Structure, Interaction and Hydrogen Bond

Abstract

Ionic liquids (ILs), as green solvents, have attracted amazing interest and their potential applications have prompted a large amount of research and investment, and some of the results have been inspiring. In recent years, in combination with cations and anions, some new ILs have been synthesized in the laboratory. However, compared with simple solid salts, the structures of ILs are complicated and their properties vary considerably. It is thus very time consuming to explore ILs experimentally when facing the huge number of ionic combinations. A molecular-based understanding can reveal the quantitative correlation between structures and properties, and is thus an important subject in the study of ILs. The unusual complexity of ionic interactions renders molecular-based interpretations difficult and gives rise to controversies about the structure of the ILs. Herein we discuss the ion-pair, cluster and X-ray crystals structures and their relationship with the properties of many typical ILs, especially imidazolium-based. In the ILs, apart from the strong electrostatic forces, non-covalent H-bonds and van der Waals (dispersion, induce forces) are examined and are shown to have a decisive effect on the properties of ILs.

Beijing Key Laboratory of Ionic Liquids Clean Process, Key Laboratory of Green Process and Engineering, State Key Laboratory of Multiphase Complex System, Institute of Process Engineering, Chinese Academy of Sciences, Beijing 100190, PR China.

1 Short Review of Ionic Liquids (ILs)

1.1 Synthesized ILs

Room temperature ionic liquids (RTILs) have been paid close attention by academia and the industries due to their environment-friendly properties, and particularly their potential in green chemistry. It is nearly 100 years since an IL, [EtNH₃][NO₃], which has a melting point of 12°C, was first reported in 1914 [1]. Then research on RTILs renewed with the discovery of alkylpyridinium or 1,3-dialkylimidazolium haloaluminate salts by groups at the Air Force Academy in the mid-1970s. These salts were formed by mixing aluminum halides with the corresponding imidazolium or pyridinium halides [2]. Unfortunately, these haloaluminate ILs are highly sensitive to atmospheric moisture. In the 1990s, these salts with imidazolium-based cations incorporated with tetrafluoroborate or hexafluorophosphate anions were stable in the atmosphere and were regarded as a significant milestone in the development of ILs [3, 4]. By careful choice of the combination of these cations and anions it was possible to synthesize a large variety of ILs. Since 2000, more than 6,000 papers related to ILs have been published, with more than 1,000 in 2004, 1,300 in 2005, and nearly 1,900 in 2006 [5].

Non-haloaluminate RTILs share a common general structure of a bulky unsymmetrical “onium” organic cation associated with a weakly coordinating inorganic anion. So far, the most popular cations are 1,3-dialkylimidazolium-, N,N-dialkylpyrrolidinium-, N,N,N,N-tetraalkylammonium-, and N-alkylpyridinium-based. Concerning their anionic counterparts, tetrafluoroborate ([BF₄]⁻), hexafluorophosphate ([PF₆]⁻), bis(trifluoromethylsulfonyl)amide ([NTf₂]⁻), and triflate ([OTf]⁻) are the most common. Figure 1 shows some cations and anions commonly used for the preparation of RTILs [6, 7].

Fig. 1

Cations and anions studied: (a) cations; (b) anions

RTILs have been attracting interest because of the many unique properties they are supposed to have. RTILs have negligible vapor pressure and are thus nonflammable and nonvolatile, largely decreasing the chance for fugitive emissions. They show high thermal stability, with decomposition temperatures around 300–500°C. They are able to solvate a large variety of organic and inorganic compounds, either polar or nonpolar. They are generally regarded as polar yet noncoordinating solvents. They display a good intrinsic conductivity and are extremely redox-robust. Actually, as a result of the intense research dealing with their physicochemical properties, it has become clear that the above-mentioned properties cannot be considered as generic properties for these media. For instance, it has been shown that many ILs decompose above 80–90°C.

1.2 Relationship between Properties and Structures of ILs

Melting Points (T _m ). Many properties of ILs are special and show prospects for applications. ILs are completely composed of ions, but the unusually low melting points exceed expectations. Melting points of many ILs are lower than 0°C; even the melting point of [Emim][N(SO₂CF₃)₂] is −16°C [8]. The low melting points are related to the structures and molecular packing of ions. An increase in size, anisotropy, and internal flexibility of the ions should lower the melting point, as shown in Table 1. The symmetry and length of alkyl chain of cations also influence the melting point, and the short N-alkyl chains and higher molecular symmetry can frequently increase the melting point. The melting point of [Mmim]Cl IL is 124.5°C and of [Bmim]Cl is 65°C. A simple rule is that cations with short N-alkyl chains (C _n  ≤ 3) form crystalline phases with high melting points, and ILs with 4 ~ 10 N-alkyl chains exhibit a broad liquid range with low melting points and a remarkable tendency to supercool. When the n in C _n is larger than 10, the salts show a complex phase behavior [10].

Table 1 Properties of the typical 1-ethyl-3-methylimidazolium-based ILs

Viscosity (η). RTILs are viscous liquids, their viscosities being one to three orders of magnitude higher than those of conventional solvents [11]; obviously that is not very helpful for some applications and it may exert a strong effect on the rate of mass transport within solution and on the conductivity of the salts. The viscosity of the RTILs strongly depends on the nature of the anion [12–14]. ILs containing [NTF₂]⁻ anion present the lower viscosity, for instance, the viscosity of [Emim][NTf₂] IL is 25 mPa·S as shown in Table 1, but viscosity is higher for RTILs containing [PF₆]⁻ nonplanar symmetric anions, and the viscosity of [Bmim][PF₆] IL is 308 mPa·S [15]. It has been suggested that the size, shape, and molar mass of the anion contribute to the viscosity with smaller, lighter, and more symmetric anions leading to more viscous RTILs [16]. Additionally, the relative basicity of the anions and their ability to form H-bonds or to allow van der Waals attractions also have a pronounced effect on viscosity. The fluorinated anions such as [BF₄]⁻, [PF₆]⁻, or [BETI]⁻ form ILs with higher viscosity due to the formation of H-bonds or stronger van der Waals forces, but salts with [N(CN)₂]⁻ and [C(CN)₃]⁻ anions exhibit low viscosity. Compared to anions, cations have secondary effects on the viscosities of ILs. For any of the cation types, increasing the length of alkyl chains results in higher viscosities because of stronger van der Waals interactions between larger cations [17, 18].

Temperature has a remarkable effect on the viscosity of ILs. Okoturo et al. found that the viscosity decreases when the temperature increases. The correlation can be described by the Vogel-Tammann-Fulchers equation:

$$ \ln \eta = \ln {\eta_{\infty }}+\frac{{{{\mathrm{ E}}_{\eta }}}}{\mathrm{ RT}} $$

(1)

where ${\eta_{\infty }}$ is the viscosity of infinite high temperature ${{\mathrm{ E}}_{\eta }}$ is the energetic barrier of ionic motion.

Conductivity (Λ). Conductivity is a property of primary importance and most electrochemical processes are related to the conductivity of the ILs. Being composed of ions, ILs are supposed to be among the most concentrated electrolytes with many charge carriers per unit volume. Some ILs exhibit higher conductivity, and the order of 10 mS·cm⁻¹ can be found in the imidazolium family (Table 1). However, quaternary ammonium ILs exhibit lower conductivity, and the highest conductivity is 2 mS·cm⁻¹, found in N, N-dialkylpyrrolidinium bis(trifluoromethylsulfonyl)amide.

