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One of the Distinctive Properties of Ionic Liquids over Molecular Solvents and Inorganic Salts: Enhanced Basicity Stemming from the Electrostatic Environment and “Free” Microstructure Qiwei Yang,† Huabin Xing,*,† Zongbi Bao,† Baogen Su,† Zhiguo Zhang,† Yiwen Yang,† Sheng Dai,‡,§ and Qilong Ren*,† †

Key Laboratory of Biomass Chemical Engineering of Ministry of Education, Department of Chemical and Biological Engineering, Zhejiang University, Hangzhou 310027, China ‡ Chemical Sciences Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37831, United States § Department of Chemistry, University of Tennessee, Knoxville, Tennessee 37966, United States S Supporting Information *

ABSTRACT: The basicity of ionic liquids (ILs) underlies many important ILbased processes including the dissolution and conversion of cellulose, the capture of CO2, and metal catalysis. In this work, we have disclosed the nature of the basicity of ILs, i.e., the difference between the basicity of IL and the basicity of the molecular solvent and inorganic salt, through a quantitative electrostatic and electronic analysis of the molecular surface for the first time. The results reveal one of the distinctive properties of ILs (enhanced basicity over molecular solvents and inorganic salts with the same basic site) stemming from their special electrostatic environment and microstructure. The enhancement is significant, from either the electrostatic aspect or the covalent aspect of basicity. The peculiar electrostatic environment of ILs leads to stronger basicity than similar molecular solvents, and the relatively freer microstructure of ILs contributes to the enhancement of basicity over their inorganic analogues. These results are highly instructive for better understanding the unique value of ILs and designing novel ILs to improve the efficiency of basicity-related processes.

1. INTRODUCTION Understanding the acidity and basicity of chemicals is one of the major topics in chemistry. Intermolecular interactions arising from the acidity and basicity of molecules are the foundation of various chemical processes and can play an important role in the development of novel chemical materials. One fascinating type of material, ionic liquids (ILs), has attracted quite a bit of attention in recent years as novel media for reaction and separation.1−5 The unique efficiency of ILs in several impressive applications relates to their basicity. ILs have shown great performance in the dissolution of cellulose,6−11 lignin,12,13 proteins,14,15 bioactive compounds,16,17 and even metallic compounds,18 providing a new platform for processes such as biomass conversion and metal catalysis. The most important dissolution mechanism relies on the hydrogenbonding interaction between the basic anions of ILs and solutes, and the solubility of solutes often shows obvious dependence on the basicity of ILs. In another intriguing application of ILs, the capture of CO2 and the separation of other gases,19−28 the interaction between the basic group of the IL and the acidic atom of the gas molecule is critical for improving the absorption capacity and separation selectivity. Furthermore, in chemical reactions using ILs as solvents, the basicity of ILs can be crucial to the activation of the reactant or © 2014 American Chemical Society

the stabilization of the intermediate and, further, to the reaction efficiency.29,30 To some extent, it is simply basicity that takes the usefulness of ILs beyond the concept of a nonvolatile solvent. Therefore, a promising route toward a better understanding of why the performance of ILs is superior to that of traditional molecular solvents or inorganic salts in chemical processes is to illustrate the nature of the basicity of ILs. Current studies on the basicity of ILs mostly focus on the determination of the experimental values of basicity. Especially, the Kamlet−Taft solvatochromic parameter, hydrogen-bonding basicity β, has been reported for many different ILs and ILsolvent mixtures together with other solvent properties.31−34 The NMR and chromatographic methods have also been applied to determine the hydrogen-bonding basicity of ILs.35,36 Other works tried to estimate or predict the hydrogen-bonding basicity of ILs by correlating the experimental β values with different theoretical parameters. Niedermeyer et al. correlated the β values of ILs with several computed IL-solute interaction parameters,37 and we have investigated the relationship Received: January 23, 2014 Revised: March 12, 2014 Published: March 12, 2014 3682

