Ionic liquids as model systems for the study of the behaviour of electrolytes and surfactants in solutions

Ion mobility and ion association of ILs in dilute solutions

The mobility and mechanism of ion pairing in aqueous solutions were systematically studied on nine imidazolium based ILs: 1-methyl imidazolium, [mim]Cl; 1,2-dimethyl imidazolium, [1,2-mim]Cl; 1,3-dimethyl imidazolium [1,3-mim]Cl; 1-ethyl-3-methylimidazolium, [C₂mim]Cl; 1-butyl-3-methyl imidazolium, [C₄mim]Cl; 1-hexyl-3-methyl imidazolium, [C₆mim]Cl; 1-octyl-3-methyl imidazolium, [C₈mim]Cl; 1-decyl-3-methyl imidazolium, [C₁₀mim]Cl; and 1-dodecyl-3-methyl imidazolium, [C₁₂mim]Cl [6]. The values of their molar electrical conductivity – still one of the most suitable methods for studying electrolyte solutions – for dilute solutions of the investigated imidazolium based ILs at 298.15 K are shown in Fig. 6, together with the data for aqueous NaCl solution and the results of data analysis within the framework of Barthel’s low-concentration chemical model (lcCM).

Fig. 6:

Molar conductivity of the investigated imidazolium based ILs chlorides [6] and NaCl [22] at 298.15 K in water; symbols denote experiment and lines lcCM calculations. Figure reproduced from reference [23] with permission of Acta Chimica Slovenica.

As can be seen here, the values for NaCl are only slightly higher than those for [mim]Cl. Thus, it can be assumed that the investigated ILs in dilute aqueous solution behave like simple 1,1-electrolytes – the molar conductivity is linearly dependent on the square root of concentration.

lcCM successfully describes the thermodynamic and transport properties of dilute solutions (c ≈ ≤0.005 M) and provides the molar conductivity of the solute at infinite dilution, Λ^∞, and the standard-state (infinite dilution) ion association constant, K _A, (details are given in the literature [24]). Using the literature values for the ionic conductivity of chloride ion at infinite dilution ( λ Cl − ∞ ( 298.15 K ) = 76.35 S cm 2 mol − 1 for example), the limiting molar conductivity was split into separate ionic contributions, λ i ∞ , and then the values of limiting ionic conductivity were estimated for all cations. Consequently, diffusion coefficients, D i ∞ , were calculated as characteristic properties of ion transport that are not affected by interionic interactions, using the following relationship.

Here k is the Boltzmann and F the Faraday constant, T is the absolute temperature, e _o is the proton charge, and z _i is the valence of the ion.

Values of D i ∞ for ILs cations, studied in [6] at 298.15 K are presented in Fig. 7 as a function of the number of C atoms in the alkyl side chain, along with values obtained from MD simulations performed at 298 K [6].

Fig. 7:

Diffusion coefficients of imidazolium-based cations at infinite dilution, D + ∞ , in water at 298.15 K as a function of the number of carbon atoms in the side chain, C _n : () experiment; () MD simulations. Figure is reproduced from reference [6] with permission from PCCP Owner Societies.

Evidently, the diffusion coefficient of [mim]⁺ is only slightly lower than that of Na⁺, D Na + ∞ ( 298.15 K ) = 13.3 × 10 − 10 m 2 s − 1 , and D + ∞ decreases with increasing length of alkyl chain (Fig. 6), but not linearly as it has been also obtained in MD simulations [6, 25].

However, the length of the side chain is not the only factor affecting the conductivity of the system. It can also be lowered by possible ion associations forming nonconducting ion pairs or even higher associations. The most energetically favourable position of Cl-relative to the cations is shown in Fig. 8 with red circles, represented with the potential of mean force (PMF) as obtained from statistical mechanics calculations within the framework of integral equation theory using the three-dimensional (3D) Reference Interaction Site Model (RISM) approach [6].

Fig. 8:

2D-maps of the PMF between cation and anion for (a) [mim]Cl; (b) [1,2-mim]Cl; (c) [1,3-mim]Cl; (d) [C₂mim]Cl; (e) [C₄mim]Cl; (f) [C₆mim]Cl; (g) [C₈mim]Cl; (h) [C₁₀mim]Cl and (i) [C₁₂mim]Cl. The most probable locations of the Cl⁻ are indicated by red circles. Figure is reproduced from reference [6] with permission from PCCP Owner Societies.