As mentioned previously, the high viscosity of RTILs has a major impact on the conductivity because the conductivity is inversely linked to the viscosity [16, 18, 19]. The less viscous [NTf₂]⁻ salts usually exhibit among the highest conductivities [20]. However, quantitative correlations have not been observed between viscosity and conductivity in most ILs. For instance, [Emim][TfO] and [BMIm][NTf₂] display similar viscosities and densities, but the conductivity of [Emim][TfO] is 0.29 mS·cm⁻¹ and [Bmim][NTf₂] is 0.91 mS·cm⁻¹, and their conductivities differ by a factor of 2. Many others factors also influence the conductivity, such as ion size and aggregation, anionic charge delocalization, and correlated ionic motions [13, 16, 21]. The Nernst-Einstein equation relates the molar conductivity to the self-diffusion coefficients of ions. However, the values predicted by

$$ \Lambda =\frac{{{{\mathrm{ F}}^2}}}{\mathrm{ RT}}({{\mathrm{ D}}_{\mathrm{ cation}}}+{{\mathrm{ D}}_{\mathrm{ anion}}}) $$

(2)

are not in good agreement with experimental measurements, and experimental values are smaller than the theoretical values, which can be rationalized by a coupled and clustered aggregation motion of cations and anions in electrically neutral configuration, but the neutral species cannot contribute to the ionic conductivity. On the other hand, increasing the length of the N-alkyl chains of cation results in a higher viscosity and a lower conductivity.

Enthalpy of Vaporization (ΔH). An important property of ILs is nonvolatility such that the ILs are green and able to substitute for environmentally contaminating volatile organic chemicals (VOCs). However, Earle et al. [22] found that many RTILs can be distilled and not decomposed at higher temperature and lower pressure. Owing to the strong ionic interactions, the molar enthalpies of vaporization of the ILs are by almost an order of magnitude higher than for typical molecular liquids, as shown in Table 1. The anionic symmetry has the obvious effect on the vaporization of IL, and the ILs with more symmetric anions, Cl⁻, [BF₄]⁻, and [PF₆]⁻, have higher enthalpy of vaporization than [NTf₂]⁻ salt. However, these ILs can decompose in the gas phase, for example, [Emim]Cl thermally decomposes when heated to 190°C in vacuum. The volatile decomposition products are 1-methylimidazole, 1-ethylimidazole, chloromethane, chloroethane, ethane, and hydrogen chloride species, but the [NTf₂]⁻ anionic salts are thermally stable and may be heated to 600 K. Thus most vapor measurements have been carried out on ILs by different methods (isothermogravimetry, Knudsen-type effusion, surface tension, etc.). In such studies it was found that increasing the length of alkyl chains of cation leads to a larger enthalpy of vaporization, the enthalpy values being 135, 155, 173, and 192 kJ·mol⁻¹ for [C₂mim][NTf₂], [C₄mim][NTf₂], [C₆mim][NTf₂], and [C₈mim][NTf₂], which indicates the long alkyl chain leads to stronger interactions between cations and anions [23, 24].

1.3 Hydrophilic and Hydrophobic

The presence of water in ILs can greatly influence the quality of ILs, including viscosity, polarity, and conductivity, and can also influence the solubility of other substance in RTILs. For example, the presence of water in [Bmim][PF₆] IL has resulted in the underestimation of the solubility of CO₂ in the ionic liquid [25]. These imidazolium-based ILs present different hydrophilicities. Anions such as halide, pseudo-halide, [BF₄]⁻, methyl sulfate, [NO₃]⁻, and [ClO₄]⁻ are hydrophilic, which significantly limits their usages, while ionic liquids with anions such as [PF₆]⁻, [SbF₆]⁻, [OSO₂CF₃]⁻, [OCOCF₃]⁻, etc. are hydrophobic and have potential in some industrial applications. However, it is also clear that the commonly called “hydrophobic” ILs are actually hygroscopic and can absorb a certain amount of water from the air [17].

T. Welton et al. have investigated the state of water in many imidazolium-based ILs by ATR and transmission IR spectroscopy [26]. They found that anions are mainly responsible for the solubility of water. In these ILs, the water is not associated together but bound by H-bonds with anions with concentrations of dissolved water in the range 0.2 ± 0.1 mol · dm⁻¹. The water molecules at these concentrations exist in symmetric 1:2 type H-bonded complexes: anion…H–O–H…anion. Additional evidence can be seen from Table 2 that the shifts of antisymmetric ν ₃ and symmetric ν ₁ bands of water molecules in the listed ILs roughly follow the relationship of Δν ₃ = 1.17Δν ₁ − 0.60. The strength of H-bonds between the H₂O molecule and anions can be estimated from the wave number shift (Δν) between ν ₃ of water in vapor phase and that dissolved in ionic liquids by Eq. 3 [27]:

$$ \Delta H=-80({\nu_3}^{vapor }-{\nu_3}^{IL })/{\nu_3}^{vapor } $$

(3)

(where ΔH is the enthalpy of H-bond, KJ/mol). It can be seen that the different anion can interact with water by the different strength of H-bonds as listed in Table 2. In addition, the water contents absorbed from the open air have a consistent trend with the strength of the H-bonds; for example, the contents of water in [Bmim][PF₆], [Bmim][BF₄], and [Bmim][(CF₃SO₃)₂N] increase by the order of 2,640, 19,500, and 33,090 ppm by Karl-Fischer at relative humidity 59 % [26]. It implies that the H-bond play an important role for the solubility of water in ILs.

Table 3 Crystallographic data of 1-ethyl-3-methylimidazoium halides ILs [32, 33]

Table 2 The shifts of the ν3 and ν1 (cm⁻¹) of water in different ionic liquids (with [Bmim]⁺ cation) compared with the corresponding water bands in vapor and the enthalpies of H-bonding [26]

Although the cation plays a secondary role for the solubility of water in bulk ILs, at the vapor-liquid surface the water mainly interacts with the imidazolium cations. The hydrophobic ILs display a dramatic reorientation at the gas-liquid interface as water contacts the surface. It had been shown by sum frequency generation (SFG) of the C–H stretching mode that C–H groups, especially the C₂–H group in ring contact with water by C–H…O H-bond interaction, produces a tilting of the cation on the surface. Where the plane of the ring will be, either parallel or perpendicular to the alkyl chain, is dependent on the protonation or methylation of the C₂ position [28, 29].

Generally speaking, the properties presented are important for the application of ILs, especially basic scientific data; however some data are difficult to measure directly. However, the properties of ILs strongly depend on the structures and anion–cation combination styles, and understanding of the structures and the ionic interaction will provide a more convenient way to estimate thermodynamic data prior to practical process design.

2 Structures of ILs

In recent years, some new ILs were synthesized in the laboratory, and the number of possible cation and anion combinations has increased significantly such that researchers believe the one trillion RTIL could possibly be prepared [30]. The synthetic problem of being able to design RTILs rationally through variation of the anion and cation still remains to be thoroughly investigated, despite attempts to correlate structures with properties. It is thus essential to develop a systematic method of selectively choosing a given ion pair to be used as a predictive tool in the rational design of new ILs.

A structurally based understanding is a great challenge because the molecular and electronic structure, charge distribution, and orbital overlap of ions, give rise to complex interactions and forces involved in the electrostatic, hydrogen bonds, and van de Waals interactions. Currently there has been substantial growth in the number of theoretical investigations pertaining to ILs, whereby scientists are attempting to predict many of the physical properties and have produced results agreeing with experimental values. Many quantitative correlations have also been established to help predict physic-chemical properties and accelerate the exploration for new ILs [31].

2.1 Dialkylimidazolium-Based Halides ILs

Unlike ionic solid materials, for example NaCl, in which the molecules and ions can be packed close to each other, leading to high lattice energy, the halide anionic ILs are liquid state at ambient temperature, although the strong electrostatic interaction – the bulkiness of both the cation and anion – prevents such packing, thereby lowering the lattice energy.

Halide anionic ILs have already been studied and the anions were found mainly to involve chloride, bromide, and iodide. However, a fatal defect of the class of ILs is their moisture problems in the atmosphere; thus their applications are very limited. In recent years, ILs (1-ethyl-3-methylimidazolium halides ([Emim]X, X = C1 or Br)) have been mixed with the corresponding aluminum(III) halide (AlX₃, X = C1 or Br) to produce the halogenoaluminate ILs, in which the anions exist as [AlX₄]⁻ and [Al₂X₇]⁻. Halogenoaluminate(III)-based ionic liquids, which may be used to investigate a wide range of industrially important chemistry, have been of particular interest [32].

