CO2 Absorption Using Fluorine Functionalized Ionic Liquids: Interplay

Article pubs.acs.org/JPCA

CO2 Absorption Using Fluorine Functionalized Ionic Liquids: Interplay of Hydrogen and σ‑Hole Interactions Soniya S. Rao and Shridhar P. Gejji* Department of Chemistry, Savitribai Phule Pune University, Pune 411 007, India S Supporting Information *

ABSTRACT: Use of ionic liquids (ILs) for CO2 capture offers certain advantages over currently used methodologies and is of growing interest. With this perspective, ILs composed of S-ethyl-N,N,N′,N′-tetramethylthiouronium ([ETT]) and 1-hexyl-3-methylimidazolium ([Hmim]) cations and tris(pentafluoroethyl)trifluorophosphate ([FEP]) anion have been investigated. The present work unravels the noncovalent interactions accompanying CO2 capture by these ILs. Electronic structure of ion pairs and their CO2 absorbed [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) (n up to 30) complexes are derived. The anisotropy in molecular electrostatic potential dictates the binding of CO2 through the interplay of (i) halogen bonding (O···F) between electron deficient σ-holes on fluorines, (ii) electrostatic C···F interactions between electron deficient carbons of CO2 and the electron-rich fluorine atoms, and the (iii) hydrogen bonding (O···H) interactions from the cation. The manifestations of these interactions on binding energies, polarizabilities, and vibrational spectra of CO2 absorbed complexes are presented. Consequent “frequency shift” accompanying hydrogen and halogen bonding exhibit complementary characteristics in the infrared spectra of CO2 absorbed complexes. Correlation of binding energies to absorbed CO2 molecules further demonstrate that [Hmim] based ILs are more efficient for CO2 capture applications.

1. INTRODUCTION

On the theoretical front the molecular dynamics (MD) simulations demonstrated that interactions from the anion contribute significantly toward dissolution of CO213,25,26 The solubilities of CO2 in several ILs based on varying anions follow the order: [NO3] < [DCA] < [BF4] ≈ [PF6] < [TfO] < [Tf2N] < [Tf3C] (with [TfO], [Tf2N], and [Tf3C] containing one, two, and three CF3 groups, respectively). It has been concluded that structural attributes such as fluorine substituted cation or anion, bromination of anion, long alkyl groups with branched chains, ether linkages, and carbonyl or ester substituents on the cation enhance the physical absorption of CO2 compared with their alkyl analogues.27−32 Maginn and co-workers examined ILs composed of 1-hexyl-3-methylimidazolium ([Hmim]+) and fluorinated phosphate anions ([PF6]− and [PF3(RF)3]− (FEP), where RF = C2F5, C3F7, or C4F9) through MD simulations. The mechanism of CO2 absorption was found to be dependent on the size and shape of the anion.28 The symmetric [PF6] anion attracts CO2 strongly due to electrostatic interactions and substitution of CF2 or CF3 groups on the [FEP] anion renders charge delocalization with diminutive electrostatic interactions between CO2 and [FEP] accompanied by increased van der Waals attractions on absortion of CO2. The asymmetric [FEP] anion possesses large free volume and provides binding sites for

Escalating concentrations of carbon dioxide (CO2) in the atmosphere has resulted in global warming and therefore viable solutions to reduce CO2 have been of growing interest.1,2 Solid adsorbents, biomimetic approaches, membrane separation, and chemical or physical absorption3−6 are the existing methodologies for CO2 capture and storage. Recent absorption methods for reducing CO2 based on amine solvents are energy intensive and far from cost-effective for large scale applications.7−9 Alternatively ILs have emerged as a fascinating class of compounds that can overcome the limitations associated with the current CO2 capturing technologies.10−14 Varying combinations of cations and anions or their chemical functionalities in ILs render an extra degree of freedom to adjust physicochemical properties such as volatility, thermal and chemical stability, solvation, and CO2 solubility.15,16 Further the “tunability” of ILs has been explored to design and synthesize task specific ionic liquids (TSILs)17 for the removal of captured CO2 via physical absorption. Interest in CO2 capture by TSILs stems from the works of Bates et al.18 with functionalization of imidazolium cation incorporating primary amine moiety. Enhanced CO2 solubility with substitution of primary, secondary, or tertiary amine groups on cations and anions have subsequently been reported in the literature.19−24 TSILs with amine groups enhance CO2 capacity; nonetheless, higher viscosities in these systems continues to be a bottleneck. © XXXX American Chemical Society

