Exploring physicochemical properties of the nanostructured Tunable Aryl Alkyl Ionic Liquids (TAAILs)

Journal of Molecular Liquids 209 (2015) 14–24

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Journal of Molecular Liquids journal homepage: www.elsevier.com/locate/molliq

Exploring physicochemical properties of the nanostructured Tunable Aryl Alkyl Ionic Liquids (TAAILs) Hossein Roohi ⁎, Khatereh Ghauri Department of Chemistry, Faculty of Science, University of Guilan, Rasht, Iran

a r t i c l e

i n f o

Article history: Received 8 January 2015 Received in revised form 24 March 2015 Accepted 1 May 2015 Available online 20 May 2015 Keywords: TAAIL Substituent effect M06-2X functional Binding energy Physicochemical property

a b s t r a c t The ionic liquids, termed TAAILs (Tunable Aryl Alkyl Ionic Liquids), are clearly different from the standard imidazolium based ionic liquids. In this work, a systematic and comprehensive study of the physicochemical properties of nanostructured TAAILs [X-PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2) was carried out by using M06-2X functional in conjunction with the 6-311++G(d,p) basis set. The binding energy, Gibbs free interaction energies, enthalpy of the formation of cations, structural parameters, topological properties of electron density, natural charges and charge transfer values were calculated. The results demonstrate that the strength of interaction between cations and anion increases with increase in the electron-accepting power of the substituents. The order of changes in electrochemical stability, melting point, conductivity, surface tension and critical-point temperature of the studied IPs were estimated. © 2015 Elsevier B.V. All rights reserved.

1. Introduction As novel and green class of chemical compounds, ionic liquids (ILs) with a wide range of applications have gained increasing interest. They have physicochemical properties including low melting points, negligible vapor pressure, extremely low volatility, non-flammability, unusual solvation, controllable hydrophobicity, resistant to oxidation and extraordinarily high thermal/chemical/electrochemical stability, particularly in the presence of air and moisture, etc. During the last 15 years, ionic liquids have been shown to be very promising “green solvents” with several advantages compared to traditional organic solvents [1–7]. Recently, Ahrens et al. were introduced a new class of imidazolium based ILs [8,9]. These new ionic liquids, termed Tunable Aryl Alkyl Ionic Liquids (TAAILs), are clearly different from all other ionic liquids and constitute a new generation of ionic liquids with a high degree of flexibility. Standard imidazolium based ionic liquids carry two alkyl chains (Csp3), which are limited to inductive effects (+I, −I). In contrast, the TAAIL concept introduces mesomeric and inductive effects (+ M, − M) by having a substituted phenyl ring (Csp2) replaced by one of the alkyl substituents in dialkyl substituted imidazolium based ILs. Combination of sp3 alkyl and sp2 aryl substituents at those nitrogen atoms of the imidazolium ring, allows a far greater variation of the ionic liquids characteristics than imagined compared to the usual ILs. Similar to the standard ILs, the anions can be chosen without restrictions [10]. TAAILs are not restricted to van der Waals interactions; they also ⁎ Corresponding author. E-mail addresses: [email protected], [email protected] (H. Roohi).

http://dx.doi.org/10.1016/j.molliq.2015.05.001 0167-7322/© 2015 Elsevier B.V. All rights reserved.

allow π–π interactions, which will be important for applications in the separation of compounds as well as for the stabilization of catalytically active metals [8]. The nature of the intermolecular forces in all materials and especially in different ILs mainly controls their physical properties. The physicochemical properties of ILs can be finely tuned by slight structural changes of the corresponding cations and anions [11,12]. To choose or design the suitable ILs for a variety of applications, it is necessary to understand their structures and interactions in various IL systems. It is, however, time-consuming and cost-intensive to study the physical properties of all ILs by experiments. Thus, study of the physical properties of ILs by computer simulation is an efficient and necessary approach. It is well known that both the proton-donating and protonaccepting groups on the para position phenyl ring change the electronic properties of substituted molecules [13,14]. Similarly, the interaction between the anion and cation can be obviously influenced by substituents at aryl ring of ILs and therefore these substituents have a significant effect on the properties of the resulting ILs. In the present work, we explore the effect of different substituents (electron-donor and electronacceptor groups) on the phenyl ring of the cation on the nanostructural parameters and related physicochemical properties of TAAILs with different cations and anion BF− 4 . Many methods have been developed to explore the anion–cation interactions of ILs, including experimental [15–23] and theoretical [24–35] methods. To the best our knowledge٫ interaction between [para-X-phenyl methyl imidazolium]+ cations and anion [BF4]− has not been characterized and no detailed studies exist on the influence of different substituents in the para position of these cations on hydrogen binding strength between their constituents. The substituents selected for this study cover almost the whole set of

H. Roohi, K. Ghauri / Journal of Molecular Liquids 209 (2015) 14–24

classical substituents: from the strongly electron-attracting (NO2٫ CN٫ CHO and F) to typical electron-donating (NH2٫ OH٫ OCH3 and CH3) ones. In this work, we have calculated the binding energies, geometrical parameters, topological properties, enthalpy and Gibbs free energy of formation of cations in the gaseous phase. In addition, we have analyzed the nature of intermolecular interactions in TAAILs by natural bond orbital (NBO) [36] and Badger's quantum theory of atoms in molecules (QTAIM) [37].