Analysis of the experimental data by lcCM also provides the values of the association constants, K _A [24]. For the imidazolium based ILs in water studied in reference [6] the K _A values were found to be small (∼2.5 ≤ K _A ≤ ∼6), as expected but significantly higher than those obtained with the same model for alkali metal halides in water [26].

For [C₄mim]⁺, [C₆mim]⁺, and [C₁₂mim]⁺, the binding free energies with Cl⁻, Δ G bind MD , were obtained also from MD simulations, using thermodynamic integration as specified in reference [6]. From the values of K _A, estimated from the electrical conductivity data, the binding free energies were calculated using the relation Δ G bind exp = − R T ⋅ ln   K A . In Fig. 9, the comparison of Δ G bind exp and Δ G bind MD is shown. Evidently, results of the MD simulations are in reasonable agreement with experiment.

Fig. 9:

Comparison of binding free energies, Δ G bind , for imidazolium based chlorides in water as a function of the number of C atoms, C _n , in the side alkyl chain as obtained from conductivity experiment, Δ G bind exp , and from MD simulations, Δ G bind MD . Adopted from reference [6] with permission from PCCP Owner Societies.

Thus, theory, MD simulations and experiments show that the association of the studied ILs as model 1,1-electrolytes in water solutions is weak, but obviously depends on the molecular structure (alkyl chain length), which also strongly influences the mobility of cations. For the study of ion association, water is not very suitable as a solvent due to its high dielectric constant. Numerous studies on the conductivity of mixtures of ILs with molecular solvents can be found in the literature. Surprisingly, despite the great practical interest in such systems, there are few systematic studies on the transport properties of binary mixtures of ILs and co-solvents in a wide range of concentrations – from dilute solutions to pure ILs. Therefore, in this review we will focus mainly on imidazolium based ILs, which are also well studied in a broader range of concentrations in aqueous and non-aqueous solutions, allowing a comparison of their properties.

Figure 10(a) shows the molar conductivity of [C₄mim]Cl in acetonitrile (AN) [27], methanol (MetOH) [5] and dimethyl sulfoxide (DMSO) [5] along with the values in water [6]. Evidently, the mobility of [C₄mim]⁺, decreases with increasing viscosity of the solvents as indicated by the ionic limiting conductivity λ [ C 4 mim ] + ∞ and the diffusion coefficients D [ C 4 mim ] + ∞ (Table 1).

Fig. 10:

Molar conductivity, Λ, of (a) [C₄mim]Cl in AN [27], water [6], MetOH [5] and DMSO [5] and (b) [C₄mim]Cl [27], [C₄mim]BF₄ [4] and TBABr in AN [28] at 298.15 K; symbols denote experiment and lines lcCM calculations. Figure reproduced from reference [23] with permission of Acta Chimica Slovenica.

Figure 10(b) shows the molar conductivity of [C₄mim]Cl [27], [C₄mim]BF₄ [4] and, for comparison, tetrabutylammonium bromide, TBABr [28], as a “classical” electrolyte in AN. While for [C₄mim]Cl only a weak association in water was observed (K _A ≅ 5–6), it is more pronounced in solvents with lower dielectric constants, as expected, but obviously also depends strongly on the anion, as can be seen from Table 1, which presents the literature data of limiting molar conductivity, Λ ^∞ , and association constants, K _A, for [C₄mim]Cl and [C₄mim]BF₄ in water, DMSO, AN, N,N-Dimethylformamide (DMF), MetOH and dichloromethane (DCM) at 298.15 K, as obtained from the lcCM model. When the value of the limiting ionic conductivity λ [ C 4 mim ] + ∞ is available, the (average) values of the diffusion coefficient D [ C 4 mim ] + ∞ are also listed

While the values of Λ^∞ are not very “sensitive” to the applied model for the analysis of conductivity data, the opposite is true for K _A. Even more, the values of K _A also depended on the upper distance limit set, where the ions are still treated as ion pairs. However, the many different interactions in ILs make them very complex (Fig. 2), so it is not unexpected that the dielectric constant is unable to adequately model solvent-solvent interactions and has often failed in qualitatively and quantitatively correlating solvent effects, It is indeed surprising that the K _A values for [C₄mim]Cl and [C₄mim]BF₄ in AN and MetOH are different although the dielectric constants of these two solvents are very similar. However, the K _A values obtained for [C₄mim]BF₄ and [C₄mim]Cl in AN are in agreement with those reported for sodium tetraphenylborate (K _A = 13.9 at 298.15 K [31]) and TBABr (K _A = 27.5 at 298.15 K [28]).