The ILs are solid at ambient temperature and their crystallographic data are collected in Table 3. The [Emim]Cl IL is an orthorhombic crystal system, but similar solid structures were found in [Emim]Br and [Emim]I. Starting with the coordinates from the iodide structure, refinement by full-matrix least squares, with non-hydrogen anisotropic atoms and hydrogen isotropic atoms, was used to solve the structure.

Figure 2 shows the packing of ions and local structures in the crystals of [Emim]Br and [Emim]I. The structures consist of different layers of anions and cations which are interconnected by the extensive H-bonds (Fig. 2a, b). Each cation is hydrogen bonded to three anions and each anion is hydrogen bonded to three cations (Fig. 2c). The local structure around the cation is superficially similar to that observed for [Emim]Cl, but the overall morphology of the crystals is somewhat different.

Fig. 2

The packing of ions in solid [Emim]Br (a) and [Emim]I (b), and the local ionic structure (c) [32]. The dashed lines showed the H-bonded interactions

Figure 3 shows the pairwise molecular structures of 1-ethyl-3-methylimidazolium and 1-butyl-3-methylimidazolium halides salts from ab initio calculations with the lowest interactive energies [34]. [Emim]Cl has four different ionic structures (Fig. 3), and the chloride ion with coplanar structures is more stable. Chloride ion, found in the plane of the ring, is associated with the hydrogen of C2 of the ring, positioned closer to the methyl substituent as opposed to the ethyl substituent, which is agreeable with crystal structure. [Emim]Br has three different ionic structures (Fig. 3). The first (Emim Br 1 in Fig. 3) was found to be the most stable structure, and the third (Emim Br 3 in Fig. 3) coalesced with the first at the MP2 levels. Unlike the chloride system, where the chloride was found to be coplanar with the ring, the bromide of the most stable structure was found to be above the ring plane, which was a significant difference with respect to the position of the bromide in the crystal structure and the crystal structure indicates that the bromide resides in the ring plane (Fig. 3). The [Emim]I have two different structures, wherein the anion is within the plane of the ring. The first is the most stable structure (Emim I 1 in Fig. 3), agreeing with the crystal structure. The H-bond distance in the calculated structure (0.285 nm) was found to agree with the literature (0.293 nm).

Fig. 3

1-Ethyl-3-methylimidazolium (a) and 1-butyl-3-methylimidazolium (b) halide ion pair structures with lowest interactive energies by ab initio calculations [34]

Gross trends relating interaction energies and melting points were found within the chloride, bromide, and iodide series (see Fig. 4). From the figure it can be seen that the iodide series with the different chain lengths exhibit a “linear” trend, in which the melting points were found to decrease with increasing chain length, and this was associated with a decrease in the magnitude of interaction energy. In the cases of the chloride and bromide anions, the melting points of the [Pmim]⁺ analogue were significantly lower than those of the [Emim]⁺ and [Bmim]⁺ analogues, although the interaction was found to increase in magnitude with increasing alkyl chain length. For the bromides, the structures were different, which introduces some scatter in the results. A trend was also found to exist in the [Bmim]⁺ halide series, where the melting point increase between [Bmim]Cl and [Bmim]Br was associated with an increase in magnitude of the interaction. The melting point was then found to decrease significantly between [Bmim]Br and [Bmim]I, and this was associated with a significant decrease in magnitude of the interaction.

Fig. 4

Correlation between melting points and the binding energies for halide-based ionic liquids: [Emim]⁺ (filled diamonds); [Pmim]⁺ (filled squares); [Bmim]⁺ (filled triangles) [34]

The preliminary investigation of a series of three imidazolium-based cations paired with three halide anions provides some initial hints as to the relationship between the interaction energies and the melting points. The result suggests that there is more than one factor contributing to the melting point behavior of ILs. It is possible that for certain systems the melting point is governed more strongly by the cation than the anion and vice versa, with a balance between Coulombic attractions of oppositely charged ions and Van der Waals repulsions of the alkyl chains on the imidazolium cation.

2.2 Fluoro-Anion-Based ILs

The main defect of these halide ILs is that they are very sensitive to the water in air, which prohibits development of further applications. The fluoro-anion ILs (e.g., anion = [BF₄]⁻, [PF₆]⁻, [(FH)_nF]⁻, [AsF₆]⁻, [OCOCF₃]⁻, [OSO₂CF₃]⁻, etc.) are typically air-stable and have a greater potential in many chemical processes.

2.2.1 X-ray Diffraction Single-Crystal Structures

• 1. ILs with Cations of Different Alkyl Lengths

ILs have shifted the conventional view on ionic molten salts because of their liquid state at room temperature. Coulombic attractive forces in ILs must be diffuse to inhibit or suppress crystallization. Varying the length or degree of branching of the alkyl chains of imidazolium cations can dramatically change their melting points. This has attributed to variations in the rotational freedom of the alkyl chains of the cations and the ability of the cations to pack efficiently (or not) within the crystalline cell. Branching in alkyl substituents can similarly lead to increases in melting point through restriction of rotational freedom, but can also, conversely, reduce melting points when branching results in enantiomeric mixtures.

Reichert et al. have examined the structure of the solid state imidazonium hexafluorophosphate ([Dmim][PF₆]) ILs with different N-alkyl chains by X-ray crystallographic analyses [35]. Figure 5 presents the local structures and ion packing in the crystal structures, in which [PF₆]⁻ anions contact with imidazolium cation via the H-bonds; however, the length of the N-alkyl chains have a remarkable effect on the packing of the ions. When the N-alkyl chains are short, C _n  ≤ 4, the [PF₆]⁻ anions contact not only with the heads (ring) of the cation, but also with the tails (alkyl chains), as shown in Fig. 4a–i. As shown in Fig. 4j–l, [PF₆]⁻ anions make contact with only the heads of the cation when the chains are longer (C _n  > 4). From the side views of the cations presented in Fig. 5, it is found that there were several locations of the [PF₆]⁻ anions above and below the imidazolium heads and in close contact with the C and N atoms of the rings. The direct interactions between the delocalized positive charge of the imidazolium ring and the anion is further indication of the Coulombic nature of ILs based on these ions. Although the electrostatic forces are dominant between cations and anions, the close contacts with a wide angular distribution indicates the C−H…F H-bonded interaction. However, as Reichert et al. said, the H-bonded interactions were weak and more Coulombic in nature. Thus, the interactions that are present in the crystal structures of the short chained salts suggest that the forces responsible for packing in a crystalline lattice are predominantly electrostatic or Coulombic in origin. This also suggests that lattice energies can be described by Eq. 4:

Fig. 5

Local structures and ions packing diagrams of (a) [C₁mim][PF₆], (b)[C₂mim][PF₆], (c) [C₂dmim][PF₆], (d)[C₂tmim][PF₆], (e) [C₄mim][PF₆], (f) [secC₄mim][PF₆], (g) [tertC₄mim][PF₆], (h) [C₄dmim][PF₆], (i) [(isoC₃)₂mim][PF₆], (j) [C₁₀mim][PF₆], (k) [C₁₂mim][PF₆], and (l) [C₁₄mim][PF₆]. The dashed lines show the close contacts and H-bonded interactions [35]

$$ {{\mathrm{ U}}_{\mathrm{ L}}}=2\mathrm{ I}\bigg[\frac{\upalpha}{{{^3}\sqrt{\mathrm{ V}}}}+\upbeta \bigg] $$

(4)

I is the ionic strength (=1), V is the molecular volume (in nm³) of the lattice, which is equal to the sum of the individualcation (V+) and anion (V−) volumes, and α and β are empirically derived parameters.