Received: December 12, 2015 Revised: January 14, 2016

A

DOI: 10.1021/acs.jpca.5b12161 J. Phys. Chem. A XXXX, XXX, XXX−XXX

Article

The Journal of Physical Chemistry A

Figure 1. Structures of (a) [ETT] and (b) [Hmim] cations and (c) [FEP] anion along with atomic labeling scheme.

accommodating more CO2. Henry constants from the COSMO-RS approach are used to compare CO2 solubilities in ILs incorporating a variety of cations and anions.32 Pyridinium, thiouronium, and imidazolium cations with varying alkyl chain length in combination with FEP anion showed low Henry’s constants demonstrating higher CO2 solubilities.33 Thiouronium and imidazolium based ILs are liquids at room temperature which are hydrophobic with higher thermal stability that makes them suitable for CO2 capture.34 Modifications in cation or anion moieties constituting the ILs, thus, should be directed toward synthesis of ILs possessing lower viscosities for improved CO2 capacity and selectivity as well. With this view, S-ethyl-N,N,N′,N′-tetramethylthiouronium ([ETT]+) and 1-hexyl-3-methylimidazolium ([Hmim]+) cations combined with tris(pentafluoroethyl)trifluorophosphate ([FEP]−) anion have been investigated in this work. It has been recognized that solubility of gas in ILs is primarily governed by interactions between the cation(s) or anion(s) constituting the IL and the distribution of CO2 molecules surrounding the ion pair(s). Absorption of CO2 is facilitated through noncovalent interactions, that is, hydrogen bonding as well as electrostatic, steric, and other dispersive interactions. Density functional calculations on brominated imidazolium based ILs exhibiting enhanced CO2 solubilities have been attributed to halogen bonding between CO2 and the ion pairs.35 The present endeavor aims at understanding hydrogen and halogen bonding interactions between CO2 and fluorinated ILs composed of [ETT]+, [Hmim]+ cations, and [FEP]− anion, as prototype examples. The anisotropy in electronic charge distribution around fluorines renders them as Lewis acid or

base facilitating different modes for interactions with CO2 simultaneously.36 Here the CO2 acts as a weak Lewis base bringing about hydrogen bonding interactions with the cation and electrostatic attractions between electron deficient carbon of CO2 and electron-rich fluorines on the anion via C···F interactions, in addition to halogen bonding (O···F) between electron deficient σ-holes on the fluorines of the anion. The effect of halogen and other molecular interactions has been systematically analyzed with the sequential addition of CO2. In light of this, the halogen bonding interactions driven by σ-hole, a positively charged region on the equatorial side of halogen (X) along the R−X covalent axis have been analyzed.37−42 The interplay of such noncovalent interactions is discussed in terms of electronic structure, binding energies, and spectral characteristics of the [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) complexes. Finally, the topological and bonding features have been harnessed through noncovalent interaction reduced density gradient (NCI-RDG) method.43

2. COMPUTATIONAL METHODS 2.1. Molecular Electrostatic Potential and Ion Pair Structures of the ILs. Atomic numbering schemes for [ETT] and [Hmim] cations and [FEP] anion are shown in Figure 1. Topological features of molecular electrostatic potential (ESP) are summarized by the critical points (CPs) in terms of rank and signature.44−46 Characteristic Vs,max (in red) and Vs,min (in blue) in ESP on the 0.001 au isodensity surfaces were located for the minimum energy structures of [ETT][FEP] and [Hmim][FEP] ion pairs within the framework of M05-2X/6311++G(d,p) theory. The CPs in [ETT][FEP] and [Hmim]B