Despite the fact that conductivity measurements on dilute electrolyte solutions are probably still the most accurate method for estimating the ion-pair association constant, K _A, at least for symmetrical electrolytes [24], such studies can only determine the overall association and therefore provide little information about the nature of the aggregate(s) formed, whereas conventional spectroscopic techniques such as NMR or Raman spectroscopy generally detect only contact ion pairs, CIPs [32]. On the other hand, dielectric relaxation spectroscopy, DRS, is sensitive to all types of ion-pairs and allows their identification and quantification, provided reasonably accurate dipole moments, μ _i, of the formed species are available or can be calculated [33, 34]. If we consider the DRS spectra for the [C₄mim]Cl [27] solution in water (Fig. 11(a)) and in AN [27] (Fig. 11(b)), we see a region corresponding to ion pairs, IP, and this region is much larger in AN than in water. Thus, the tendency to form ion pairs is obviously stronger in AN, and the values of K _A, obtained from the conductivity data using the lcCM model, are reasonable. In turn, one could conclude that ILs behave like ordinary (“classical”) simple electrolytes in solutions, with ion association being more pronounced in solvents with lower dielectric constants.

Fig. 11:

Relative permittivity, ε′(ν) (■) and dielectric loss, ε″(ν) (●) spectrum of a representative [C₄mim]Cl solution in (a) water (c = 0.4618 mol dm⁻³) and (b) in AN (c = 0.4658 mol dm⁻³) at 298.15 K. The symbols represent experimental data, the lines show the corresponding fit, and the shaded areas indicate the contributions of each process to ε″(ν). Adopted from reference [27] with permission of the PCCP Owner Societies.

Electrical conductivity of ILs in concentrated solutions

Conductivity data published by Nordness et al. [35] for concentrated aqueous solutions of [C₂mim]Cl, [C₄mim]Cl, [C₆mim]Cl and [C₈mim]Cl at 298.15 K are shown in Fig. 12. To reproduce this type of concentration dependence, the empirical four-parameter Casteel–Amis equation [36] is often used. In general, data fitting is performed on the molality scale. For the present systems, the Casteel–Amis equation can be conveniently fitted on the mole fraction scale of IL, x _IL, i.e.

where, κ _max is the highest electrical conductivity for a given mixture, x _{IL, max} is the mole fraction of IL at which the electrical conductivity of the mixture reaches a maximum at a given temperature, and n and m are parameters of the fit. The Casteel–Amis parameters are given in reference [23] and the fitted conductivity values are shown in Fig. 12. In the inset of Fig. 12(a), the values for NaCl solutions are included for comparison.

Fig. 12:

(a) Specific conductivity of some imidazolium based ILs chlorides at 298.15 K in water as derived from literature data [35]. Inset: Specific conductivity of some ILs chlorides and NaCl at 298.15 K in water [37]. Symbols represent experimental values and the lines represent the Casteel–Amis fit (eq. (2)) of the experimental data with the parameters reported in reference [23]. Figure reproduced from reference [23] with permission of Acta Chimica Slovenica. (b) Viscosities of some ILs chlorides at 298.15 K in water. Drawn according to data from the literature [35].

It is obvious that the conductivity in all the systems shown in Fig. 12(a) increases sharply in the dilute range, which is due to the increased number of free ions in the solutions. However, with the addition of ILs, it reaches a maximum at nearly the same mole fraction, i.e., at about x _IL ≅ 0.05 and then decreases. This tendency is thought to be the result of competition between the increasing number of free ions that can contribute to conductivity and the increasing viscosity of the system, which hinders ion mobility (Fig. 12(b)). Again, conductivity decreases with increasing alkyl chain length of the cation, while viscosity is higher for ILs with a longer side chain. When ions are added to the system, ion diffusion is hindered by increased viscous forces, leading to a decrease in ion mobility and thus diffusion.