15

[56]. The calculations were carried out with Gaussian 03 [57] and GAMESS [58] program packages. To examine the nature of interactions, the electronic properties for stationary points were calculated by using natural bond orbital (NBO) analysis. The NBO analysis was carried out on the M06-2X/6311++G(d,p) wave function using version 3.1 of NBO program [59]. Topological properties of electron charge density [electron density, ρ(r), Laplacian of electron density, ∇2ρ (r), and electronic energy density, H(r)] were calculated using Bader's theory [60] at M06-2X /6311++G(d,p) level of theory by the AIM2000 program package [61].

2. Computational details Density functional theory (DFT) is wildly used in the past decade to study the energetic properties of ionic liquids [38–44], and the performance of a range of DFT functionals has been recently analyzed to identify the most adequate in providing ionic liquids binding energies [45–51]. Here, we have performed the DFT calculations using M06-2X [52,53] method in conjunction with the 6-311++G(d,p) basis set [54, 55]. All of initial structures were fully optimized with above method in gas phase. The calculated vibrational frequencies have been used to characterize stationary points and calculation of zero-point vibrational energy (ZPVE). The vibrational frequencies were calculated without scaling. For the calculation of binding energies, basis set superposition errors (BSSE) were considered using the counterpoise method (CP)

[H–PhMIM][BF4]

[OH–PhMIM][BF4]

[CHO–PhMIM][BF4]

3. Results and discussion A series of ILs with various cations and BF− 4 anion are investigated. The cations include para-X-phenyl methyl imidazolium (X = NH2, OH, OCH3, CH3, H, F, CHO, CN, NO2). The various sites around the substituted cations are evaluated for bond formation by BF− 4 anion. The optimized structures of the most stable ion pairs (IPs) from interaction between various cations and BF− 4 anion are depicted in Fig. 1. In these ion pairs, F atoms of BF− 4 anion act as proton acceptors and C–H bonds of the substituted cations as proton donors. Generally, a structural measure of H-bonding interaction is the distance between the proton on the donor group and the acceptor atom to be less than the sum of their

[Me–PhMIM][BF4]

[NH2–PhMIM][BF4]

[CN–PhMIM][BF4]

[OMe–PhMIM][BF4]

[F–PhMIM][BF4]

[NO2–PhMIM][BF4]

Fig. 1. Optimized structures of [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN, NO2) ion pairs at the M06-2X/6-311++G(d,p) level of theory.

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H. Roohi, K. Ghauri / Journal of Molecular Liquids 209 (2015) 14–24

van der Waals radii. Herein, the van der Waals distance for F⋯H is 2.670 Å [62]. Accordingly, there are four F⋯H H-bonds in the [XPhMIM][BF4] IPs as shown by dashed lines in Fig. 1. One H-bond is formed between the F atoms and C2–H6 atom of the ring and three of them are formed between F atoms and H atoms of the Me (two Hbonds) and phenyl groups (one H-bond). For NO2 and CN substituted IPs, four H-bonds are established between the F and H atoms of the Me (two H-bonds) and phenyl groups (two H-bonds). 3.1. Binding energy The calculated electronic and Gibbs Free binding energies of the ion pairs (IPs) at M06-2X level of theory are listed in Table 1. For the calculation of binding energies, basis set superposition errors (BSSE) and zero-point vibrational energies (ZPVE) were considered. The interaction energy (BE = −ΔEele) and BSSE corrected interaction energies (ΔEBSSE ele ) of the IPs were calculated according to the equations given below: ΔEele ¼ EIP −ðEX−Cation þ EAnion Þ

ð1Þ

ΔE0 ¼ ΔEele þ ΔZPVE

ð2Þ

ΔEBSSE ¼ ΔE0 þ BSSE: 0

ð3Þ

According to the data given in Table 1, the electronic binding energies are relatively large, ranging from 86.9 to 96.5 kcal/mol at M062X/6-311++G(d,p) level of theory. The Gibbs Free bind energies (GFBE = −ΔG) are ranged from 71.8 to 81.2 kcal/mol. The change of the electrostatic strength due to the charge delocalization in aromatic ring is reflected on the relative values of BEs. As can be seen in Table 1, aromatic-based ILs having the electron-donating substituents present lower interaction strength compared to the electron-accepting ones. The quantum chemical studies have shown that the substituent constants (Hammett's constants) [63] correlate with several electronic or physicochemical properties. According to the Hammett's interpretation, the σp values represent the sum of the total electronic effects (resonance plus inductive/field effects) [64]; positive values are associated with electron-withdrawing substituents, whereas negative values correspond to electron-donating ones. The linear correlations (R2 = 0.95) between BEs and Hammett's constants σp at M06-2X/6-311++G(d,p) level of theory are shown in Fig. 2. It can be seen that the values of BEs for ion pairs having electronaccepting substituents are greater than those of electron donating ones. In other words, the strong electron-donating/accepting group of the substituents decreases/increases the H-bonding strength. Electron-accepting substituents strengthen interaction between cation and anion in ion pairs and in turn increase the interaction energy. Based on the BSSE and ZPVE corrected BEs, strength of the inter-ionic Table 1 Interaction energies (kcal/mol) calculated for most stable forms of the [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2) ion pairs at M06-2X/6-311++G(d,p) level. Method

a b c d

X

BSSE

ΔZPVE

ΔEele

a

b

NH2 OH OCH3 CH3 H F CHO CN NO2

1.6 1.6 1.6 1.6 1.6 1.6 1.6 2.1 2.1

1.5 1.1 1.2 1.8 1.4 1.2 1.5 1.4 1.1

−86.9 −88.8 −87.9 −88.6 −89.7 −91.8 −93.3 −95.7 −96.5

−85.3 −87.7 −86.7 −86.8 −88.3 −90.6 −91.8 −94.2 −95.4

−85.3 −87.2 −86.3 −87.0 −88.1 −90.1 −91.7 −93.5 −94.4

ΔE0 = ΔEele + ΔZPVE. ΔEBSSE ele = ΔEele + BSSE. ΔEBSSE = ΔE0 + BSSE. 0 BSSE and ZPVE corrected Gibbs Free energy.