• 2. ILs with different anions

The ILs containing fluoroanions are air-stable, with a consequent greater potential application in chemical processes. Except for the common ILs with [BF₄]⁻ or [PF₆]⁻, some other ILs containing fluoroanions have been reviewed in recent papers by Matsumoto and Xue et al. [36, 37]. The melting points of these fluoroanionic ILs are around freezing point, and thus the solid state structures are important to the understanding of the ILs’ behavior. Table 4 lists the crystal structural parameters for the fluoro-containing ILs with typical 1-ethyl-3-methylimidazolium ([Emim]⁺) cation.

Table 4 Cell parameters and melting points of 1-ethyl-3-methylimidazolium ILs with fluoroanions [36]

In the crystal, the [Emim][FHF] is layered and it contains the smallest fluorocomplex anion [FHF]⁻ in the series. The flat imidazolium rings are stacked and arranged in parallel with the interlayer distance of 3.376 Å. In each layer, very short hydrogen bond distances between the H atom of the cation ring and the F atoms of the anion (1.951, 2.166, and 2.226 Å ) are observed [38]. In the [Emim][BF₄] IL, [Emim]⁺ cations form a pillar with the β-carbon of the ethyl group sticking out of the imidazolium-ring plane [39, 40]. The H(methylene)-π interaction sustains the cation–cation interactions with a distance of about 2.86 Å between the H atom in methylene group and the imidazolium ring centroid. [Emim][AsF₆] and [Emim][SbF₆] are isostructural with [Emim][PF₆], whereas [Emim][NbF₆] exhibit a different structure [3, 41]. Figure 6 shows the ion packings in [Emim][AsF₆] and [Emim][NbF₆]. In the [Emim][AsF₆] structure, [Emim]⁺ cations and [AsF₆]⁻ anions stack alternately along the β-axis to form pillars. On the other hand, in the [Emim][NbF₆] structure, the anion appears in a zigzag arrangement along the α-axis where the nearest fluorine atoms have a distance of 3.441 Å and the cations adopt a pillar-like stacking along the same axis. H-bonds often play an important role in the melting points of ionic compounds. For salts with fluorocomplex anions, in spite of closer H…F contacts in [Emim][NbF₆] than those in [EMIm][PF₆], [EMIm][AsF₆], and [EMIm][SbF₆], the melting points of the former two compounds are much lower than those of the latter three compounds. [Emim][BF₄] also exhibits a low melting point despite relatively strong H-bonds in its lattice. These observations suggest that the strength of hydrogen-bonding is not always a decisive factor in their melting points.

Fig. 6

Ions packing in solid (a) [Emim][NbF₆] and (b) [Emim][AsF₆] ILs [42]

2.2.2 The Structures of Ion Pair and Ion Cluster

• 1. The Structure of the Ion Pair

It has been verified experimentally that the ions are pairwise in the gas phase [43, 44]. Although in liquid state, the ions are not strictly composed of ion pairs; cations and anions can contact closely due to strong electrostatic attraction. Theoretical calculations have been performed in an attempt to understand the structure of ion pairs and interaction between cation and anion [10, 45].

We have investigated in detail the ion-pair structures of the two typical ILs, [Emim][BF₄] and [Bmim][PF₆], by ab initio DFT calculations [46]. In the [Emim][BF₄] ion pair, [BF₄]⁻ anion can locate at five positions of cation as shown in Fig. 7. Then the five structures of ethyl-front, methyl-back, back, methyl-front and ethyl-back as initial configurations were optimized at B3LYP/6-31++G** theory level. Finally, the three lowest energy geometries were obtained. The [Emim][BF₄](I) was obtained from the ethyl-front, methyl-back, and back configurations, in which [BF₄]⁻ anion moved to the upper part of the ring from the lateral part and located near to the C₂–H group. The [Emim][BF₄](II) was obtained from the methyl-front configuration and [BF₄]⁻ anion also moved to the upper part of the ring. [Emim][BF₄](III) was obtained from the ethyl-back conformer, but [BF₄]⁻ anion did not move and only localized near to the C₅–H group.

Fig. 7

The optimized structures of [Emim][BF₄] ion-pair when anions are located at five different positions of cation. The arrows indicate the initial locations of the anion. The dashed lines indicate the H-bonds in these ion pairs

A shallow potential energy well was present when the [BF₄]⁻ anion localized at the back and the methyl-front of imidazolium ring and experienced the minimal electrostatic attraction to the significant stronger electrostatic attraction due to the essential neutral charge distribution on the C_4/5 atoms and the positive charge on the C₂ atom. However, when localizing at the ethyl back of imidazolium ring as shown in Fig. 7, it is difficult for the [BF₄]⁻ anion to move due to steric exclusion derived from the ethyl side chain and the electron repulsion derived from the π-bond between the C₄ and C₅ atoms.

To date, many researchers have discussed the possible interactions between cation and anion and proposed that H-bonds also widely exist in ILs and represent the most explicit interaction besides electrostatic attraction [47–51]. According to the criterion for forming an H-bond – that the distance between H on donor atom and basic acceptor atom, R_H…A, is less than the sum of their respective van der Waals radii – herein the van der Waals distance is 2.670 Å [36]. There are four F…H H-Bonds to form in [Emim][BF₄](I) as drawn by dashed lines in the right of Fig. 4. Two H-bonds form between the F atoms labeled F₂ and H(C₂) atoms of the ring, and H(C₆) atom of the alkyl chain. The distances are 2.02 and 2.54 Å, respectively. The other two occur between the F₃ atom and H(C₂) atom of the ring, and H(C₈) atom of the alkyl chain. The distances are 2.29 and 2.21 Å, respectively. A very similar H-bonded structure can be found in [Emim][BF₄](II), but there is no H-bond between F₃ atom and H(C₂) atom. When [BF₄]⁻ anion locates at the ethyl back of imidazolium ring as [Emim][BF₄](III), shown in Fig. 4, only two H-bonds form, and one is between the F₂ atom and H(C₂) atom of the ring, and another is between the F₃ atom and H(C₆) atom of the alkyl chain. Their distances are 1.91 and 2.05 Å, respectively.

For the [Bmim][PF₆] ion pair, the similar five initial configurations, butyl-front, methyl-back, back, methyl-front and butyl-back, were optimized at the B3LYP/6-31++G** level. The three lowest energy structures were obtained as shown in Fig. 8. [Bmim][PF₆](I) was obtained from the butyl-front and the methyl-front configurations. [Bmim][PF₆](II) and [Bmim][PF₆](III) were obtained from the back and the butyl-back conformers, respectively. However, no stable point was found at the potential energy surface for the methyl-back configuration. For the back configuration, the [PF₆]⁻ anion moves above the imidazolium ring and the steric exclusion of the long butyl chain is responsible for the location compared with the shorter ethyl chain on the [Emim]⁺ cation.

Fig. 8

The optimized structures of [Bmim][PF₆] ion pair with anions located at the five possible different positions around the cation. The arrows represent the initial locations. The dashed lines indicate the H-bonds in these ion pairs

There are five H-bonds in the [Bmim][PF₆](I) ion-pair. One is between the F atom labeled as F₂ of [PF₆]⁻ anion and the H(C₂) atom of imidazolium ring and the distance is 2.49 Å. Two are between the F₃ atoms and the H(C₂) atom, and the H(C₆) atom of alkyl chain. The distances are 2.12 and 2.52 Å, respectively. The other two are between the F₄ and H(C₂) atoms, and the H(C10) atom of the alkyl chain. The distances are 2.14 and 2.26 Å, respectively. When the [PF₆]⁻ anion locates at butyl-back, there are four H-bonds as [Bmim][PF₆](III) in Fig. 8. One H-bond can form between the F₂ atom and the H(C₅) of ring, and the distance of 1.97 Å is the shortest. The other three are between two F atoms and H(C₅) and H(C₆) of the alkyl chain, and the distances are 2.47, 2.30, and 2.32 Å, respectively. However, when the [PF₆]⁻ anion locates at the back of imidazolium ring, the [PF₆]⁻ anion moves above imidazolium ring to form the [Bmim][PF₆](II) and there is no H-bond to form.