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Kamlet−Taft solvatochromic experiments were 4-nitroaniline and N,N-diethyl-4-nitroaniline, purchased from Sigma-Aldrich and Oakwood Products, Inc., respectively. Other chemicals were all commercially obtained and used without further purification. [Emim]Gly was synthesized and purified according to the literature,48 using 1-ethyl-3-methylimidazolium bromide ([emim]Br) and glycine as the starting material. First, the bromide anion of [emim]Br was exchanged with a hydroxyl anion to produce 1-ethyl-3-methylimidazolium hydroxide ([emim]OH), by eluting the [emim]Br aqueous solution through strongly basic anion exchange resin (hydroxyl ion type). The AgNO3 test was performed to ensure the complete exchange. Second, the aqueous solution of [emim]OH was added dropwise into excessive aqueous solution of glycine, and then the mixture was stirred at the room temperature for 24 h. The resulting solution was evaporated under vacuum at 60 °C to remove the water, after which the crude product was dissolved in acetonitrile/methanol solution (90/10, v/v) and filtered to remove the excessive glycine. The final product was dried under vacuum at 55 °C for about 48 h, and the purity was confirmed by NMR analysis on a Bruker DMX-500 spectrometer. 1H NMR (DMSO-d6, δ/ppm): 1.36 (t, 3H), 2.78 (s, 2H), 3.88 (s, 3H), 4.27 (q, 2H), 7.86 (s, 1H), 7.96 (s, 1H), 10.18 (s, 1H). Kamlet−Taft Solvatochromic Experiments. Aliquots of probe solution in dichloromethane with a certain concentration were added into a vessel. After the evaporation of dichloromethane, some IL was added. The obtained probe-IL mixture was stirred thoroughly, and then was transferred into a cuvette and placed in the sample cell of a UV−visible spectrophotometer (Shimadzu UV-2550). After that, the maximum absorption wavelength of the sample was measured. The temperature of the cell was maintained at 308.15 ± 0.1 K by an external temperature controller (Shimadzu, S-1700). The average of five repeated scans was taken as the final value of maximum absorption wavelength, with an uncertainty of 0.5 nm. The Kamlet−Taft hydrogen-bonding basicity β was calculated according to eq 1:49

between the basicity strength of ILs and the cation−anion interaction energy.38 Despite the achievements of the published works, the understanding of the nature of ILs’ basicity at the electronic level, which features IL as unique materials, is still very limited. Especially, the difference between the basicity of ILs and the basicity of inorganic salts and molecule solvents as well as the underlying mechanism is unclear, which seriously hinders the understanding and application of ILs. Herein, we have disclosed the nature of the basicity of ILs, i.e., the difference between the basicity of IL and the basicity of molecular solvent and inorganic salt, through a quantitative electrostatic and electronic analysis of molecular surface for the first time. Two intrinsic properties of ILs themselves, the minimum surface electrostatic potential (Vs,m) and the minimum average local ionization energy on the surface (Is̅ ,m) were employed as powerful tools to directly elucidate the basicity of ILs at the electronic level. The results, combined with experimental evidence, have demonstrated that ILs have enhanced basicity over molecular solvents and inorganic salts with the same basic site, stemming from the electrostatic environment and “free” microstructure of ILs.

2. COMPUTATIONAL METHODS The geometry optimization of the studied molecules and ionic systems was carried out by the Gaussian 03 software,39 at the B3LYP/6-311++G(d,p) level of theory.40,41 All the optimized structures were confirmed to be minima on the potential energy surface via vibrational frequency analysis. The calculations on the surface electrostatic potential and the surface average local ionization energy were performed by WFA 1.0 program,42 using B3LYP/6-311++G(d,p) optimized geometries, wave functions and electrostatic potential cube files from Gaussian 03. According to the suggestion of Bader et al.,43 the molecular surface was defined to be the 0.001 au (atomic unit, equals electron/bohr−3) outer contour of its electronic density ρ(r). A visualization tool SurRender was used to obtain the figures of electrostatic potential and average local ionization energy mapped on the ρ(r) = 0.001 au contour.42 The gas-phase basicity is defined as the negative of the Gibbs free energy change (ΔG) for the protonation reaction B + H+ → BH+ at 298.15 K and 1 atm, where B refers to the investigated substance. The gas-phase enthalpy and entropy of proton in the standard conditions are 1.48 and 0.026 kcal mol−1,44 respectively, which produces a gas-phase free energy of proton of −6.28 kcal mol−1. The ΔG values of NaCl and [Bmim]Cl were corrected by the counterpoise method for the basis set superposition error.45