ΔE0

ΔEBSSE ele

c

ΔEBSSE 0

d

−83.7 −86.1 −85.1 −85.2 −86.7 −88.9 −90.2 −92.1 −93.2

−71.8 −74.9 −73.7 −72.2 −75.4 −77.5 −78.2 −80.3 −82.1

ΔG°g

Fig. 2. Correlation between the bonding energy (BE) values and Hammett's constants (σp) of [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN, NO2) ion pairs.

interactions in the ILs decreases in the following order: [NO 2 – PhMIM][BF 4] N [CN–PhMIM][BF4 ] N [CHO–PhMIM][BF 4] N [F– PhMIM][BF 4] N [PhMIM][BF4 ] N [CH3–PhMIM][BF4 ] N [OH– PhMIM][BF4] N [OCH3–PhMIM][BF4] N [NH2–PhMIM][BF4]. There is a simple interpretation for stability order observed in these type ILs. Regarding the total electronic effects of substituents on the ionic interactions, electron-donating substituents on the aromatic ring decreases the overall net charge on the cation ring through electron release, when compared to the electron-accepting substituents, leading thus to unfavorable anion–cation interactions. 3.1.1. Correlation between BE and the experimental properties The melting point of ILs can be correlated with the structure, composition and intermolecular forces. The melting point determines the lower limit of the liquidity and also the thermal stability of ILs. There is a correlation between BE and physical properties of ILs for dialkylimidazolium ILs. The reported results [65] show that the melting point of [Alkylmim+][BF− 4 ] ILs increases as the BE increases. The influence of the substituents in the para position of aryl ring on the melting points of the TAAILs has been investigated by Strassner et al. [66]. They showed that electron-withdrawing groups (NO2, halogens) tend to higher melting points than electron-donating (Me, OMe, OEt) substituents in the paraposition of the phenyl ring, which is in good agreement with the greater BE obtained for ILs having electron-withdrawing groups. Thus, based on relative stability of ion pairs, it can be predicted that the melting point of studied ILs decreases in the following order: [NO2–PhMIM][BF4] N [CN–PhMIM][BF4] N [CHO–PhMIM][BF4] N [F–PhMIM][BF4] N [PhMIM][BF4] N [CH3–PhMIM][BF4] N [OH–PhMIM][BF4] N [OCH3– PhMIM][BF4] N [NH2–PhMIM][BF4]. Ionic association reduces the number of ions in solution and, in turn, reduces the electrical conductivity of the solution. It is predicted that the order of conductivity of ILs is in contrast to the order of melting points of ILs. Based on calculated BEs, it is estimated that the conductivity decreases more with electron-drawing substituents and increases with electron-donating ones. In addition, it has been shown that surface tension property follows a trend analogous to the relative interaction energies determined by mass spectrometry as well as calculated by DFT methods [45]. From the comparison of BEs, it seems reasonable to assume that the surface tension of TAAILs having electron accepting substituents is greater than those of electron donating ones. The critical temperature is basically a measure of the strength of intermolecular interactions. Thus, it is predicted that the increase in binding energy is accompanied by an increase in the critical-point temperature. From the comparison of BEs, it can be estimated that the critical-point temperature of TAAILs having electron accepting substituents to be greater than those of electron donating ones.

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Table 2 The optimized geometrical parameters (distances in Å and angles in °) at M06-2X/6-311++G(d,p) level of theory. Parameter

Cation X=H

C16–C17 C17–C18 C16–N1 N1–C5 C4–N3 N3–C2 C2–N1 N3–C7 C17–H22 C5–H23 C4–H24 C2–H6 C7–H8 C7–H9 C7–H10 B11–F12 B11–F13 B11–F14 B11–F15 F12⋯H22 F12⋯H6 F14⋯H6 F14⋯H8 F13⋯H8 F13⋯H9 F13⋯C7 F13⋯N3 F12H22C17 F12H6C2 F14H6C2 F14H8C7 B11F13N3 B11F13C7 C17C16N1C2 Parameter

C16–C17 C17–C18 C16–N1 N1–C5 C4–N3 N3–C2 C2–N1 N3–C7 C17–H22 C5–H23 C4–H24 C2–H6 C7–H8 C7–H9 C7–H10 B11–F12 B11–F13 B11–F14 B11–F15 F12⋯H22 F12⋯H6 F13⋯H8 F13⋯H9 F14⋯H22 F14⋯H6 F14⋯H8 F13⋯C7 F13⋯N3 F12⋯N1 F14⋯C2 F12H22C17 F12H6C2 F14H22C17 F14H6C2 F14H8C7

CH3

1.389 1.390 1.442 1.380 1.378 1.330 1.334 1.467 1.084 1.077 1.077 1.079 1.088 1.089 1.088

−51.2

OCH3

1.387 1.390 1.441 1.379 1.378 1.331 1.333 1.466 1.084 1.077 1.077 1.079 1.088 1.089 1.089

OH

1.394 1.379 1.440 1.379 1.378 1.331 1.333 1.466 1.084 1.077 1.077 1.079 1.088 1.089 1.089

−51.2

NH2 1.392 1.383 1.440 1.380 1.378 1.331 1.334 1.466 1.084 1.077 1.077 1.079 1.088 1.089 1.089

−52.8

−53.9

Anion

Ion pair

[BF− 4 ]