• 2. Understanding the Interaction at Electronic Level

By NBO analysis it was found that in [Emim][BF₄] ion-pairs the occupation on the N₃ orbital always increases, while the occupations on the π _C2=N1 and π _C4=C5 orbitals vary with the positions of anion. In [Emim][BF₄](I) and [Emim][BF₄](II) ion pairs, the occupation on the π _C2=N1 orbital decreases, but on the π _C4=C5 the orbital increases. In [Emim][BF₄](III) ion-pair, the occupation of the π _C4=C5 orbital decreases, and the occupation on the π _C2=N1 orbital exhibits no significant change. Furthermore, the molecular orbitals provide an essential insight into the ion pairs. Figure 9 shows the orbitals and energies for both [Emim][BF₄](I) and [Emim][BF₄](III) ion-pairs. The HOMO of the isolated anion is degradation p orbitals (energy E = −0.256), and the LUMO of the cation is the anti-π (π*) bond orbital (energy E = −0.189). However, in the ion pairs there are not the expected p-π orbital interactions, but the HOMO of the [Emim][BF₄](I) is the highest occupied p orbital of anion and the lowest empty π*_(C2–H) orbital of cation, and the LUMO of [Emim][BF₄](I) exhibits an anti-σ (σ*) bond orbital (energy E = −0.0522). In the [Emim][BF₄](III) ion pair, the effective orbital overlap also occurs and the HOMO (energy E = −0.304) and the degradation HOMO-1 (energy E = −0.302) are the highest occupied p orbitals of anion and the lowest empty σ*_(C5–H) and σ*_(C6–H) orbitals of cation. The σ-type interaction rather than p-π overlap indicates a covalent bonding feature, but the formation requires a stronger intermolecular attraction. The electrostatic attraction drives anion to approach cation, and finally locates at the largest positive charge position of the cation. Although the repulsion of the electron between anion and cation increases the total energy, the orbital effective overlap and the partial electrons transfer result in an energetic reduction, and further stabilize the ion pairs.

Fig. 9

The orbitals and energies diagram for [Emim][BF₄] ion-pairs. (Due to the same orbital diagram with [Emim][BF₄](I) ion-pair, the [Emim][BF₄](II) ion-pair were not drawn)

Obviously from the two classes of ion-pairs with multiple atomic fluoro-anions, we find that more than one H-bonds can form between cation and anion, which is different from ionic liquids with a single atomic anion, like the dialkyimidazolium chloride ([Emim]Cl), and with only one H-bond between cation and anion [45–47].

• 3. The Structure of Ion Cluster

As reported [52–54], the ions clustering by ionic close contacts reflects the structural arrangement in the bulk phase of RTILs [55]. The clusters may become microscopic models to describe the structures of the bulk ionic liquids or their aqueous solutions. Figure 10 shows an experimental FTIR spectrum of [Emim][BF₄] IL and the calculated IR vibrational frequencies for the ([Emim][BF₄])_n ion clusters with different number, n = 2, 3, 4, 5 at B3LYP/6-31 + G** level. The strong agreement on major bands provided the possibility to explore the detailed structural information of [Emim][BF₄] IL, and these experimental vibrations can also be assigned rationally by the structures of these ion clusters.

Fig. 10

Measured FTIR spectrum of [Emim][BF₄] at room temperature conditions compared to the vibrational modes of the corresponding ion clusters ([Emim][BF₄])_n with n = 2, 3, 4, 5 calculated by DFT at B3LYP/6-31 + G** level (all of the calculated bands were corrected by the factor 0.964–0.967)

The band at 3,000–3,500 cm⁻¹ is assigned to the symmetric stretches of three C–H bonds of the imidazolium ring. Two weaker peaks around 1,600 and 1,500 cm⁻¹ are assigned to the breadths of the rings and the scissors of the alkyl chains, respectively. The peaks around 1,208 cm⁻¹ are assigned to the interplanar wagging vibrations of the C–H bonds on rings. The strongest peak occurs around 1,136 cm⁻¹ and is assigned to the stretches of B–F bonds of [BF₄]⁻ anion, while these weak peaks around 750–880 cm⁻¹ are the out-of-planar wagging vibrations of the rings.

Structural investigation into the clusters has found that the ion arrangements are ordered and similar to the crystal structure shown by X-ray [40, 42]. The structures of clusters are shown in Fig. 11. It was found that these ions are connected by the H-bonds (dashed lines) together to form a three-dimensional network [47]. In the network, the ion packing is influenced very little by the intramolecular bonding structures, and the imidazolium cations are arranged in alternative layers and the [BF₄]⁻ anions are sandwiched between two cations but do not locate above the rings and prefer to be near the C-H groups of rings in the clusters. The cation–cation distance is around 5.10 Å between H(C₂) atom and the imidazolium ring center. No strong interaction between the ring proton and π-electron cloud of the imidazolium ring is observed and the ethyl groups stick out of the imidazolium ring planes (C₂–N₁–C₆–C₇ torsion angle around 74°). [BF₄]⁻ anions form a pillar with a B…B distance of 5.7 Å in two different anions. Neutron diffraction studies on some RTILs indicated that charge ordering endured in the liquid phase.

Fig. 11

The ion clusters of ([Emim][BF₄])_n with n = 2, 3, 4, 5 by tube fashion. The dashed lines denote the forming H-bonds

Larger nano-clusters have been observed in imidazolium-based ILs with alkyl chain length more than 4. Based small-wide-angle X-ray (SWAXS), Hardacre et al. have found low-Q peaks in [C _n mim][PF₆] (n = 4, 6, 8) ILs,which are correlated with the heads (imidazolium rings) of cations and indicate the heads aggregate together [56, 57]. These observations have been supported by molecular dynamic simulation. Figure 12 shows the ion packings of [C _n mim][PF₆] (n = 2, 4, 6, 8, 10, 12) ILs observed by molecular dynamic simulation [58]. It can be seen that the ions do not aggregate when connected with short chains (n = 2). When the lengths of the alkyl chains are increased (n = 4, 6, 8), the ions start to aggregate, and the polar domains (red color, the charged anions and imidazolium ring parts) and the non-polar domains (green color, the non-charged alkyl chains) are obvious. When the lengths of the alkyl chains were increased to n = 12, the aggregation is very obvious, and the larger clusters form. In the clusters, the anions and imidazolium rings aggregate together to form the polar domains by charged attraction, and the non-charged long alkyl chains “twist” together to form non-polar domains.

Fig. 12

The nano-clusters observed [C _n mim][PF₆](n = 2,4,6,8,10,12) by molecular dynamic simulation [3]

Fig. 13

The changes of low-Q peaks of [C₆mim]Br and [C₁₀mim]Br along with the increasing temperatures from −100°C to 150°C. The inserted shows the SWAXS Q values at different temperatures

Except for the structural factors, the nano-clusters can be influenced by the temperature, pressure, and concentration in binary systems. Triolo et al. found that the low-Q peaks decrease with increasing temperatures in solid ILs, but the low-Q peaks increase with increasing temperatures in liquid ILs [59]. Figure 13 shows the low-Q peak curve of the [C₆mim]Br and [C₁₀mim]Br ILs with increasing temperatures [60]. It can be seen that the low-Q peaks of [C₁₀mim]Br decrease with increasing temperatures in the −100~−50 °C range, but the low-Q peaks increase linearly with increasing temperature in the −50~150°C range. Nevertheless, over the complete temperature ranges the low-Q peaks are very low, indicating that the nano-clusters exist throughout, even at a temperature as high as 150 °C, as shown by the Q values obtained at different temperatures. For [C₆mim]Br the situation is not the same. The low-Q peaks increase linearly with temperature, increasing over the 0~150 °C range, but the change is not obvious.

In general, the ions in the bulk phase are not randomly packed and the electrostatic force, van der Waals, and H-bonds rationalize the preferable arrangement.