Incidentally, this behavior is not new – the same has been observed for “ordinary simple” electrolytes, where the maximum was often not reached due to limited solubility, as shown here for NaCl in water (inset in Fig. 12(a)). Therefore, ILs could be very useful in the study of concentrated electrolyte solutions, especially since the entire concentration range can often be covered. Figure 13(a) shows some literature data on the conductivity of concentrated solutions of [C₄mim]BF₄ in AN, MetOH, DMSO, DCM, and propylene carbonate (PC, η = 2.512 mPas, ε = 64.96 at 298.15 K) as reported by Stoppa et all [38]. The conductivity of three imidazolium based tetrafluoroborates in AN from the same work is shown in Fig. 13(b) along with data on TBABr in AN [39] for comparison.

Fig. 13:

Specific conductivity, κ, of (a) [C₄mim]BF₄ in AN [38], MetOH [38], DMSO [38], DCM [38], and PC [38], and (b) [C₂mim]BF₄ [38], [C₄mim]BF₄ [38], [C₆mim]BF₄ [38], and TBABr [39], in AN at 298.15 K as a function of molar fraction of IL, x _IL. Symbols represent experimental values and lines the Casteel–Amis-type fitting (eq. (2)) of the experimental data with parameters reported in reference [23]. Figure reproduced from reference [23] with permission of Acta Chimica Slovenica.

From Fig. 13 it can be seen that the conductivity follows the typical pattern of concentrated electrolyte solutions already discussed in Fig. 12(a) for aqueous solutions. After a rapid increase in the low concentration range, κ passes through a well-defined maximum. As can be seen from the data, presented in Fig. 13(b), κ for ILs in AN decrease with increasing alkyl chain length of the cation, as was also observed for the aqueous solutions of IL_S presented in Fig. 12(a). While in aqueous solution the position of the conductivity maximum was more or less near x _IL ≅ 0.05, here the maximum is reached at higher concentration and shifts towards lower x _IL with increasing side alkyl chain length of the cation.

Figure 13(a) clearly reveals that κ for [C₄mim]BF₄ depends strongly on the solvent: it decreases with increasing viscosity (AN < MetOH < DMSO < PC). However, DCM does not follow this order: despite the fact that its viscosity is very low (0.415 mPa s at 298.15 K), the conductivity of [C₄mim]BF₄ solutions in DCM is close to that in DMSO and PC at much higher solvent (Table 1). This behaviour can be attributed to higher ion association in DCM due to the lower dielectric constant. The value, determined by Borun and Bald [30] (K _A = 479 at 298.15 K, Table 1) in dilute solutions, clearly confirms this assumption. Even more – also in other solvents the decrease of conductivity after the maximum could be attributed to a stronger ion association. Unfortunately, there is no theory to describe this behaviour. Therefore, ILs could be very useful in the study of concentrated electrolyte solutions, for which reliable theories do not yet exist.

However, ILs with longer alkyl chain(s) in the cation do not follow this pattern but show the typical dependence of the solution conductivity on the molality of the surfactant (see Fig. 14(a)), so that at higher concentrations they form micelles and act as surface active ionic liquids.

Fig. 14:

(a) Specific conductivity, κ, of [C _x mim]Cl with longer alkyl chain at 298.15 K. drawn from literature data [40]. The arrows show the breaks at cmc; (b) enthalpograms for the same SAILs at 288.15 K in water [8]. The solid lines represent the fits according to the model function. Here, cmc can be estimated from the inflection point. Figure reproduced from reference [8] with permission from Elsevier.

Surface active ionic liquids (SAILs) in solutions

Ionic liquids have recently attracted much attention as a new class of surfactants. Indeed, ionic liquids with long alkyl chains behave similarly to conventional surfactants forming aggregates (micelles) in water at a concentration known as the critical micelle concentration, cmc. Figure 14 shows the behavior of [C₁₀mim]Cl, [C₁₂mim]Cl, [C₁₄mim]Cl and [C₁₆mim]Cl in aqueous solutions as determined by conductivity measurements (Fig. 14(a)) [40] and isothermal titration calorimetry, ITC (Fig. 14(b)) [8] which is typical for surfactant systems in aqueous solutions.

From Fig. 15(a) it can be seen that cmc decreases with the length of the alkyl side chain in the imidazolium cation at [C _x mim]Cl, as it is common for “classical” surfactants. An inset in Fig. 15(a) for [C₁₂mim]Cl clearly shows the U-shape for the temperature dependence of cmc, which is also typical for ionic surfactants if the experiment is performed in a sufficiently large temperature range.