Article

The Journal of Physical Chemistry A [FEP] ion pairs provide potential CO2 binding sites. Initial structures of [ion pair]·n(CO2) conformers were obtained using the procedure outlined below. A sphere of radius ∼1.5 Å around these minima or saddles in ESP was generated. The carbon atoms of CO2 were then placed in their neighborhood (with varying θ for a given Φ). Further C atoms falling within the cation−anion binding region were discarded. In the next step, oxygen atoms were aligned in parallel/perpendicular fashion around the selected C imposing constraints for O···O, C···C, and C···O separations between the CO2 molecules followed by minimizing the electronic repulsions subsequently. The [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) (n = 1− 6, 8, 10, 12, 16, 20, and 30) conformers thus obtained were initially subjected to optimizations at the B3LYP/6-311G(d,p) level of theory47,48 employing the GAUSSIAN-09 program.49 To account for underlying hydrogen bonding and dispersive interactions facilitating CO2 capture, we further employ the hybrid meta-GGA (generalized gradient approximation) exchange correlation M05-2X functional combined with the 6311++G(d,p) basis set.50−55 The stationary point structures of [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) (n = 1−6, 8, and 10) complexes obtained were confirmed to be local minima from the vibrational frequency calculations. Normal modes were assigned through the visualization of displacement of atoms around their equilibrium position using the GAUSSVIEW-5 program.56 The potential energy distributions (PED) were derived for each stationary point structure. Binding energies were obtained by subtracting the sum of energies of individual ions and number of CO2 molecules making up the complex in their free state (isolated) from those of [ETT][FEP]·n(CO2) or [Hmim][FEP]·n(CO2) systems. To delve into molecular interactions further, topological parameters were derived with the QTAIM approach 57 employing AIMAll program.58 The frequency shifts of characteristic vibrations in the infrared spectra were analyzed from second order perturbation estimates of electron transfer delocalization energies (E(2)) in natural bond orbital (NBO) analyses.59 Noncovalent interaction (NCI) reduced density gradient (RDG) method43 has further been utilized to unravel the intermolecular interactions in [ETT][FEP]·n(CO2) or [Hmim][FEP]·n(CO2) systems. To corroborate the inferences from the NCI method, energy decomposition analysis (EDA) was carried out.

Figure 2. Optimized structures of (a) [ETT][FEP] and (b) [Hmim][FEP] ion pairs. Hydrogen bonding interactions are depicted as broken lines.

(∼0.3 Å) compared to the free cation. The interacting C−Fa bonds reveal an elongation of 0.03 Å and distinguish themselves from the noninteracting ones. A three-center four-electron πsystem in [Hmim] cation akin to imidazolium cations has earlier been reported.62 The most acidic H1 proton interacts with anion (cf. Figure 2b) via bifurcated C−H···F hydrogen bonds (the corresponding separations are being 470 and 2.248 Å). Thus, pentafluoroethyl groups facilitate more hydrogen bonding interactions and render symmetric structure to the [ETT][FEP] unlike the [Hmim][FEP] ion pair, which provides large free space for further CO2 absorption. 3.2. Electrostatic Potentials in [ETT][FEP] and [Hmim][FEP] Ion Pairs. The noncovalent interactions in biological and chemical systems encompassing hydrogen bonding, π−π stacking, cation−π and anion−π interactions, and lone pair−π interactions have widely been studied in the literature.63−66 Halogen bonding interactions with the halogen atom as acceptor are of particular interest owing to their strength and directionality, which compares well with those of hydrogen bonding. Understanding of molecular interactions has proven useful for designing materials with potential applications in supramolecular architectures, crystal engineering, molecular