3 Non-covalent Interaction in ILs

3.1 Binding Energy of Ion Pair

S. Zahn et al. compared the binding energies of one [Mmim]Cl with the NaCl ion pair by the symmetry-adapted perturbation theory (SAPT) method, considering different contributions in analogy to a multipole expansion (Fig. 14) [61]. For the NaCl (red diamonds in Fig. 14a) the dispersion term is negligible, whereas this contribution is comparable in magnitude to the induction term for the two conformers of the ionic liquid pair [Mmim]Cl (blue and green diamonds). The total energy of NaCl consists of only electrostatic, exchange, and induction contributions (see Fig. 14b; the curve with red squares almost exactly matches the curve with diamonds). We can make another interesting observation concerning the minima. Whereas the NaCl features the minima for all curves exactly at the equilibrium distance (at zero, see black dotted vertical line), this is, surprisingly, not the case for the [Mmim]Cl pairs (see black arrows). This means that the equilibrium distance is not exclusively determined by the most important attractive force, namely, the electrostatic interaction.

Fig. 14

(a) Contributions of different interactions to the total attractive (electrostatic + induction + dispersion) interaction energy. (b) Plot of the interaction energy versus distance [61]

This finding demonstrated that the ions interact in the repulsive region of electrostatic potential only if one considers that the hypothetical potential consists of electrostatic and exchange terms. This implies that ILs are not as ionic as one may naively imagine. Different contributions partly compensate each other, resulting in a very shallow potential energy curve. Consequently the system is highly flexible and liquid like, and is also able to adjust easily to different situations, which might also explain the good solvating properties of ILs.

Table 5 also lists the binding energies of [Emim]⁺-based complexes with different counter-anions. The binding energies follow the trend [CF₃CO₂]⁻ > [BF₄]⁻ >[CF₃SO₃]⁻ > [Tf₂N]⁻ ~ [PF₆]⁻. However, the experimental properties (melting point, density, self-diffusion coefficient, and molar conductivity) have no defined correlation with these binding energies, which indicates that many properties of RTILs are not determined just by interaction energy.

Melting point is the basic property of ILs and reflects the molecular packing and interactions of cations and anions. A rough correlation of melting points (T _m, K) and binding energies (E, kJ · mol⁻¹) for dialkylimidazolium chloride, bromide, dialkylimidazolium tetrafluoroborate, and dialkylimidazolium hexafluorophosphate ILs is shown in Fig. 17 [47]. A linear relationship was assumed for these ILs and their melting point decrease when the sizes of the N-alkyl side chains increase. Seven species of the Cl⁻ and Br⁻ series, except for [Pmim]Br, were correlated and the linear relationship is very pronounced. The melting points range from 300 to 400 K while the interaction energies lie between −375 and −320 kJ · mol⁻¹. A similar trend was observed for four species of the [PF₆]⁻ series, but the linear relationship is very gross. For four species of [BF₄]⁻, the linear relationship between melting points and interaction energies is also pronounced, but the difference between successive melting points is much greater than the difference between successive energies: the melting points are in a range 400–200 K while the interaction energies are between −350 and −340 kJ · mol⁻¹. Therefore, it is indicated that the more atoms the anions include, the more complicated are the interactions between cations and anions.

Molar conductivity is an important property of RTILs for their application as electrolytes for electrochemical devices. S. Tsuzuki et al. [18] reported the measurement of the molar conductivity of RTILs composed of [Bmim]⁺ cation. They found that the molar conductivity obtained from impedance measurement (Λ_imp) is smaller than that obtained from ion diffusivity measurement by NMR (Λ_NMR). They pointed out that the Λ_imp/Λ_NMR value provides useful information on the ionic dissociation/association dynamics in the ionic liquid (Table 6). They reported that the Λ_imp/Λ_NMR values of RTILs composed of [Bmim]⁺ follow the order [PF₆]⁻ >[BF₄]⁻ > [Tf₂N]⁻ > [CF₃SO₃]⁻ > [CF₃CO₂]⁻. A small Λ_imp/Λ_NMR value shows that an ion prefers to move with a counterion. Λ_imp/Λ_NMR is an important parameter for designing highly conductive ionic liquids. The RTILs composed of round-shaped anions ([PF₆]⁻ and [BF₄]⁻) have larger Λ_imp/Λ_NMR values than those composed of the other rod-shaped anions. The round-shaped anions dissociate and associate with cations more easily in the RTILs. The interaction energy of the [Emim]⁺ complex with [PF₆]⁻ anion(−78.4 kcal · mol⁻¹) is smaller than that of the [BF₄]⁻ complex (−85.2 kcal · mol⁻¹). The [BF₄]⁻ binds with [Emim]⁺ more strongly than the [PF₆]⁻, which well explains the smaller Λ_imp/Λ_NMR value for [Bmim][BF₄] than that for [Bmim][PF₆]. The interaction energy of complexes with [Tf₂N]⁻, [CF₃SO₃]⁻, and [CF₃CO₂]⁻ are⁻78.8,⁻82.6, and −89.8 kcal · mol⁻¹, respectively. The order of Λ_imp/Λ_NMR of three RTILs composed of [Bmim]⁺ and the rod-shaped anions coincides with the reverse order of the magnitudes of the interaction energies of the [Emim]⁺ complexes.

Table 5 Binding energies and properties of ILs

Comparison of the calculated interaction energies of the ion pairs and experimental Λ_imp/Λ_NMR values of RTILs suggests that three factors play important roles in determining the Λ_imp/Λ_NMR value: (1) RTILs have large Λ_imp/Λ_NMR values if the interaction between cation and anion is small; (2) RTILs that consist of round-shaped ions have large Λ_imp/Λ_NMR values – probably the round shape of the ions enables a high mobility of the ions in RTILs; (3) RTILs have large Λ_imp/Λ_NMR values if the orientation dependence of the interaction is small. A large anisotropy of the interaction would probably decrease the mobility of the ions.

3.2 Hydrogen Bonds

3.2.1 Structural Characteristic of H-Bonds

The particular resonance structure and high π electronic delocalization of the imidazolium cation result in large positive charges on the C₂–H, C₄–H and C₅–H groups that are the main interactive positions with anion [45, 62]. As discussed previously, H-bonds are explicit non-covalent interactions in ILs, but are not the same with different anions. For the anions with one atom, like the [Emim]Cl ion pair as shown in Fig. 15a, only one C–H…Cl H-bond forms between cation and anion. For anions with multiple atoms, like [Emim][BF₄] in Fig. 15b, more than one C–H…F H-bond forms between cation and anion [47].

Fig. 15

The H-bond denoted by a dashed line in imidazolium-based ionic liquids. (a) The optimized geometry and molecular orbital of [Emim]Cl ion pair by DFT method. (b) The optimized geometry and molecular orbital of [Emim][BF₄] ion pair

The electronic structure investigation found that the HOMO of Cl⁻ can interact with the LUMO of the cation to form a σ-type orbital overlap such that the Cl⁻ anion can localize near to the C₂–H group as shown in the [Emim]Cl ion pair [47, 63]. A similar σ-type orbital overlap can also be found in the [Emim][BF₄] ion pair. Therefore the orbital overlap results in the significant energy reduction, further rationalizing the preferable location of the anion.

3.2.2 Evidence of H-Bonds in ILs

Many ILs have been studied by experimental spectroscopy (IR, Raman, NMR, neutron scattering, etc.) and results have indicated that the ions can array orderly over a larger range of temperatures [36]. The earliest X-ray studies found that there existed a discrete H-bond in halide-based ILs (anions = Cl⁻, Br⁻, and I⁻) [64]. The evidence from X-ray and IR spectra from [Emim]Cl IL have demonstrated unambiguously that the H-bond not only exists but the is also most explicit interaction between the ions (Table 6). The Cl⁻ anions are located with the cation by H-bonds rather than at random.