X=H

CH3

OCH3

OH

NH2

1.4088 1.4088 1.4088 1.4088

1.388 1.391 1.437 1.382 1.377 1.327 1.338 1.467 1.085 1.076 1.076 1.082 1.087 1.088 1.089 1.418 1.413 1.430 1.369 2.005 2.379 2.080 2.148 2.623 2.441 2.787 2.875 162.3 112.4 130.3 137.9 107.1 109.2 −44.4

1.390 1.388 1.437 1.381 1.377 1.328 1.337 1.467 1.085 1.076 1.076 1.082 1.087 1.088 1.089 1.417 1.415 1.429 1.369 2.026 2.363 2.078 2.140 2.604 2.459 2.785 2.872 161.2 116.1 131.6 139.3 108.3 109.5 −44.1

1.394 1.382 1.437 1.381 1.376 1.328 1.337 1.466 1.085 1.076 1.076 1.082 1.087 1.088 1.090 1.417 1.414 1.430 1.369 2.017 2.388 2.084 2.143 2.610 2.446 2.781 2.865 159.6 111.5 129.7 138.4 107.1 109.2 −47.1

1.391 1.385 1.436 1.381 1.376 1.328 1.337 1.467 1.085 1.076 1.076 1.082 1.087 1.088 1.089 1.418 1.414 1.430 1.369 2.017 2.373 2.080 2.142 2.608 2.451 2.781 2.861 159.8 113.2 130.5 138.7 107.7 109.4 −46.4

1.390 1.385 1.436 1.381 1.376 1.329 1.337 1.466 1.085 1.076 1.076 1.082 1.087 1.088 1.090 1.417 1.414 1.429 1.370 2.027 2.379 2.086 2.142 2.604 2.457 2.785 2.872 160.3 114.4 130.9 139.2 107.9 109.4 −45.5

1.391 1.382 1.439 1.380 1.378 1.332 1.333 1.466 1.085 1.077 1.077 1.078 1.088 1.089 1.089

−54.0

Cation

Anion

Ion pair

X=F

[BF− 4 ]

X=F

1.390 1.387 1.441 1.380 1.378 1.330 1.334 1.467 1.084 1.077 1.077 1.079 1.088 1.089 1.089

CHO 1.388 1.390 1.441 1.380 1.379 1.330 1.335 1.467 1.084 1.077 1.077 1.079 1.088 1.089 1.088

CN

NO2 1.389 1.387 1.440 1.381 1.379 1.329 1.335 1.468 1.084 1.077 1.077 1.079 1.088 1.089 1.088

1.390 1.388 1.439 1.381 1.379 1.329 1.336 1.468 1.084 1.077 1.077 1.079 1.088 1.089 1.089 1.4088 1.4088 1.4088 1.4088

1.390 1.389 1.436 1.382 1.377 1.327 1.338 1.467 1.085 1.076 1.076 1.082 1.087 1.088 1.089 1.419 1.414 1.429 1.368 1.995 2.372 2.612 2.456 3.773 2.073 2.145 2.783 2.855

CHO 1.388 1.391 1.435 1.383 1.377 1.326 1.339 1.467 1.086 1.076 1.076 1.082 1.087 1.088 1.089 1.419 1.415 1.429 1.367 1.988 2.370 2.616 2.465 3.727 2.067 2.143 2.789 2.859

160.8 113.7

163.5 116.1

130.8 138.6

131.6 138.9

CN

NO2

1.391 1.387 1.430 1.382 1.378 1.323 1.333 1.465 1.085 1.076 1.076 1.076 1.088 1.086 1.090 1.435 1.419 1.407 1.370 2.380 3.228 2.549 2.446 2.503 2.326 2.468 2.753 2.812 2.731 2.646 151.4

1.391 1.388 1.429 1.383 1.378 1.322 1.334 1.465 1.085 1.076 1.076 1.077 1.088 1.086 1.090 1.419 1.408 1.435 1.370 2.379 3.227 2.611 2.405 2.516 2.326 2.478 2.759 2.806 2.729 2.643 113.8

151.4

151.2

122.9

121.4

(continued on next page)

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Table 2 (continued) Parameter

Cation X=F

B11F13N3 B11F13C7 B11F12N3 B11F14C2 C17C16N1C2

−54.2

CHO

−50.6

CN

−51.0

NO2

Anion

Ion pair

[BF− 4 ]