β = −0.357vNA + 0.369vDENA + 0.949

(1)

where νNA and νDENA are the maximum absorption wavenumbers of 4-nitroaniline and N,N-diethyl-4-nitroaniline, respectively.

4. RESULTS AND DISCUSSION When one pays attention to the basicity of chemicals from the microscopic view, two basic physical phenomena can be found. One is the electrostatic attraction of the basic site to the positively charged atom or group of another molecule. The other is the donation of electrons from the basic site to the orbitals of the second molecule.50−52 As a consequence, the concept of basicity can be roughly divided into two components: the electrostatic component and the covalent component, respectively. In other words, the electrostatic and covalent effects together contribute to the appearance and the strength of the basicity. This concept provides a bridge between the extrinsic base−acid interactions and the intrinsic properties of the basic molecule. Therefore, in this work, the basicity of ILs was investigated and compared with those of their analogues from both the electrostatic and covalent aspects for the first time.

3. EXPERIMENTAL SECTION Materials. The used ILs (>99%) were all purchased from Green Chemistry and Catalysis, LICP, CAS. (China) or Chengjie Chemical Co., Ltd. (Shanghai, China), except the amino acid IL 1-ethyl-3-methylimidazolium glycine ([Emim]Gly). All ILs were dried under vacuum at 50−100 °C for at least 24 h before use, and the water contents of ILs were determined by a Karl Fischer Titration. The water mole fraction of [Bmim]Cl was 0.02, and the water mole fractions of other ILs were 0.03−0.07. The latter values seem not very low; nevertheless, its effect on the experimental hydrogen-bonding basicity values should be very small, according to the previous literature.34,46,47 Lead chloride (99.99%) was purchased from Aladdin Reagent Co., Ltd. The indicator probes for the 3683