Fig. 15:

Temperature dependence of (a) cmc (data derived from reference [8]) and (b) the standard enthalpy of micellization, Δ_M H ^θ for [C₁₀mim]Cl, [C₁₂mim]Cl, [C₁₄mim]Cl and [C₁₆mim]Cl in water [8]. Figure reproduced from reference [8] with permission from Elsevier.

From the temperature dependence of the enthalpy of micellization, Δ_M H ^θ (Fig. 15(b)), obtained from the ITC experiment using a suitable model equation (see [8] for details), it appears that the process is endothermic at lower temperatures and exothermic at higher temperatures, which is usually the case for ionic surfactants.

This behavior can be attributed to two opposing effects that influence the process of micelle formation [41]. The first effect is related to the hydration of the polar group of the surfactant molecule and plays a key role in the low temperature range, where the formation of hydrogen bonds is more likely, and the strong hydration of the polar head “penalizes” the micelle formation process. As the temperature increases, the hydration of the polar group decreases, allowing micelle formation to occur at lower concentrations. The second effect is related to the hydrophobic hydration of the alkyl chain, which changes the structure of the water surrounding the hydrophobic tail. It is known that the structure of water is strongly dependent on temperature.

For this reason, the structure of water surrounding the hydrophobic chains becomes looser and the effect of hydrophobic interaction becomes weaker as the temperature increases. This phenomenon leads to a shift of the micellization process to the higher concentration range. The competition between these two effects is the reason for the observed characteristic U-shaped dependence of cmc on T, with a minimum near the temperature where Δ_M H ^θ = 0 and the temperature is in the range of room temperature (Fig. 15).

The great potential of SAILs in studying the thermodynamics of the micellization process in solutions due to possible variations in the structure of cations, side chain length, and counterions has already been clearly demonstrated [8, 19, 20]

The effect of counterions on the micellization process was studied in aqueous solution of [C₁₂mim]Cl, [C₁₂mim]Br and [C₁₂mim]I and also on SAILs of [C₁₂mim]⁺ with other counterions, shown in Fig. 16.

Fig. 16:

Structures of investigated systems in [8]: (a) [C₁₂mim]⁺, (b) acetate, OAc, (c) trifluoroacetate, TFA, (d) salicylate, Sal, (e) methanesulfonate, OMs, (f) trifluoromethane sulfonate, OTf, and (g) toluene sulfonate, OTs anions in [C₁₂mim]X systems. Figure reproduced from reference [8] with permission from Elsevier.

According to Fig. 17(a), presenting the cmc for all systems studied in ref. [8] at 298.15, it depends strongly not only on the length of the alkyl side chain but also on the counterion.

Fig. 17:

(a) Critical micelle concentrations, cmc, for all systems studied in [8] at 298.15 K. (b) Thermodynamic parameters of micellization for the systems studied in water at 298.15 K: standard enthalpy, Δ_M H ^θ, Gibbs free energy, Δ_M G ^θ, and entropy contributions, TΔ_M S ^θ, as determined by the fitting procedure given in [8]. Figure reproduced from reference [8] with permission from Elsevier.

As mentioned earlier, the micellization process for SAILs is generally endothermic at low temperatures and exothermic at high temperatures, but counterions have been shown to play an extremely important role [8], as shown in Fig. 17(b). Apparently, counterion have a greater effect on Δ_M H ^θ than on Gibbs free energy, Δ_M G ^θ, and entropy Δ_M S ^θ of micellization.

Even more, it has been confirmed that the hydrophobicity of the counterions appears to contribute to the change in heat capacity and the reduction of the surface area accessible to water when the nonpolar group is removed from contact with water during the micellization process.

The main driving force for the formation of micelles is the apparent disaffinity of water and the nonpolar (interacting) surfaces known as hydrophobic effect. This effect is reflected in the large negative standard heat capacity changes ( Δ M c p θ << 0, evident from the temperature dependence of Δ_M H ^θ presented in Fig. 15) which are the hallmark of processes directing by removing of nonpolar surface from water, including the transfer of nonpolar solutes from water to a nonaqueous phase and the folding, aggregation/association, and ligand-binding reactions of proteins [42].