X=F

−49.6

3.2. Structural parameters The main structural parameters calculated at M06-2X/6311++G(d,p) level of theory for TAAILs are given in Table 2. The optimized structure of ion pairs (with the exception of ILs having electron accepting groups CN and NO2) corresponds to a specific geometry with the anion located above the imidazolium (IM) ring, close to the most-acidic C2–H6 group and interacting with the CH groups of Me and phenyl ring. Therefore, it is predicted that the hydrogen bonds to be formed between F atoms of anion and CH of the IM and aryl rings as well as Me group. The F14 atom of BF− 4 lies near the C2–H6 group of IM ring. The C2H6⋯F14 H-bonding distance in [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2) IPs is 2.086, 2.080, 2.084, 2.078, 2.080, 2.073, 2.067, 2.326 and 2.326 Å, respectively. In ILs having electron accepting groups CN and NO2, anion is moved to the upper part of IM ring and far from the CH acidic group so that the Hbond angle also deviates from the typical H-bond angle. The IM and phenyl rings in cations are not coplanar so that dihedral angle between them decreases upon ion pair formation. The value of dihedral angle for NO2 substituent is smaller than others. As can be seen, the bonds involved in H-bonding significantly changes in the IPs with respect to the free ions. The C2–H in all IPs (with the exception of ILs having electron accepting groups CN and NO2), C17– H22 of aryl ring and all B–F bonds involved in H-bonding are lengthened with respect to those of free ions. In ILs having substituents CN and NO2, C2–H6 bond of IM ring is shortened upon ion pair formation. In other words, the effect of IP formation on the C2–H6 bond length of these two IPs is the opposite of those found in other IPs. In all TAAILs, two C–H bonds of the Me group in which have a suitable direction towards the anion are involved in C–H⋯F interactions. The C–H bonds of Me group (C7–H8 and C7–H9) as well as non-participating B– F15 bond of anion are contracted and C7–H10 bond of Me is elongated, compared to the free ions. Correlations between B–F as well as C2–H6 bond lengths and Hammett's constants σp are shown in Fig. 3. The results of geometrical parameters indicate that the H-bond interactions are the major intermolecular structural feature between cations and anions. However, since bonds involved directly in the interaction between the cations and anion do not change significantly, the high BEs obtained can be attributed to the electrostatic attraction between different ions rather than due to weak effect of H-bonding interaction. Although, orientation of anion relative to the cation in all IPs is not suitable for a strong H-bonding interaction, however, H-bond and van der Waals interactions cannot be ignored. Here, a cooperative interaction between the electrostatic attraction and H-bonds was predicted. The similar situation for [EMim+][BF− 4 ] has been reported in [67]. The C16–N1 bond connecting the two rings of the cation is contracted upon formation of IPs, so that the contraction for electron accepting substituents is greater than other ones. Fig. 4 shows the change in C16–N1, C2–N1, C2–N3 bond lengths as well as C2–N1– C11–C12 dihedral angle against the Hammett's constants σp. As can be seen, IP formation increases (decreases) C2–N1 bond length in ILs including the substituents with σ b 0 (σ N 0) with respect to that of free cation. Besides, C2–N3 bond length in all TAAILs decreases upon ion pair formation. There is a simple analysis: the presence of an electrondonating substituent in para position of phenyl ring leads to decrease in mobility of the lone pair at the nitrogen atom N1 towards the phenyl

CHO

108.2 109.5

108.9 109.5

−45.6

−42.5

CN

NO2

91.2 109.5 117.4 95.6 −41.2

91.4 109.4 117.2 95.6 −40.6

ring and a decrease of mesomeric effect. In consequence, it is equivalent to the decrease of π-electron delocalization in the ring. 3.3. Vibrational frequency analysis IR techniques can use for studying the strength of interaction energies between cations and anions in ionic liquids. The structural changes influenced by the formation of an ion pair are often accompanied by shifts in vibrational frequencies with respect to the separable ions. Analysis of the vibrational frequencies shows that the ion pair formation does not lead to a significant shift of vibrational frequencies of C–H bonds involved in interaction. The stretching vibrational frequencies of the C2–H6 in the cations having the substituents NH2, OH, OMe, Me, H, F, CHO, CN and NO2 are 3292, 3287, 3294, 3275, 3276, 3286, 3280, 3272 and 3296 cm−1, respectively, that change to 3252, 3253, 3255, 3251, 3281, 3253, 3248, 3296 and 3302 cm−1 upon IP formation. As can be observed, IP formation is accompanied by a minor change in the vibrational frequency of C2–H6 bond. Because the weak H-bonding interactions existing between differently charged species, H-bonding has a weak effect on the electronic structure of IPs and, in turn, on the vibrational frequencies. The results given in Table 3 show that the change in the symmetric vibrational frequency of the C7–H8 bonds of Me group is greater than those of C2–H6 and C17–H22 bonds involved in interaction with anion. 3.4. AIM analysis Topological parameters derived from Bader's quantum theory of AIM are useful quantities to characterize the chemical bonds. Topological criteria are also useful in detecting the existence of H-bond interactions [68–70]. For tunable ionic liquids, the calculated values of the total electronic density, ρ(r), its Laplacian, ∇2ρ(r), and electronic energy density, H(r), electronic kinetic energy density, G(r), and electronic potential energy density, V(r) at the bond critical points (BCPs) of chemical bonds at M06-2X/6-311++G(d,p) level of theory are listed in Tables S1 and S2 as Supplementary data. The molecular graphs including the bond critical points (BCPs), ring critical points (RCPs), cage critical points (CCPs) and bond paths for all ion pairs are shown in Fig. 5. Besides all of the expected BCPs, the topological features of electron density unveil the additional BCPs in the region between anion and cation of ion pairs. As can be observed, there are seven BCPs, eight RCPs and two CCPs in inter-ionic region of IPs including the CN and NO2 substituents while all other IPs have six BCPs, six RCPs and one CCP. The sum of electron densities, ρ(r), at all bond, ring and cage critical points in the inter-ionic regions of [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2) IPs is 01466, 0.1483, 0.1474, 0.1479, 0.1477, 0.1500, 0.1507, 0.1628 and 0.1629 au, respectively. The results show that the electron density increases as the substituent changes from strong electron-donor to strong electron-acceptor. A good correlation is found between the sum of the electron densities in inter-ionic region and BE. This correlation is shown in Fig. 6(a). Inspection of AIM results shows that the electron density at C2–H6 BCP for IPs having substituents CN (0.2906 au) and NO2 (0.2907 au) is greater than those found in other IPs, which is in good agreement with the smaller C2–H6 bond length obtained for these IPs. The molecular graphs for IPs containing the substituents CN and NO2 do not show any HBCP between H6 and F atoms of anion, in contrast to the other IPs.