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kcal mol−1) is just about 30 kcal mol−1. In fact, we have obtained direct experimental evidence for the enhanced basicity of [Bmim]Cl over NaCl for the first time, using the optical probe method developed by Duffy et al. for the measurement of “optical basicity” of glasses and metallic compounds (see the Supporting Information for details).57 It is found that the ultraviolet spectra of the probe, Pb2+, in [Bmim]Cl revealed an absorption maximum at 284 nm due to the 1S0 → 3P1 transition of Pb2+, much larger than the corresponding wavelength in NaCl (273 nm).58,59 Because the Pb2+−Cl− interaction leads to a red shift of the 1S0 → 3P1 transition of free Pb2+, the larger maximum absorption wavelength of Pb2+ in [Bmim]Cl than that in NaCl clearly demonstrates that the basicity of [Bmim]Cl is indeed stronger than the basicity of NaCl. Additionally, we also investigated the difference in basicity between NaOH and OH−-based IL. The calculated gas-phase basicity of NaOH ion pair was 255.1 kcal mol−1. Although this value was higher than that of [Bmim]Cl ion pair, it was lower than that of the OH−based IL ion pair, tetrabutylphosphonium hydroxide ([P4444]OH) (284.1 kcal mol−1), indicating again that the basicity of ILs could be stronger than that of inorganic salts with the same anion. It is of great interest to elucidate why the basicity of ILs can be notably enhanced over that of common analogues, and the electrostatic and electronic analyses on the molecular surface provide a possible approach for the exploration. Generally, the most apparent characteristic of ILs may be the strong electrostatic environment within them, which is quite different from the nearly neutral environment within molecular solvents. Thus it is speculated that the electrostatic environment probably contributes to the enhancement of basicity. To validate this hypothesis, we tracked the structural transition from a neutral molecule to an ion pair, focusing on how the electrostatic and electronic characteristics of compounds changed (Figure 2). The selected molecule was alanine, which bears a primary amino group in the structure. For the alanine molecule, the calculated Vs,m and Is̅ ,m values were −32.4 kcal mol−1 and 7.56 eV, respectively, close to those of amines. Nevertheless, when this molecule became excessive in number of electrons as an alaninate anion, both the Vs,m and Is̅ ,m values became much lower. This suggests that the potential of the amino group to attract the acidic site of a solute molecule via electrostatic interaction and donate its electron pair was notably enhanced by the negative ionization, even though the ionization was induced by the change of the carboxyl group rather than by the amino group itself. Further, when the alaninate anion paired with an 1-ethyl-3-methylimidazolium cation, an ion pair of the IL 1-ethyl-3-methylimidazolium alaninate ([Emim]Ala) formed. It was observed that although the Vs,m and Is̅ ,m values near the amino group of the [Emim]Ala pair were higher than those in the isolated alaninate anion, due to the counteraction effect from the positive charge of cation and the cation−anion electrostatic interaction, the values of both were still much smaller than those of the alanine molecule. In fact, a difference can be found not only in the minimum value of the surface properties. From Figure 2, it can be seen that the domains near the amino group of [Emim]Ala show much more blue or green color than those of the alanine molecule, indicating that the surface electrostatic potential and the surface average local ionization energy near the amino group of [Emim]Ala are generally smaller than those of alanine. Therefore, it is reasonable to conclude that, compared with the case for the neutral molecule, the capability of electrostatic and covalent

Usually, when the types of basic group are similar, the strengths of the basicity of different molecules are also similar, unless there is a strong inductive or conjugative effect. However, surprisingly, when we gained insight into the basicity of ILs, it was observed that compared with a molecular solvent or an inorganic salt with a similar basic group, the ILs had basicities that were notably enhanced from either the electrostatic or the covalent aspect. Typically, a comparative study was carried out between the most representative basic solvents, amines, and the amino acid ILs, which bears an amino group at the anion, using Vs,m and Is̅ ,m to evaluate the intrinsic electrostatic character and covalent character of basicity,42,53,54 respectively (Figure 1a,b). For several common aliphatic

Figure 1. Comparison between aliphatic amines (yellow) and amino acid ILs (green) at the B3LYP/6-311++G(d,p) level: (a) minimum surface electrostatic potential, Vs,m; (b) minimum surface average local ionization energy, Is̅ ,m; (c) gas-phase basicity. Amines: 1, ethylamine; 2, diethylamine; 3, n-butylamine; 4, tert-butylamine; 5, triethylamine; 6, dimethylamine; 7, cyclohexylamine; 8, ethylenediamine. ILs: 9, [Emim]Gly; 10, [Emim]Ala; 11, [Emim]Met; 12, [Emim]Val. Emim = 1-ethyl-3-methylimidazolium ion. Gly = glycinate anion. Ala = alaninate anion. Met = methioninate anion. Val = valinate anion.

