorphous substrate will orient its ions in order to mirror the (uneven) local electrical field felt by it at the different regions of the surface (Fig. 12). In other words, a charge-template phenomenon is the driving force for stratification: the strong adsorption of the IL ions on a flat surface that is not necessarily uniform at a molecular level is efficiently accomplished because the IL ions can adapt to the (uneven) electrical field produced by the solid substrate. The use of the solid substrate as a charge template by the IL ions reflects again the versatility of ILs as adsorbates of any solvation media – as we have seen in the previous sections, in a bulk phase, the ions of ILs are known to interact selectively with different molecular solutes, adopting different spatial arrangements around them and promoting unique fluid-phase demixing phenomena.
The unique, complex, and versatile nature of ILs can be correctly modeled and discussed at a molecular level using ab initio and MD simulation data.
From the interactions point of view, ILs display an intricate balance between electrostatic, dispersion, and H-bond interactions that can cause nanosegregation at an atomistic level and the formation of different mesoscopic subphases. The length and functionalization of the alkyl side chains exhibited by many ILs play a particularly important role in the definition of different morphologies.
When (IL + molecular species) solutions are considered, the structural richness of ILs and the versatility of its ions in terms of possible interactions cause an enormous variety of possible fluid-phase behaviors: ILs act as charge templates for different solute molecules by allowing their solvation in different domains, aligning their ions according to the molecule’s dipoles or quadrupoles, or establishing additional H-bonds. The concept of IL ions as effective charge-templates for different molecular species can also be extended to the adsorbtion of IL films at substrate surfaces.
A collection of invited papers based on presentations at the 33rd International Conference on Solution Chemistry (ICSC-33), Kyoto, Japan, 7–12 July 2013.
Financial support provided by Fundação para a Ciência e Tecnologia (FCT) through projects FCT-ANR/CTM-NAN/0135/2012 (including a post-doctoral grant awarded to KS), PTDC/CTM-NAN/121274/2010 and PEst-OE/QUI/UI0100/2013.
K. Shimizu, M. F. C. Gomes, A. A. H. Pádua, L. P. N. Rebelo, J. N. C. Lopes. J. Mol. Struct.: THEOCHEM946, 70 (2010).Google Scholar
K. Shimizu, A. A. H. Pádua, J. N. C. Lopes. J. Phys. Chem. B114, 15635 (2010).Google Scholar
W. Smith, T. R. Forester. The DL_POLY Package of Molecular Simulation Routines (v.2.2), The Council for The Central Laboratory of Research Councils, Warrington, Daresbury Laboratory (2006).Google Scholar
M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria, M. A. Robb, J. R. Cheeseman, J. A. Montgomery Jr., T. Vreven, K. N. Kudin, J. C. Burant, J. M. Millam, S. S. Iyengar, J. Tomasi, V. Barone, B. Mennucci, M. Cossi, G. Scalmani, N. Rega, G. A. Petersson, H. Nakatsuji, M. Hada, M. Ehara, K. Toyota, R. Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda, O. Kitao, H. Nakai, M. Klene, X. Li, J. E. Knox, H. P. Hratchian, J. B. Cross, V. Bakken, C. Adamo, J. Jaramillo, R. Gomperts, R. E. Stratmann, O. Yazyev, A. J. Austin, R. Cammi, C. Pomelli, J. W. Ochterski, P. Y. Ayala, K. Morokuma, G. A. Voth, P. Salvador, J. J. Dannenberg, V. G. Zakrzewski, S. Dapprich, A. D. Daniels, M. C. Strain, O. Farkas, D. K. Malick, A. D. Rabuck, K. Raghavachari, J. B. Foresman, J. V. Ortiz, Q. Cui, A. G. Baboul, S. Clifford, J. Cioslowski, B. B. Stefanov, G. Liu, A. Liashenko, P. Piskorz, I. Komaromi, R. L. Martin, D. J. Fox, T. Keith, M. A. Al-Laham, C. Y. Peng, A. Nanayakkara, M. Challacombe, P. M. W. Gill, B. Johnson, W. Chen, M. W. Wong, C. Gonzalez, J. A. Pople. Gaussian 03, revision C.05; Gaussian, Inc.: Wallingford, CT, 2004.Google Scholar
J. N. C. Lopes, A. A. H. Pádua, K. Shimizu. J. Phys. Chem. B112, 5039 (2008).Google Scholar
K. Shimizu, D. Almantariotis, M. F. C. Gomes, A. A. H. Pádua, J. N. C. Lopes. J. Phys. Chem. B114, 3592 (2010).Google Scholar
T. H. Dunning Jr. J. Chem. Phys.90, 1007 (1989).Google Scholar
R. A. Friesner, R. B. Murphy, M. D. Beachy, M. N. Ringnalda, W. T. Pollard, B. D. Drunietz, Y. Cao. J. Phys. Chem. A103, 1913 (1999).Google Scholar
A. A. H. Pádua, J. N. C. Lopes. J. Phys. Chem. B110, 3330 (2006).Google Scholar
Y. Wang, G. A. Voth. J. Am. Chem. Soc.127, 12192 (2005).Google Scholar
S. M. Urahata, M. C. C. Ribeiro. J. Chem. Phys.120, 1855 (2004).Google Scholar
A. J. L. Costa, M. R. C. Soromenho, K. Shimizu, I. M. Marrucho, J. M. S. S. Esperança, J. N. C. Lopes, L. P. N. Rebelo. ChemPhysChem13, 1902 (2012).Google Scholar
J. Restolho, J. L. Mata, K. Shimizu, J. N. C. Lopes, B. Saramago. J. Phys. Chem. C115, 16116 (2011).Google Scholar
J. N. C. Lopes, M. F. C. Gomes, A. A. H. Pádua. J. Phys. Chem. B110, 16816 (2006).Google Scholar
L. Pison, J. N. C. Lopes, L. P. N. Rebelo, A. A. H. Pádua, M. F. C. Gomes. J. Phys. Chem. B112, 12394 (2008).Google Scholar
M. T. Zafarani-Moattar, H. Shekaari. J. Chem. Eng. Data50, 1694 (2005).Google Scholar
J. Łachwa, J. Szydłowski, A. Makowska, K. R. Seddon, J. M. S. S. Esperança, H. J. R. Guedesa, L. P. N. Rebelo. Green Chem.8, 262 (2006).Google Scholar
M. B. Shiflett, A. Yokozeki. J. Chem. Eng. Data53, 2683 (2008).Google Scholar
K. Shimizu, M. F. C. Gomes, A. A. H. Pádua, L. P. N. Rebelo, J. N. C. Lopes. J. Phys. Chem. B113, 9894 (2009).Google Scholar
F. J. Deive, A. Rodríguez, A. B. Pereiro, K. Shimizu, P. A. S. Forte, C. C. Romão, J. N. C. Lopes, J. M. S. S. Esperança, L. P. N. Rebelo. J. Phys. Chem. B114, 7329 (2010).Google Scholar
M. Aroney, D. Izsak, J. W. Lefevre. J. Chem. Soc. 1407 (1962).Google Scholar
P. Nockemann, K. Binnemans, B. Thijs, T. N. Parac-Vogt, K. Merz, A. V. Mudring, P. C. Menon, R. N. Rajesh, G. Cordoyiannis, J. Thoen, J. Leys and C. J. Glorieux, J. Phys. Chem. B113, 1429 (2009).Google Scholar
U. Domańska, M. Laskowska. J. Solution Chem.37, 1271 (2008).Google Scholar
A. J. L. Costa, M. R. C. Soromenho, K. Shimizu, I. M. Marrucho, J. M. S. S. Esperança, J. N. C. Lopes, L. P. N. Rebelo. J. Phys. Chem. B116, 9186 (2012).Google Scholar
A. J. L. Costa, M. R. C. Soromenho, K. Shimizu, J. M. S. S. Esperança, J. N. C. Lopes, L. P. N. Rebelo. RSC Adv.3, 10262 (2013).Google Scholar
K. Shimizu, A. Pensado, P. Malfreyt, A. A. H. Pádua, J. N. C. Lopes. Faraday Discuss.154, 155 (2012).Google Scholar
R. Hayes, S. Z. El Abedin, R. Atkin. J. Phys. Chem. B113, 7049 (2009).Google Scholar
R. Atkin, G. G. Warr. J. Phys. Chem. C111, 5162 (2007).Google Scholar