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19

Fig. 3. B–F (a) and C2–H (b) bond lengths versus substituent constants (σp) in [X–PhMIM][BF4] ion pairs.

Therefore, there is no C2–H6⋯F H-bonding interaction in IPs including stronger electron-accepting substituents. On the other hand, electron density at C17–H22 and C7–H8 B.P. of the IPs containing the substituents CN and NO2 is smaller than other IPs and for the C7–H9 BCP is greater than those of the others. It is interesting to note that the ρ(r) at the C2–N1, C2–N3, N3–C4 and N1–C16 B.P. of IPs having the substituents CN and NO2 is also bigger than those calculated for other IPs. We have also observed a linear correlation between the sum of electron densities at B11–F12, B11–F13 and B11–F14 B.P. and BEs for all IPs with the exception of the IPs including the substituents CN and NO2 (Fig. 6b). 3.5. Natural bond orbital (NBO) analysis A better understanding of charge transfer interactions in H-bonded complexes is provided by NBO analysis. The formation of a hydrogen bond implies that a certain amount of electronic charge is transferred from the lone pairs of the proton acceptor to the antibonding orbitals of proton donor molecule [71–73]. The results of NBO analysis including charge transfer (CT) values, natural charge and occupancy of NBOs at

M06-2X/6-311++G(d,p) level of theory are given in Table S3 of Supplementary data file. The NBO analysis shows that the LP(F) → σ*(C– H) donor acceptor interactions are the most important intermolecular interactions between anion and cation. Charge transfer energy E(2) corresponding to the LP(F) → σ*(C17–H22) interaction is 1.49, 1.45, 1.34, 1.64, 1.62, 1.76, 2.24, 0.20 and 0.19 kcal/mol and that of LP(F) → σ*(C2–H6) interaction is 0.76, 0.78, 0.79, 0.77, 0.79, 0.80, 0.80, 0.13 and 0.12 kcal/mol in [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2), respectively. Besides, E(2) energy of LP(F) → σ*(C7–H8) interaction is 1.68, 1.78, 1.70, 1.74, 1.73, 1.76, 1.81, 0.50 and 0.45 kcal/mol, respectively. The NBO analysis clearly depicts that the LP(F) → σ*(C2–H6) interaction is weaker than other two interactions, which is in good agreement with the corresponding C–H bond length. The sum of LP(F) → σ*(C–H) donor acceptor interaction energies is 3.93, 4.01, 3.83, 4.15, 4.14, 4.32, 4.85, 0.83 and 0.76 kcal/mol, indicating that the charge transfer from proton acceptor to proton donor through these interactions in ILs including CN and NO2 substituents is weaker than other ones. Although the BE is greater for ion pairs with CN and NO2 electron-attracting substituents than other substituents, the sum of

Fig. 4. Correlation between the C16–N1 (a), C2–N1 (b), C2–N3 (c) bond lengths as well as C2–N1–C16–C17 dihedral angle (d) and σp in [X–PhMIM][BF4] ion pairs.

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Table 3 The M06-2X/6-311++G(d,p) stretching vibrational frequency of the C–H and B–F bonds in ion pairs and corresponding isolated ions. υ/cm−1

X = NH2

OH

OCH3

CH3

H

F

CHO

CN

NO2

Ion pair B–F C7–H8 C17–H22 C2–H6

769 (765)a 3086 3203 3253

769 3085 3218 3253

769 3089 3218 3255

769 3088 3201 3251

768 3091 3211 3281

769 3086 3222 3253

769 3091 3196 3248

766 3083 3221 3296

767 3086 3250 3302

Cation C7–H8 C17–H22 C2–H6

3097 3201 3292

3097 3217 3287

3095 3218 3294

3083 3198 3276

3092 3205 3276

3097 3220 3286

3094 3201 3280

3096 3216 3272

3099 3235 3296

a

Isolated anion.

three LP(F) → σ*(C–H) charge transfer energies is greater in ion pairs with OH, F and CHO substituents. Thus, from the charge transfer energies, it can be predicted that the smallest amount of electron density is transferred from the lone electron pair of F atoms to anti-bonding orbital of C–H bonds in ILs having CN and NO2 substituents. According to the results given in Table S3, electronic density of σ*(C2–H6) is 0.0179, 0.0181, 0.0181, 0.0184, 0.0189, 0.0189, 0.0200, 0.0140 and 0.0140 au in [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2), respectively. Although, electronic density of σ*(C–H) is greater in ion pairs with respect to free cations, the electronic density of σ*(C2–H6) for ion pairs with CN and NO2 substituents is

lesser than other ion pairs, which is in good agreement with the lesser E(2) energy obtained for CN and NO2 substituents with respect to other substituents. The increase in electron density of σ*(C–H) bonds leads to bond elongation in the IPs as compared to those in the free cation. There is a correlation between the sum of charge transfer energies E(2) and the sum of occupancy of σ*(C–H) anti-bonding orbitals involved in interactions. This correlation is shown in Fig. 7. As can be seen, increase in electronic density of antibonding orbitals upon formation of ion pairs is accompanied by an increase in the E(2) energy. From comparison of occupation numbers of N1–C2, C2–N3 and N3– C7 bonds, it was found that the double bond character of N–C bond decreases more with respect to free cations for ion pairs with CN and NO2 electron-attracting substituents, which is in good agreement with the change in corresponding bond lengths. For conventional H-bonds, charge of the H atoms involved in Hbonding becomes more positive. A common observed represent upon ion pair formation in our studied ILs is the positive charge on the H atoms (H6, H7, H8 and H22) involved in H-bonding interactions increases. Substituents in the para position of cations as well as H-bonding interaction with BF− 4 anion can influence the electron density of atoms involved in interactions. The natural charges calculated at M06-2X/6-311++G(d,p) level of theory are given in Table S3. The positive charge of H6, H7, H8 and H22 atoms in C–H bonds increases upon ion pair formation for all substituents so that this increase for electron-accepting substituents is smaller than electron-donating ones.