amines, the values of Vs,m or Is̅ ,m were similar, implying similar basicities. Nevertheless, for amino acid ILs, the values of Vs,m or Is̅ ,m near the amino group became much lower, indicating a much stronger electrostatic or covalent basicity for the amino group in the IL than for those in the molecular amines. The lowest Vs,m and Is̅ ,m values of the studied amines were −37.7 kcal mol−1 (n-butylamine) and 6.64 eV (triethylamine), respectively, but the Vs,m and Is̅ ,m values of the studied amino acid ILs were all lower than −52.5 kcal mol−1 and 6.20 eV. A similar result was obtained for the comparison of an IL with an inorganic salt. The Vs,m and Is̅ ,m values of a NaCl ion pair were −50.2 kcal mol−1 and 6.75 eV, respectively. But the Vs,m and Is̅ ,m values of a [Bmim]Cl ion pair were −62.8 kcal mol−1 and 5.74 eV, notably lower than those of NaCl, indicating that the electrostatic or covalent basicity of [Bmim]Cl was stronger than that of NaCl. Arising from the intrinsically increased electrostatic and covalent basicity, the overall basicityor rather the overall apparent capacity of ILs to interact with an acidic moleculewas enhanced. We calculated the absolute value of the Gibbs free energy change of the protonation reaction, also known as the gas-phase basicity,55,56 for the compounds mentioned (Figure 1c). It was found that the gasphase basicities of the aliphatic amines were all about 210 kcal mol−1, but the basicities of the amino acid ILs showed values much higher than 245 kcal mol−1. The gas-phase basicity of the NaCl ion pair was 202.7 kcal mol−1, but that of the [Bmim]Cl ion pair reached up to 236.9 kcal mol−1. These increases are significant, because the difference between the gas-phase basicity of most aliphatic amines and that of acetone (187.0 3684

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Figure 3. Minimum surface electrostatic potential Vs,m (green) and the minimum surface average local ionization energy Is̅ ,m (red) of the NaCl ion pair against the Na···Cl distance, computed at the B3LYP/6-311+ +G(d,p) level.

pair, a longer cation−anion distance could be expected for the [Bmim]Cl ion pair than for the NaCl. Interestingly, the calculated average distance between the chloride anion and each atom of the [Bmim]+ cation was 4.32 Å, and the distance between the chloride anion and the mass center of the cation was 2.91 Å, both much larger than the equilibrium Na···Cl distance in a NaCl ion pair. This result indicates that the longer cation−anion distance in [Bmim]Cl is probably an important reason for its smaller Vs,m and Is̅ ,m values, and further, its stronger basicity. In fact, the phenomenon that [Bmim]Cl has a larger cation−anion distance, along with lower Vs,m and Is̅ ,m values, than NaCl is found not only for the simplest system (an ion pair) but also for more complex systems, such as the ionic dimer and trimer (Figures 4 and 5 and Figure S2 (Supporting

Figure 2. Change of molecular surface properties during the structural transition from alanine to alaninate anion to 1-ethyl-3-methylimidazolium alaninate ([Emim]Ala), computed at the B3LYP/6-311+ +G(d,p) level: (a) electrostatic potential at the 0.001 au contour of the electron density and its minimum Vs,m near the amino group; (b) average local ionization energy at the 0.001 au contour of the electron density and its minimum Is̅ ,m near the amino group. Color ranges for electrostatic potential, in kcal mol−1: blue < −30.0 < green ≤ 5.2 < yellow < 19.6 < red. Color ranges for average local ionization energy, in eV: blue < 8.68 < green < 11.23 < yellow < 13.79 < red. The positions of Vs,m and Is̅ ,m are denoted by the red arrow and cyan dot on the surface.

interactions of the amino group as a basic site can be greatly enhanced by the electrostatic environment within ILs when the amino group is located at the anion, leading to the enhanced basicity. This discussion has shed light on the difference between ILs and molecular solvents, but it cannot elucidate the difference between an IL and an inorganic salt that also has a strong electrostatic environment. Thus there should be other reasons. Apparently, inorganic salts are solid but ILs are liquid near room temperature. Thus the reason may lie in the spatial structure of the compounds. One important parameter for the spatial structure of an ionic compound is the distance between the cation and anion. To give insight into the effect of cation− anion distance on basicity, we investigated the Vs,m and Is̅ ,m of a NaCl ion pair at different Na···Cl distances. Much to our surprise, the Na···Cl distance affects the values of Vs,m and Is̅ ,m significantly. As shown in Figure 3, when the Na···Cl distance increased from a value smaller than the equilibrium distance (2.38 Å), both the Vs,m and Is̅ ,m values decreased rapidly, in a nearly linear fashion. Thus because the Vs,m and Is̅ ,m values of a [Bmim]Cl ion pair were notably lower than those of a NaCl ion

Figure 4. Optimized geometries of the ion pair, dimer, and trimer of [Bmim]Cl at the B3LYP/6-311++G(d,p) level.