3. RESULTS AND DISCUSSION 3.1. Ion-Pair Structures of [ETT][FEP] and [Hmim][FEP]. Ion pairs have widely been employed to model liquidus behavior and physicochemical properties of ILs.60,61 The initial structures of the ion pairs were generated from the CPs, and the ion pair conformers up to 5 kJ mol−1 higher in energy relative to its lowest energy structure were further subjected to optimizations within the M05-2x/6-311++G(d,p) framework of density functional theory. The lowest energy conformers of [ETT][FEP] and [Hmim][FEP] ion pairs converged to only the conformer shown in Figure 2. Selected bond distances and angles in these ion pairs are compared with those of the free cations and anion in Table 1. Hydrogen bonding distances are also given. The anion interacts with methyl and ethyl groups of the cation through C−H···F hydrogen bonding interactions in [ETT][FEP] ion pairs, the separations being in the range of 2.239−2.564 Å (cf. Table 1). Electron density redistribution accompanying formation of ion pair results in the shortening of C1−N2 and C1−N3 bonds C


Article

The Journal of Physical Chemistry A Table 1. Selected Bond Distances (in Å) of [ETT]+ and [Hmim]+ Cations, [FEP]− Anion, and [ETT][FEP] and [Hmim][FEP] Ion Pairs, Including Hydrogen Bonding Distances [ETT]+ C1−N2 C1−N1 C1−S N1−C4 N1−C5 N2−C2 N2−C3 P−F1 P−F3 P−F2 C6−F17 C6−F18 C1−F5 C2−F6 C3−F9 C1−H1 C1−N1 C1−N2 N2−C4

[FEP]−

1.330 1.333 1.743 1.467 1.464 1.468 1.468 1.649 1.660 1.650 1.333 1.327 1.360 1.330 1.361

[ETT][FEP]

[Hmim]+

[Hmim][FEP]

1.327 1.330 1.752 1.469 1.458 1.466 1.462 1.659 1.690 1.632 1.348 1.374 1.358

1.074 1.329 1.331 1.465 Hydrogen Bond Distances (Å) [ETT][FEP]