Table 6 Experimental spectroscopic studies on interactions and H-bonds in some typical ILs

For the fluoro-anion-based ILs (e.g., anion = [BF₄]⁻, [PF₆]⁻, [(FH)_nF]⁻, [AsF₆]⁻, [OCOCF₃]⁻, [OSO₂CF₃]⁻, etc.), X-ray studies have indicated that F atoms of the anions can form C–H…F H-bonds with the imidazolium cation, in particular C–H₂, C–H₄ and C–H₅ groups on the ring as listed in Table 6 [36]. NMR studies have also provided structural depictions for [Mmim][PF₆] IL [72], in which the asymmetric unit contains one ion-pair, and the imidazolium cations form a weakly C–H…π hydrogen-bonded zigzag chain along the (0 0 1) direction. The [PF₆]⁻ anions are located between these chains and make closest contacts with the hydrogen atoms of the cation. The liquid phase of the [Dim][PF₆] shows a similar structure to the crystal state by neutron diffraction [79].

In crystal structures of dialkylimidazolium salts containing [OSO₂CF₃]⁻ or [OCOCF₃]⁻ anions, for example [Bmim][OSO₂CF₃] IL, the imidazolium ring and the -SO₃ group of the anion are connected by the C–H…O H-bonds to form a two-dimensional NaCl-like arrangement in each layer, and the butyl chain of the cation and the -CF₃ group of the anion protrude out of the layer in the same direction. Therefore, this structure contains a layer of butyl and -CF₃ groups sandwiched between the two layers. In the liquid state this more ordered structure is not obvious, but except for the C–H…O H-bonds between the cation and the -SO₃ group of the anion, the -CF₃ group of the anion can connect with the other cation by C–H…F H-bonds [39]. For [Dim][Tf₂N] ILs, both crystallographically distinct [Tf₂N]⁻ anions adopt a cis-conformation [80]. The crystal lattice of [Dim][Tf₂N] consists of alternating two-dimensional sheets with an –AA–B–AA– pattern.

Neutron diffraction studies on some RTILs have indicated that charge ordering endures in the liquid phase. In the [Mmim]Cl and [Mmim][PF₆] salts, the molecular packing in the first two or three solvation shells are similar both in the crystal and in the liquid, although the atomic distances are altered in the liquid [70, 73]. In comparison, the [Mmim][Tf₂N] salt showed smaller charge ordering [79] because of the diffuse charge density and larger size of the [Tf₂N]⁻ anion. Moreover, its liquid structure is poorly correlated to its crystal structure [80]. This was attributed to the conformational flexibility of the [Tf₂N]⁻ anion [79, 80]. Nevertheless, it remains clear that ILs form “quasimolecular” structures through three-dimensional H-bonded networks. The network is maintained to a great extent even in solutions having low dielectric media, making the RTILs highly organized media.

Further evidences have demonstrated that by substituting a hydrogen atom for a -CH₃ group, which is incapable of H-bond formation, the bands associated with stretches or bends of H-bonds disappear completely. R. Ludwig et al. [48, 76, 77] had demonstrated that the bands that relate to H-bonds disappeared in [Emim][Tf₂N] IL when the proton atoms on the imidazolium-ring were replaced by -CH₃ groups. With the increased abilities and strengths of H-bonds, the [Tf₂N]⁻ anion was found in the cis conformation. Switching off these local interactions leads to the energetically favored trans conformation. These structural effects caused by H-bonds should have significant influence on IL properties, such as melting points and viscosities, as discussed later. Moreover, the vibrational frequencies of H-bonds themselves can be observed directly by far-IR spectrum even for the complicated [Dmim][Tf₂N] IL as listed in Table 6. Wavenumbers above 150 cm⁻¹ were assigned to intramolecular bending modes of cations and anions in the ILs, and the wavenumbers below 150 cm⁻¹ were assigned to the bending and stretching vibrational modes of H-bonds [81].

The solvent properties of these ILs have also been investigated using chromatographic techniques [82, 83]. It is generally found that the ILs may be considered to be polar phases with the solvent properties being largely determined by the ability of the salt to act as an H-bond donor and/or acceptor. Table 1 lists the H-bonded donation abilities (HBDA, α) to denote qualitatively the acidity of ionic liquids. The α of the pyridinium-based and the quaternary amino ILs are obviously smaller than that of the imidazolium cation, which indicates that the interactions of H-bonds are not remarkable in these ILs.

3.2.3 Local H-Bonded Network

As previously described, involved in the interactions of many ions in bulk ILs, the forming H-bonds can be extended to form a three-dimensional network as in liquid water [84]. The network can even be observed on a smaller nanometer scale domain or in an ionic cluster in which non-homogeneity of local structure is found in ionic liquids. A network of ammonium nitrate IL (ANILs) is shown in Fig. 16, in which cations and anions are connected together by H-bonds (O atom as acceptor on anion and N–H as donor on cation) [48].

Fig. 16

The ion clusters and the local H-bonded network in ILs. (a) The ion cluster of dimethylammonium nitrate (DMAN) and the measured low-frequency vibrational FTIR spectrum of at 353 K compared to the vibrational modes of the corresponding PIL clusters [(DMA)(NO₃)]x with x = 1, 2, 3, 4, and 6 calculated by DFT at the B3LYP/6-31 + G* level of theory [77]. (b) Schematic representation of H-bond network IN [Emim]Cl IL. All of the H-bonds were are denoted by dashed lines [85]

Fig. 17

The correlation of melting points and interaction energies for dialkylimidazolium ionic liquids: Cl and Br salts series (filled triangles); [BF₄] salts series (filled circles); and [PF₆] salts series (filled diamonds) [62]

Crystallographic studies had also found that the H-bonded networks exist below melting points or at glass state temperatures in many ILs [35, 36, 86]. A remarkable three-dimensional H-bonded network of the [Emim]Cl [86] from X-ray and DFT calculations is shown in Fig. 16b, in which the asymmetric unit contains four Emim⁺…Cl⁻ ion pairs. The [Emim]⁺ ions cluster in four distinct layers perpendicular to the c axis. Similarly, the arrangement of Cl⁻ anions is a layered one. Each Cl⁻ interacts with three Emim⁺ cations and each [Emim]⁺ is associated with the three nearest Cl⁻ ions. The distance from Cl⁻ to a ring carbon atom averages 3.55 Å. The Cl⁻ anions are situated in H-bonded positions rather than randomly. In the [Bmim][PF₆] IL, the cations form an H-bonded zigzag chain motif via methyl hydrogen atoms and π-electrons, and [PF₆]⁻ anions are located between these chains with closest contacts with the methyl hydrogen atoms. Such connection results in the formation of an H-bonded network [72, 87].

Furthermore, molecular dynamical calculations have proved that ions clusters, when viewed over the course of a typical simulation (generally several nanoseconds), remain locked in a local place and the cations and anions cease to behave as discrete units but act rather as larger supermolecules, H-bonds being responsible for the spatial heterogeneity [52, 88, 89]. Experimental optical Kerr effect spectroscopy [90] also demonstrated the occurrence of heterogeneous domains in 1,3-dialkyl imidazolium salts. The translational and rotational motion of these larger units and their interactions with neighboring molecules explain the high experimental viscosities [91].

4 Effect of Structure on Application of ILs

4.1 Extraction and Separation

Some ILs have been used as solvents in separation processes due to their recovery and recycling properties. Rogers et al. [92] have reported the partition of substituted-benzene derivatives between water and the hydrophobic [Bmim][PF₆]. Their studies have included the removal of sulfides [93] and nitrides [94] from diesel and gasoline, the separation of aromatics from aliphatics [95], and the removal of pollutants (e.g., phenols, dyes, organic acids) from water, and they have demonstrated the abilities of ILs [94, 96, 97].

Some natural bioactive homologues with high structural similarity have been effectively separated by ILs. Cao et al. [97] have shown that ILs can be used to separate the soybean isoflavone aglycones, known as genistein, daidzein, and glycitein, as shown in Fig. 18. Furthermore, we investigated the interactive mechanism and found that ILs can form H-bonds with the homologues, but the magnitudes of interactions decrease in the order genistein > daidzein > glycitein, which was consistent with the distribution coefficients [98].