Information) and detailed information of Vs,m/Is̅ ,m calculation method in the Supporting Information). In the NaCl dimer (Figure S2, Supporting Information), the distance between each Cl− and each Na+ is 2.55 Å. In the [Bmim]Cl dimer (Figure 4), the average distance between each Cl− and the atoms (mass centers) of the two cations was 4.86 Å (3.96 Å). In the NaCl trimer, the average distance between the typical chloride anion (Cl− 2 in Figure S2, Supporting Information, 3685

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Figure 5. Minimum surface electrostatic potential Vs,m (left) and the minimum surface average local ionization energy Is̅ ,m (right) of the ion pair, dimer, and trimer of NaCl (yellow) and [Bmim]Cl (green), computed at the B3LYP/6-311++G(d,p) level.

Figure 6. (a) Plot of the experimental Kamlet−Taft hydrogen-bonding basicity β of different ILs against their minimum surface electrostatic potential Vs,m. The solid line stands for correlation with a linear equation. (b) Plot of experimental β values (βexp) against those calculated with a multiple linear regression model based on Vs,m and the minimum surface average local ionization energy Is̅ ,m (βcal). The solid line is the diagonal line. 1, [Bmim]C3H7CO2; 2, [Emim]Gly; 3, [Bmim]CH3CO2; 4, [Bmim]Cl; 5, [Bmim]CF3CO2; 6, [Bmim]CH3SO4; 7, [Bmim]HSO4; 8, [Bmim]NO3; 9, [Bmim](CN)2N; 10, [Bmim]CF3SO3; 11, [Bmim]BF4; 12, [Bmmim]BF4; 13, [BPy]BF4; 14, [Omim]NTf2; 15, [HOEmim]NTf2; 16, [Hmim]NTf2; 17, [Bmim]NTf2; 18, [Bmim]PF6. [Bmim] = 1-butyl-3-methylimidazolium ion. [Emim] = 1-ethyl-3-methylimidazolium ion. [Omim] = 1octyl-3-methylimidazolium ion. [Hmim] = 1-hexyl-3-methylimidazolium ion. [Bmmim] = 1-butyl-2,3-dimethylimidazolium ion. [HOEmim] = 1-ethoxyl-3-methylimidazolium ion. [BPy] = N-butylpyridinium ion. Gly = glycinate anion. NTf2 = bis(trifluoromethylsulfonyl) imide anion. The experimental β values of ILs 1, 2, 4, 6, 9, 10, 11, 12, 14, 15, 16, 17, and 18 were taken from refs 33, 34, and 61−65. Vs,m and Is̅ ,m were calculated at the B3LYP/6-311+ +G(d,p) level.