H7···F1 H5···F5 H1···F11 H17···F13 H7···F5 H13···F11 H13···F2 H11···F19 H9···F19 H1···F3

contrary to those for oxygens. Figure 3b represents the ESP mapped surface in the [Hmim][FEP] ion pair. As shown, an envelope of negative ESP brings about the partial shielding of electron-deficient regions originating from σ-holes on the C−F covalent bonds. ESP of the ion pair can be envisaged through segregation of the electron-rich and electron deficient regions depicted in Figures 3c,d. The (bare) nuclei contributions to electrostatic potential prevail around the atom, and consequent formation of the covalent bond brings about polarization of the electron density leaving behind an electron-deficient region on the outer side with an increase of the electron density around equatorial sides. A positive σ-hole (Vs, max > 0) in the molecular system engages in establishing σ-hole interactions with negative sites occupied on another or the same molecule. This can be deciphered from Figure 3c. As shown, the concentrated blue σholes are at the apex of C−F bonds on the pentafluoroethyl groups of the anion. The following inferences may be drawn. The anisotropy in the ESP distribution of fluorine atoms of the ion pair facilitates (i) O···F interactions from the positive σholes at the apex along the extension of covalent C−F bonds and (ii) C···F interactions between the fluorines and electropositive carbon atoms (Cδ+) of the CO2 molecule. These σ-hole interactions are mutually dependent and the CO2 affinity of ion pair arises from head-on O···F halogen bonding and C···F interactions resulting from the compensation of negative and positive ESP. Hence, anisotropic electronic distribution of the fluorinated anion is crucial in CO2 absorption. Parallel inferences are drawn for the [Hmim][FEP] system. 3.3. Complexation of CO2 with Individual [ETT]+ and [Hmim]+Cations and [FEP]− Anion. M05-2x/6-311++G(d,p) optimized structures of isolated free ions with CO2 molecules are derived. The CO2 molecule facilitates weak hydrogen bonding with [ETT] and [Hmim] cations. The anion mainly interacts with the CO2 molecule via strong O···F halogen bonding with a deviation of ∼2° from linearity thereby showing larger affinity toward CO2. MD simulations on CO2 absorption reveal gradual diffusion of gas molecules to ILs.82 It therefore would be instructive to understand the CO2 binding to isolated cation and anion as well as the ion pair at the CO2/ IL interface under equilibrium conditions. Thus, charge distribution and binding of CO2 molecules surrounding the ion pair provide insights for solvation at the molecular level. As an initial step, we derive to the electronic structure and spectral features of [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) (n = 1−6, 8, 10,12, 16, 20, and 30) with subsequent addition of CO2. 3.4. Structures of [ETT][FEP]·1(CO2) or [Hmim][FEP]· 1(CO2) Systems. As shown in Figure 4a, the CO2 molecule participates in CO···H hydrogen bonding interactions with the [ETT]+ cation simultaneously facilitating σ-hole (O···F and C···F) interactions via CF2/CF3 groups and P−F1 end. An elongation of 0.003 and 0.008 Å is observed for the respective C6−H12 and P−F1 bonds (cf. Table 1). On the other hand, the interacting C−Fa bonds are shortened consequent to O···F halogen bonding interactions compared to the unabsorbed ion pair. The σ-hole angle (C−Fa···LB), defined as the angle between the covalent C−F bond of the anion and the CO2, turns out to be 136.7°, and reveals remarkable deviation from linearity. The deviation of CO2 from linearity (177.8°) can be attributed to these interactions. The optimized structure of [Hmim][FEP]·1(CO2) is shown in Figure 4b. The elongation of C−Ha (0.005 Å) bonds and the shortening of interacting C−

1.682 1.636 1.679 1.321 1.351 1.356 1.330 1.363 1.072 1.327 1.326 1.463

[Hmim][FEP] 2.390 2.436 2.337 2.564 2.261 2.414 2.290 2.489 2.380 2.596

H1···F18 H9···F6 H1···F1 H9···F2 H12···F9 H10···F15 H4···F18 H1···F15 H10···F2

2.248 2.421 2.451 2.366 2.580 2.545 2.546 2.470 2.511

recognition, and biological systems.67−74 Halogen bonding comprises of electronegative halogen that renders attractive interactions with a negative site and can be envisaged as a subset of “σ-hole” bonding. The concept of σ-hole introduced by Clark et al.75 describes the regions of positive electrostatic potential on the electronically depleted outer portion of the half filled p-type orbital of halogens. These positive regions bring about electrostatic interactions with lone pairs and π-electrons on the same or other molecules giving rise to highly directional halogen bonding. The σ-hole results when the halogen atom participates in a covalent σ-bond and subsequent electron density redistribution leaves behind a diminutive negative charge region on its outer noninteracting side. The resulting charge anisotropy engenders interactions between the positive σ-holes and the negative sites.41,76−81 The condensed phase systems are accompanied by the hydrogen bonding, σ-hole, and other dispersive interactions.40,79−81 The interplay of two or more such interactions should prove crucial for the absorption of CO2 by ILs. The [ETT][FEP] and [Hmim][FEP] based ILs reveal the coexistence of all such interactions on CO2 absorption. ESP topography can be used to delve into the nature of such interactions facilitated through fluorine atoms of the anion. The ESP mapped isodensity surfaces (0.05 au) of CO2 displays electron deficient regions around the central C atom D


Article


Figure 3. ESP mapped isodensity surface (0.005 au) in (a) isolated CO2 molecule and (b) [Hmim][FEP] ion pair. Large negative ESP (golden region) surrounding electron-deficient positive ESP (greenish-blue). (c) Positive ESP isosurface (0.621 a.u) at the apex of C−F covalent bonds of the [FEP] anion in the [ETT][FEP] ion pair. (d) Negative ESP regions in the ion pair are shown.