Spolar et al. [43] analysed the thermodynamic data for the transfer of hydrocarbons and organic amides from water to the pure liquid phase to obtain contributions to the thermodynamics of folding from the reduction in water-accessible surface area. Although the removal of nonpolar surface makes the dominant contribution to the standard heat capacity change of folding ( Δ fold c p θ ), it was show that inclusion of the contribution from removal of polar surface allows a quantitative prediction of Δ fold c p θ within the uncertainty of the calorimetrically determined value and the Δ fold c p θ can be calculated according to the expression

where ΔA _np represents the part of the water-accessible nonpolar and ΔA _p denotes the part of water-accessible polar area of molecule surface that is removed from contact with water during protein folding.

Micellization and protein folding are two different processes that occur in solutions, but they do share some similarities in terms of the interactions that drive their formation and stability. In micellization, the hydrophobic tails of the surfactant molecules come together to form a hydrophobic core, while in protein folding, the hydrophobic amino acid side chains are typically buried in the protein’s interior to minimize their contact with water. Thus, eq. (3) can be adopted also for modelling of Δ M c p θ which can be attributed to the removal of water molecules from contact with nonpolar surface area during micelle formation [44].

Since the hydrophilic head groups of nonionic surfactants remain hydrated during micelle formation, it can be assumed that the “theoretical” contribution of the change in water-accessible nonpolar surface area to the change in heat capacity during micelle formation, Δ M c p θ ( th np ) reflects only the change in exposure of the hydrophobic tails to water. Consequently, eq. (3), when applied for micellization process, reduces to

This approach has proven useful for a number of nonionic [45, 46], some cationic surfactants [47, 48] and has also been applied to SAILs.

According to Richards [49], the surface area accessible to water for a methylene group is 30 Ų and 88 Ų for a methyl group. Thus, ΔA_np of the hydrophobic tails of the surfactants investigated in ref. [8] is 358, 418, 478 and 538 Ų for the C₁₀, C₁₂, C₁₄ and C₁₆ alkyl chains, respectively, giving values of Δ M c p θ (th_np) = −479, −560, −640 and −720 J K⁻¹⋅ mol⁻¹ for C₁₀, C₁₂, C₁₄ and C₁₆ alkyl chains, respectively. Comparison of Δ M c p θ (th_np), estimated according to eq. (4) with Δ M c p 0 (exp), as shown in Fig. 18(a), indicates that there is an obvious discrepancy between Δ M c p θ (th_np) and Δ M c p θ (exp). This discrepancy can be explained in two ways:

1. It can be assumed that there are still water molecules inside the micelle after micellization, and thus the ΔA_np estimated with the Richards [46] approach mentioned above is overestimated. That is, the “real” ΔA_np is smaller than that according to Richards.

2. Equation (4) holds well for nonionic surfactants [44, 45], but is probably not valid for ionic surfactants. The positive contribution of removal of the water-accessible polar surface to the change in heat capacity during micellization should also be considered and cannot be neglected, consistent with the findings on the change in heat capacity during protein folding [43].

Fig. 18:

Comparison of theoretical (shaded) and experimental (full columns) values of heat capacity changes in the micellization process, Δ M c p θ , for (a) [C _n mim]Cl and (b) [C₁₂ mim]X. Figure reproduced from reference [8] with permission from Elsevier.

The influence of the counterion on Δ M c p θ for [C₁₂mim]X systems [8], is shown in Fig. 18(b). For [C₁₂mim]Cl, [C₁₂mim]Br and [C₁₂mim]I it is obvious that the difference between Δ M c p θ (exp) and Δ M c p θ (th_np) increases with increasing size of the anion. It is plausible to assume that the removal of the polar surface accessible to water also increases with the increase in the size of the counterion binding to the micelle, but it can be concluded that the ΔA _p cannot be neglected here or/and that the hydrophobic parts of the counterions are incorporated into the micelles. In this case, it can be assumed that the contribution to Δ M c p θ comes from the removal of the hydrophobic part of the counterion from contact with water after micellization. Thus, the “hydrophobic” anions (OTs⁻, OTf⁻, TFA⁻ and Sal⁻) can be partially incorporated into the micelle and therefore the hydrophobic part of the anion can contribute to the ΔA_np. The incorporation of counterions has already been confirmed for [C₁₆mim]OTs by Singh et al. using various methods [50] and for [C₁₂mim]OTf using DRS [51].

SAILs have also been shown to be excellent model systems to demonstrate the influence of counterion isomerism on the micellization of surfactants in water [20]. This was studied on 1-dodecyl-3-methylimidazolium ([C₁₂mim]⁺) benzoate (Bz⁻), ortho- (o-HBz⁻), meta- (m-HBz⁻), and para-(p-HBz⁻) hydroxybenzoate (Fig. 19) [20].