Fig. 5. Molecular graphs of ion pairs. Nuclei and critical points (bond and ring) are represented by big and small circles, respectively.

H. Roohi, K. Ghauri / Journal of Molecular Liquids 209 (2015) 14–24

21

Fig. 6. (a) Correlation the BEs and sum of the electron densities in inter-ionic region in all IPs. (b) Correlation between the BEs and sum of electron densities at B11–F12, B11–F13 and B11– F14 B.P. in all IPs with the exception of the IPs including the substituents CN and NO2.

The C2 atom in [X–PhMIM][BF4] (NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2) ion pairs is involved in the F⋯H–C interaction and acts as a proton donor. In both cations and IPs, an increase in electronaccepting power of the substituents is accompanied by an increase in the positive charge of the C2 atom. The value of charge on C2 atom of cation increases upon IP formation so that the increase for electronaccepting substituents is bigger than other ones. The relative acidity of C–H groups has an important role in interaction between cation and anions of ILs. The results given in Table S3 reveal that the relative acidity (sum of the charges on the C and H atoms) of the three C–H groups of IM ring is different, so that the relative acidity of C2–H group is greater than C4–H and C5–H groups. Thus, it is expected that the anions are located near the C2–H group. Although all cationic parts of IPs (Me, phenyl and IM rings) carry positive charge, the charge distribution is quite different. For example, the positive charges carried by Me, phenyl and IM rings of IP (cation) are 0.3318 (0.3299), 0.3029 (0.3408) and 0.3276 (0.3293) for X = NH2, 0.3344 (0.3337), 0.2914 (0.3268) and 0.3356 (0.3395) au for X = H and 0.3330 (0.3388), 0.2738 (0.3110) and 0.3639 (0.3502) au for X = NO2. As can be seen, the positive charge is mostly located on the IM ring so that Me group and IM ring carry most of the positive charges (ca. 68%) in IPs. Comparison of atomic charges shows that formation of the Hbonding involves some charge transfer (CT) between the interacting monomers. Analyses of the NBO data allow us to investigate the amount

Fig. 7. Correlation between sum of the occupancy of σ*(C17–H22), σ*(C2–H6) and σ*(C7– H8) anti-bonding orbitals and sum of the corresponding E(2) for the LP(F) → σ*(C–H) donor–acceptor interactions.

of charge transfer due to electron delocalization between species. Here, the charge transfer is defined as the difference between the sum of atomic charges on complexed and isolated anion. Results of population analysis show that the charge is transferred from BF− 4 anion to cation [X–PhMIM+]. The CT values obtained for [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F, CHO, CN and NO2) ion pairs are 0.0378, 0.0385, 0.0379, 0.0388, 0.0387, 0.0398, 0.0410, 0.0293 and 0.0293 au, respectively, which is in good agreement with the greater E(2) energy obtained for F and CHO electron-attracting substituents with respect to other substituents. The weak H-bonding between anion and cation is responsible for a partial electron transfer. The results illustrate that the CT value for electron-accepting substituents (X = CN and NO2) is smaller than other ones. Fig. 8 shows a correlation between the BE and CT for [X–PhMIM][BF4] (NH2, OH, OCH3, CH3, H, F and CHO) ion pairs. This figure reveals that an increase in CT is accompanied by an increase in BE. 3.6. Frontier molecular orbital analysis In addition to NBO population analysis, canonical molecular orbital analysis provides us a necessary understanding for the charge transfer interactions. ELUMO and ELUMO of ion pairs and the corresponding ions calculated at M06-2X/6-311++G(d,p) level of theory are listed in Table S4. Fig. 9 shows the frontier molecular orbital (FMO) energy diagram for [X–PhMIM][BF4] (X = NH2, H and NO2) ion pairs. As can be seen in Fig. 9, HOMO of anion consists of three deformed p orbitals (EHOMO = −6.9397 au) localized on the four F atoms and LUMO of cations is a π* MO in nature (EHOMO = −3.8346, −4.1704 and −5.0031 au). HOMOs are localized on the IM and phenyl rings of cations rather

Fig. 8. Correlation between the values of BEs and CT for [X–PhMIM][BF4] (X = NH2, OH, OCH3, CH3, H, F and CHO) ion pairs.

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H. Roohi, K. Ghauri / Journal of Molecular Liquids 209 (2015) 14–24

Fig. 9. Frontier molecular orbital (FMO) energy diagrams for [X–PhMIM+][BF− 4 ] (X = NH2, H and NO2) ion pairs.