fluorine to the π* orbital of CO2, characteristic of the halogen bonding interactions. An influx of charge density from the lone pair on oxygens of CO2 to the C−Ha covalent bond participating in hydrogen bonding further can be noticed. Second order perturbation components of energy are given in Table S3 of the Supporting Information. The C−Ha covalent bonds in [ETT][FEP]·1CO2 are elongated consequent to hydrogen bonding contrary to the interacting C−Fa bonds participating in halogen bonding. Here the covalent bonds reveal increased polarization as a combined result of hydrogen and σ-hole interactions. Noteworthy enough, the halogen bonding interactions reveal diminutive occupancies for the corresponding σ* orbitals as opposed to the hydrogen bonding. Moreover, the s-character of C atoms shows complementary behavior, and is accompanied by increased electron-deficient character of the carbon and fluorine atoms (cf. Table 2). The hyperconjugation and rehybridization effects can be envisaged through hydrogen bonding and σ-hole interactions. An increase

Fa bonds is evident. Structural parameters are displayed in Table 2. These interactions are further analyzed in terms of the electron density at the bond critical point (ρbcp) and its Laplacian (∇2ρ) in the QTAIM theory reported in Table 3. Koch and Popelier83 categorized hydrogen bonding (O···H) interactions based on the ρbcp parameters in the range 0.002− 0.035 au, with the corresponding Laplacian values being 0.014− 0.139 au. QTAIM considerations also predicted the presence of halogen bonding interactions. The strength of O···H, O···F, and C···F interactions are correlated to ρbcp parameters. The alkyl chain protons of the [Hmim] cation facilitate more hydrogen bonding interactions with CO2. NBO analyses were applied to gain deeper insights for counterbalancing of O···H, C···F, and O···F interactions. Binding of CO2 to ion pairs revealed transfer of electron density from donor (Lewis base) to the acceptor (cation or anion) accompanied by a charge transfer from the lone pairs on E


Article


are in consonance with the Bent’s rule84,85 which states that “The p-character tends to concentrate in orbitals with weak covalency (arising from either electronegativity or overlap considerations), and s-character tends to concentrate in orbitals with strong covalency (matched electronegativities and good overlap)”. Thus, the hydrogen and halogen bonding interactions engender a concomitant increase of the electrondeficient nature of H and F atoms, respectively, and the decrease and increase of s-character in their respective hybrid orbitals renders large affinity for CO2. The [Hmim][FEP]· 1(CO2) complex led to similar conclusions. 3.5. Structures of [ETT][FEP]·n(CO2) or [Hmim][FEP]· n(CO2) Aggregates. To further understand the phenomenon of CO2 capture, we derive the [ETT][FEP]·n(CO2) (n = 2−6, 8, 10, 12, 16, and 30) aggregates. The underlying hydrogen, halogen, and C···F σ-hole distances with the corresponding ρbcp and their Laplacian parameters given in Table 3 reaffirm the presence of C···F and O···F interactions from the anionic end. A deviation up to ∼3° from linearity for the CO2 was noticed. As inferred (Figure 5) relatively strong σ-hole interactions are observed on subsequent addition of CO2. Further addition of CO2 beyond n = 8 facilitates O···H hydrogen bonding via the cationic end of the ion pair. A comparison of [ETT][FEP] and [Hmim][FEP] aggregates up to n = 8 reveals more hydrogen bonds as well as the σ-hole interactions for the latter. On subsequent addition of CO2 (n > 10), the binding sites in the ESP surrounding the [FEP] anion are consumed and are no longer available; hence, CO2 binds to the cation of the ion pair through hydrogen bonding. Subsequently, molecular recognition of CO2 from the cationic end of the ion pair becomes crucial. The distribution of CO2 around the [ETT][FEP] ion pair with 16 CO2 molecules can possibly be considered as first solvation shell (cf. Figure 6a,b). As far as [Hmim][FEP]· n(CO2) systems are concerned, a comparison of n = 12 and n = 16 absorbed molecules reveals that the alkyl chain on the cation brings about more hydrogen bonding interactions (cf. Figure 6c,d). The CO2 solvation further extends beyond n = 16. This is evident from Figure 6b,d. On further absorption of CO2, the accessibility of σ-hole is shielded by the surrounding electron rich belt of fluorine causing repulsions toward the CO2. The cumulative effect of σ-holes, negative ESP, and steric hindrance from the absorbed molecules results in the formation of “‘potential wells”’ or “voids’ through which the CO2 escapes. This becomes apparent for n = 20 CO2 molecules in both [ETT] and [Hmim] aggregates portrayed in Figure 7. Further addition of CO2 engenders saturation of its solvation