Fig. 19:

Structures of the counterions together with the representation of the surface distribution function, SDFs, of water around them as obtained by molecular dynamic simulations: (a) Bz⁻, (b) o-HBz⁻; (c) m-HBz⁻, and (d) p-HBz⁻,included in the study of the influence of counterion isomerism on the micellization of [C₁₂mim]⁺ in water [20]. Figure reproduced from reference [20] with permission from Elsevier.

It was found that the temperature dependence of cmc also shows a nearly U-shaped form (Fig. 20(a)) but depends strongly on the presence and position of the –OH group in the counterion. The micellization process is endothermic at low temperatures and becomes exothermic at higher temperatures – a common phenomenon for ionic surfactants – for all investigated systems, only the micellization of [C₁₂mim]o-HB is exothermic over the whole temperature range studied. Namely, it is known that o-HBz⁻ causes the formation of more compact, elongated micelles due to a stronger interaction of o-HBz⁻ with the surfactant molecule and its subsequent incorporation. This can be explained by the different hydration, which is presented by the surface distribution function, SFD, of water around the studied anions as obtained by molecular dynamic simulations (Fig. 19). It could be assumed that the strong hydration of the hydroxyl group prevents the incorporation of the m-HBz⁻ and p-HBz⁻ isomers into the micelle, leading to the highest exothermic micellization process in the case of o-HBz⁻. Further details can be found in [20].

Fig. 20:

Temperature dependence of (a) cmc and (b) ∆_M H ^θ _, for [C₁₂mim]HBz, [C₁₂mim]o-HBz, [C₁₂mim]m-HBz, and [C₁₂mim]p-HBz in water. The solid lines in (a) represent the corresponding polynomial fits, cmc = A + BT + CT ² (the coefficients are given in [20]). The solid lines in (b) represent linear fits. Figure reproduced from reference [20] with permission from Elsevier.

SAILs thus offer exceptional and valuable opportunities to alter the structures of counterions, which cannot be achieved with “ordinary” surfactants. This is likely to become even more evident when studying the aggregation properties of pyridinium-based SAILs to investigate the influence of the isomerism of the surfactant cation on the micellization process, which was studied on N-dodecyl-pyridinium bromide, [N-C₁₂-py]Br, and its three (ortho, meta, and para) methyl-substituted derivatives, N-dodecyl-2-methylpyridinium bromide, [N-C₁₂-2-mpy] Br, N-dodecyl-3-methylpyridinium bromide, [N-C₁₂-3-mpy]Br), and N-dodecyl-4-methylpyridinium bromide, [N-C₁₂-4-mpy]Br [21]. Their structures are shown in Fig. 21.

Fig. 21:

Structures of (a) [N-C₁₂-py]⁺, (b) [N-C₁₂-2-mpy]⁺, (c) [N-C₁₂-3-mpy]⁺, and (d) [N-C₁₂-4-mpy]⁺ together with the representation of the surface distribution function, SDFs, of water around them as obtained by molecular dynamic simulations included in the study of the influence of cation isomerism on the micellization in water [21]. Figure reproduced from reference [21] with permission from Elsevier.

It can be seen from Fig. 22(a) that the temperature dependence of cmc on T is again U-shaped, but the cmc values are in a very narrow range (cmc = 11.0 ± 1.5 mmol dm⁻³). With a less sensitive technique, one would probably not notice any difference. Similarly, for the Δ_M H ^θ its temperature dependence (= Δ M c p 0 ) the values are very similar. However, [N-C₁₂-2-mpy]⁺ was found to have the highest (absolute) value of Δ M c p 0 [21]. This could be attributed to the dehydration of the methyl group in ortho-position (Fig. 22(b)) and even to the dehydration of the pyridinium ring during micelle formation. In contrast, for other systems, it can be assumed that the alkyl chain is mainly dehydrated during the micelle formation process, while the pyridinium ring remains completely in contact with water.

Fig. 22:

Temperature dependence of (a) cmc for investigated SAILs in water. Solid lines represent the corresponding polynomial fits, cmc = A + BT + CT ² (the coefficients are given in [21]); (b) Δ_M H ^θ for investigated SAILs in water. Figure reproduced from reference [21] with permission from Elsevier.

Obviously, the isomerism of the cation has less influence overall micellization process than the isomerism of the counterions.