than anion BF− 4 . Thus, electron density over the F atoms of anion decreases upon ion pair formation, indicating the CT occurs from the anion to cations. The nature of LUMO of mentioned IPs is different. The LUMO in [X–PhMIM][BF4] (X = NH2 and H) is located across the phenyl ring, while in [NO2–PhMIM][BF4] is distributed on phenyl ring and partially on the IM ring. The stability of ionic liquids with respect to high voltage differences is important to use in electrochemical applications. Ionic liquids are decomposed when the voltage differences applied are larger than their electrochemical window (4–6 V). The electrochemical stability of ionic liquids depends principally on the resistance of the cation against reduction and the resistance of the anion against oxidation. Quantum chemical calculations can be employed to predict the electrochemical stability of ionic liquids. It has been shown that the electrochemical stability of ionic liquids is correlated with the electron affinity A = −ELUMO. The low ELUMO value indicates that the electron accepting ability of the molecule is higher [74–78]. Since the LUMO of studied ILs is located completely on the cation, the resistance of the cation against reduction principally determines the stability of the ionic liquid with respect to reduction. The linear correlation between LUMO energy of cations versus Hammett's constant (σp) is depicted in Fig. 10. From this result, it can be predicted that the resistance of the cation against reduction increases as the substituents change from the electron-accepting to electron donating substituents. Thus, electrochemical stability of the ILs towards reduction is greater for ILs having electron donating substituents. In other words, ILs with electron-donating substituents have the more positive reduction potential than those of ILs including electron-accepting ones. Accordingly, it is predicted that the ILs with electron-donating substituents are more suitable electrolyte than other ILs in electrochemical reactions.

formation of ILs [81]. The thermodynamic quantities are estimated with the help of the following reactions: 1C10H12N3 + 13O2 → 10CO2 + 6H2O + 3/2 N2 (X = NH2)

1C10H11N2O + 49/4O2 → 10CO2 + 22/4H2O + 1 N2 (X = OH)

1C11H13N2O + 55/4O2 → 11CO2 + 26/4 H2O + 1 N2 (X = OMe)

1C11H13N2 + 57/4O2 → 11CO2 + 26/4H2O +1 N2 (X = Me)

1C10H11N2 + 51/4O2 → 10CO2 + 22/4H2O + 1 N2 (X = H)

1C10H10N2F + 49/4O2 → 10CO2 + 18/4H2O + 1 N2 + 1HF (X = F)

3.7. Enthalpy and Gibbs free energy of combustion and formation of cations The enthalpy and Gibbs free energy of combustion and formation of cations in ILs are important thermodynamic data. We have used CBSQB3 composite method [79,80] to obtain the standard molar enthalpies of combustion and standard molar enthalpies of formation of cations. Besides, the similar calculations are performed to calculate the Gibbs free energies. This method has been used for calculation of heat of

Fig. 10. Correlation between the LUMO energies of [X–PhMIM+] and Hammett's constants (σp).

H. Roohi, K. Ghauri / Journal of Molecular Liquids 209 (2015) 14–24

1C11H11N2O + 53/4O2 → 11CO2 + 22/4H2O + 1 N2 (X = CHO) 1C11H10N3 + 27/2O2 → 11CO2 + 5H2O + 3/2 N2 (X = CN) 1C10H10N3O2 + 23/2O2 → 10CO2 + 5H2O + 3/2 N2 (X = NO2). The standard enthalpy of combustion reaction, ΔH∘c,298, and standard enthalpy of the formation of cations, ΔH∘f,g,298, were computed using the calculated enthalpy changes of the above chemical reactions and experimental ΔH∘f,g,298 of H2O, CO2 and HF (for X = F). The thermodynamic quantities obtained using the CBS-QB3 method are given in Table 4. The calculated values of ΔH∘c,298 and ΔH∘f,g,298 are ranged from − 1756.3 (X = CHO) to − 2007.4 (X = OMe) kcal/mol and 502.7 (X = F) to 622.6 (X = CN) kcal/mol, respectively. We have also calculat∘ ed ΔGf,g,298 and ΔG∘c,298 of cations using CBS-QB3 method. The results are presented in Table 4. The positive values of ΔG∘f,g,298 demonstrate that the formation of cations from the corresponding separated elements at 298 K is not a spontaneous process.

4. Conclusions The influence of various substituents in the para position of [X– PhMIM][BF4] TAAILs on their electronic, structural and thermochemical properties was investigated in gas and solution phases using the quantum chemical calculations at M06-2X level of theory. The change in binding strength due to variation of the substituents in the para position was well displayed by change in interaction energy, structural parameter, electron density properties, natural charge, charge transfer and percentage of p-character of C atom in C–N and C–H bonds. Our results predicted that the hydrogen bonds to be formed between F atoms of [BF− 4 ] and CH groups of the Me, IM and aryl rings. The results show that the aromatic-based ILs having the electron-donating substituents present lower interaction strength compared to the electron-accepting ones. There is found a correlation between change in BEs and changes in melting points, conductivity, surface tension and critical-point temperature of IPs. A good correlation is observed between the sum of the electron densities in inter-ionic regions of IPs and BE values. NBO analysis shows that the smallest amount of electron density is transferred from the lone electron pair of F atoms to anti-bonding orbital of C–H bonds in ILs having CN and NO2 substituents. It is predicted that the ILs with electron-donating substituents are more suitable electrolyte than other ILs in electrochemical reactions.

Appendix A. Supplementary data Supplementary data to this article can be found online at http://dx. doi.org/10.1016/j.molliq.2015.05.001.

Table 4 ΔH∘c,298, ΔG∘c,298, ΔH∘f,g,298 and ΔG∘f,g,298 (kcal/mol) values obtained for cations by using the CBS-QB3 composite method. X

ΔH∘c,298

ΔG∘c,298

ΔH∘f,g,298

ΔG∘f,g,298

NH2 OH OMe CH3 H F CHO CN NO2

−1947.4 −1769.3 −2007.4 −1820.7 −1905.2 −1853.7 −1756.3 −1763.3 −1962.5

−1983.5 −1807.6 −2046.1 −1854.9 −1944.5 −1895.9 −1803.4 −1802.9 −2008.0

565.1 503.8 550.9 595.8 561.1 502.7 551.5 622.6 525.6

624.3 558.7 614.8 652.9 610.7 552.8 606.0 672.3 586.6

23

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