Fig. 18

Soybean isoflavone aglycone homologues

Another example is to separate the α-tocopherol from the four tocopherol homologues. The four homologues have different H-bonded acidities resulting from their structural differences in the number and position of methyl groups on the chromanol head. [Bmim]Cl, [Bmim][CF₃SO₃], and [Bmim][BF₄] ILs can extract α-tocopherol successfully from biphasic material in the order [Bmim][BF₄] < [Bmim][CF₃SO₃] < [Bmim]Cl under the same conditions [96]. This order is consistent with the that of these ILs’ hydrogen-bond basicity strength [99]. This phenomenon further disclosed that the selectivity was based on the hydrogen-bonding interaction between ILs’ anion and the -OH group on the tocopherols.

4.2 Dissolution of Cellulose

Biocompatible composites generated from renewable biomasses are regarded as promising materials that could replace traditional polymers and reduce global dependence on fossil fuels [100–102]. Rogers et al. applied ILs as green solvents to dissolving cellulose, which has opened up a new area for application of ILs in bio-energy [103]. The experiments have proved that many ILs, such as [Bmim]Cl, [Amim]Cl, [Emim]Cl, [BMPy]Cl, [Bmim][OAc], [Amim][OAc], [Emim][OAc], [Emim][XS], and [N₁₁₁C₂H₄OH][OAc], can effectively dissolve the biomass from wood, bagasse, wool, silk, chitin, etc. [100, 103–106]. Experiments and calculations have both proved that the formation of H-bonds between ILs and biomolecules play crucial roles in the dissolving process [107–109]. By means of¹³C and^35/37Cl NMR relaxation measurements, R. C. Remsing et al. demonstrated conclusively that the solvation of carbohydrates by [C₄mim]Cl involves stoichiometric H-bonding between the hydroxyl protons of the cellulose and the chloride ions of the IL. There may be two steps for the dissolution; the first is to break the H-bonds between anions and cations, then the active chloride anions interact with the hydroxyl groups of cellulose by forming H-bonds. It is also thought that [C₄mim]⁺ cations may possibly form H-bonds with cellulose and play a non-negligible role in the dissolution process.

Figure 19 showed the proposed dissolution mechanism of cellulose in ILs, and how the oxygen atoms and hydrogen atoms of the –OH groups on cellulose form electron donor-acceptor complexes. The anion can combine with hydrogen atoms of cellulose by H-bonds, whilst the imidazolium cation interacts with the oxygen atoms by hydrogen bonding. Comparing with water, we found that the binding energies of the H-bonds between cellulose and ILs are three times stronger than that between cellulose and water, indicating that cellulose is more easily dissolved in [Emim][OAc] [109]. The presence of water in the IL was shown to decrease significantly the solubility of cellulose, presumably through competitive hydrogen-bonding to the cellulose microfibrils which inhibits solubilization. This phenomenon can also corroborate that hydrogen bonding is a key interaction in the dissolving process of cellulose in ILs.

Fig. 19

A proposed mechanism of dissolution of cellulose in ionic liquids. Dashed lines show the hydrogen bonds

The increasing basicity of the anion resulting in the higher solubility of biomass is also in agreement with the H-bonding mechanism. C. S. Pereira recently reported that a class of biocompatible and biodegradable cholinium-based ILs, the cholinium alkanoates, showed a highly efficient and specific dissolution of the suberin domains from cork biopolymers. They found that the cholinium alkanoates of increasing alkyl chain length (ethanoate, butanoate, hexanoate) showed augmented dissolution efficiency since the basicity of the anion increases with chain length [110]. It indicates clearly that the H-bond basicity of the solvent system is essential for the dissolution of cellulose.

4.3 As Catalyst

Diels-Alder cycloaddition reactions have been explored using various ILs substituting for water. The reaction of cyclopentadiene with methyl acrylate and methyl vinyl ketone in IL is shown in Fig. 20 [111]. The studies have found that many ILs can influence the selectivity of the endo/exo. The theoretical investigation had demonstrated that the ability of the imidazolium cation to act as H-bond donor is responsible for the selectivity. The formation of H-bonds from the cation to dienophile is a Lewis acid–base interaction, as it is well known that Lewis acid catalysts can have a dramatic effect on both the rates and selectivities of Diels–Alder reactions [112].

Fig. 20

Diels–Alder reaction by ILs catalysis and the H-bonds (Lewis acid) between imidazolium cation with the carbonyl oxygen of methyl acrylate

4.4 Gas Absorption

Studies have shown that many environmental pollution gases, such as CO₂, SO₂, and NH₃, can dissolve freely in ILs [25, 113–115]. The absorption of CO₂ in [Bmim][PF₆] reaches a mole fraction of 0.6 at 25°C and 8.0 MPa [116], and [Bmim][BF₄] and [TMG][BF₄] ILs can absorb 1–2 mole SO₂ per mole of IL reversibly at the ambient pressure and temperature [117].

Figure 21 shows a physico-absorptive mechanics of CO₂ in [Bmim][PF₆] IL. When CO₂ molecules enter the bulk phase of the IL, the O atoms of CO₂ form H-bonds with C–H groups of the imidazolium-ring, and C atoms (with positive charges) can be attracted by anions due to electrostatic interaction. The line-like nonpolar structure of CO₂ is thus changed to give a polar molecule (the angles of the ∠O–C–O are distributed between 175 ~ 178° according to ab initio molecular dynamic calculations, so the change of CO₂ structure results in easier absorption in ILs. The theoretical studies have also indicated that the “spontaneously forming cavities” by H-bonds are vital to accommodate CO₂ and SO₂ molecules when they enter the bulk of ILs [118, 119]. However, particularly Ludwig et al. reveal that the experimentally observed “anomalous” temperature dependence of CO₂ in ILs is not due to cavity effects but is caused by attractive solute–solvent interactions. The contribution of the van der Waals interactions is found to dominate [120, 121].

Fig. 21

Proposed mechanics of CO₂ absorption in ILs. The relative positions and the H-bonds change when CO₂ accommodates into the IL physically. Red dashed lines show the relative position and green dashed lines show the hydrogen bonds. Arrows show the electrostatic interactions

5 Perspective

In the past two decades, explorations of ILs have experienced a fast development and the knowledge base of interdisciplinary data is rapidly being generated for ILs. This will fuel innovative ideas to take these liquids far beyond the realm of mere solvent. However, due to the complications and diversities in ILs’ family, it is very difficult to clarify all of the structural factors exerting effects on their properties.

The molecular interactions between ions depend on their geometry and charge distribution. Unlike in simple salts, the interactions are controlled by long-range Coulombic forces between the net charges of the ions. As for ILs with molecular ions, their bulky size and asymmetric charge distribution softens the Coulombic forces and generates highly directional interactions of shorter range. Therefore, almost all studies on the structures of ILs have involved the non-covalent interactions between cations and anions. In most ILs, in particular imidazolium-based ILs, the H-bonds present the structural characteristics and form the three-dimensional local network to connect the discrete ions. The melting points, viscosity, hydrophilic and hydrophobic properties, gas absorption, dissolution of cellulose, etc, are closely related to the interaction energy and H-bonds.

As for other chemistries, for example nano science and biotechnology, ILs give people the potential to deal with some unresolved problems and fulfill a task in green chemical processes. Structural investigations can further provide a deep understanding of the unique properties of ILs and valuable hints for the design of new ILs. Based on this, we believe the following aspects are the most important for future research on ILs. First is to know the quantitative relationship between the structures and the properties of ILs, which will point us in the direction of how we may design efficient ionic liquids for a special application. Another aspect is to study how to scale up the systems with ILs for the chemical engineering processes. Additionally, with the development of materials science and engineering, ILs have found a new stage on which to play a role.