which is surrounded by three cations) and the three cations is 2.74 Å. In the [Bmim]Cl trimer, the average distance between the typical chloride anion (Cl− 42 in Figure 4, which is surrounded by three cations) and each atom (mass center) of the surrounding three cations was 5.73 Å (4.94 Å). Along with the larger cation−anion distance in [Bmim]Cl, the Vs,m and Is̅ ,m values of the [Bmim]Cl dimer and trimer were all higher than those of NaCl analogues, similar to the case of ion pair. More importantly, as shown in Figure 5, although the Vs,m/Is̅ ,m values of NaCl and [Bmim]Cl both increased when the studied structure varied from the ion pair to the dimer and further trimer, the increase in [Bmim]Cl systems was slower than in NaCl systems, so that the difference of Vs,m/Is̅ ,m between NaCl and [Bmim]Cl became more obvious. The Vs,m difference between a NaCl trimer and a [Bmim]Cl trimer was 23.4 kcal mol−1, and the Is̅ ,m difference was 2.12 eV, both higher than the corresponding differences between the two types of ion pairs (12.4 kcal mol−1, 1.00 eV). In other words, the enhanced basicity of [Bmim]Cl over NaCl was more obvious when the investigated system was more complex. The different cation−anion distance is just one reflection of the differential spatial structures of NaCl and [Bmim]Cl. On the whole, the microscopic structure of NaCl is quite compact and orderlyone ion is tightly surrounded by six ions of the opposite charge. However, the structure of [Bmim]Cl is much less compact and more flexible, even in the crystal state.60 As a consequence, compared with NaCl, the anion in [Bmim]Cl will be more free, and its electrostatic field and electrons are less affected by the cations. The Vs,m and Is̅ ,m of a perfectly free chloride anion are −137.4 kcal mol−1 and 1.09 eV, respectively, much lower than those of NaCl and [Bmim]Cl. Thus to a large extent, it can be argued that it is the freer anion in [Bmim]Cl that makes the Vs,m and Is̅ ,m of [Bmim]Cl lower than those of its inorganic analogue, NaCl, and then brings on the enhanced basicity. The above studies demonstrated the difference in basicity strength between ILs and the molecular or ionic analogues, as well as the underlying mechanism. After that, further interest was attached to the basicity variance of different ILs. We investigated the basicity of different ILs from a statistical viewpoint, using the intrinsic parameters Vs,m and Is̅ ,m as independent variables. In the case of molecular solvents, the hydrogen-bonding basicity of different solvents can correlate well with the Vs,m values of the solvent molecules.54 However, when we tried to correlate the Kamlet−Taft hydrogen-bonding basicity β of different ILs with their Vs,m values, the correlation was found to be poor (Figure 6a), whereas a much better correlation was obtained when both the Vs,m and the Is̅ ,m parameters were employed as fitting variables in a multiple

linear regression (Figure 6b). The regression model to calculate β was given in eq 2: βcal = 1.781 − 0.011Vs,m − 0.206Is,m ̅ (n = 18, R2 = 0.73, F = 20.4)

(2)

where Vs,m and Is̅ ,m were the minimum surface electrostatic potential and the minimum average local ionization energy on the surface, respectively. Because the Is̅ ,m refers to the covalent character of basicity, the results probably suggest that, from a statistical viewpoint, the hydrogen-bonding basicity of ILs has a more covalent character than molecular solvents. This phenomenon is probably also caused by the peculiar electrostatic environment of ILs, like the enhancement of the basicity strength.

5. CONCLUSION In summary, we have shown that, compared with molecular solvents or inorganic salts with a similar basic group, the ILs have basicities mainly from the anion that are notably enhanced from either the electrostatic aspect or the covalent aspect of basicity. The peculiar electrostatic environment of ILs leads to stronger basicity than for similar molecular solvents, and the relatively freer microstructure of ILs contributes to the enhancement of basicity compared with their inorganic analogues. These results not only reveal the nature of the basicity of ILs, which is of great importance for a better understanding of the unique value of ILs, but also are highly instructive for the design of ILs and even other ionic materials. They show a way to enhance the basicity of ILs and improve 3686

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the performance of IL-mediated processes by making the anion of the IL more free.

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ASSOCIATED CONTENT

S Supporting Information *

Further information about the experimental methods and results of the ultraviolet absorption spectrum of [Bmim]Cl:Pb2+. Details about the Vs,m and Is̅ ,m calculation for the ion pair, dimer, and trimer of NaCl and [Bmim]Cl. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*H. Xing: e-mail, [email protected]. *Q. Ren: e-mail, [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Natural Science Foundation of China (21222601, 21106127), the Zhejiang Provincial Natural Science Foundation of China (LR13B060001), and the Program for New Century Excellent Talents in University of China (NCET-13-0524). S.D. was supported by the Division of Chemical Sciences, Office of Basic Energy Sciences, U.S. Department of Energy.

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REFERENCES

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