Figure 4. Optimized structures of (a) [ETT][FEP]·1(CO2) and (b) [Hmim][FEP]·1(CO2) ion pairs. Hydrogen bonding interactions are depicted as broken lines.

of the electron deficient character of the H atom (from the cation of the ion pair) can be observed on CO2 binding. The polarization of the C−H(F) Lewis acid bond (% at C atom) and the s-character of the C atom increases as a result of hydrogen and halogen bonding (cf. Table 2). These inferences

Table 2. Bond Distances (in Å), % s-Character for C Centers, Polarization, Net Atomic Charges, and Population of Antibonding Natural Orbitals Facilitating Hydrogen- or σ-Bonding Interactions from NBO Analyses in [ETT][FEP]·1(CO2) and [Hmim][FEP]·1(CO2)a r

a

% s character (C)

C4−H7 C8−F18 C10−F8

1.084 (1.081) 1.325 (1.327) 1.332 (1.333)

27.43 (27.32) 22.98 (23.04) 22.48 (22.70)

C1−H1 C5−F15 C6−F16 C6−F18

1.072 1.335 1.370 1.351

39.01 20.61 22.07 21.61

(1.071) (1.336) (1.373) (1.353)

(38.88) (20.70) (22.09) (21.73)

polarization [ETT][FEP]·1(CO2) 0.7936 (0.7922) 0.5298 (0.5302) 0.5285 (0.5301) [Hmim][FEP]·1(CO2) 0.7994 (0.7786) 0.5162 (0.5177) 0.5284 (0.5283) 0.5210 (0.5214)

charge on X (X = H or F)

σ*

0.2525 (0.2532) −0.3387 (−0.3360) −0.3515 (−0.3421)

1.9881 (1.9885) 1.9554 (1.9955) 1.9955 (1.9955)

0.2702 0.3929 0.3488 0.3812

(0.2683) (−0.3878) (−0.3496) (−0.3785)

1.9800 1.9933 1.9937 1.9951

(1.9808) (1.9935) (1.9952) (1.9954)

The values in parentheses represent those of ion pairs. See text for details. F


2.726

2.463 2.775 2.780 2.448−2.709

2.439−2.505

2.523−2.721

2.429−2.772 2.428−2.783

2.641 2.649 2.435 2.758 2.780 2.490−2.898 2.481−2.779

2.427−2.786 2.492−2.789 2.549−2.689

1

2

4

6

8 10

1

G

6 8 10

3 4

2

3

r (Å)

no. of CO2, n

0.0030−0.0089 0.0022−0.0084 0.0045−0.0060

0.0062 0.0072 0.0061 0.0083 0.0085 0.0037−0.0069 0.0047−0.0071

0.0036−0.0083 0.0047−0.0084

0.0043−0.0086

0.007−0.008

0.008 0.007 0.004 0.004−0.009

0.006

ρbcp (au)

O···H

+0.0226 +0.0282 +0.0268 +0.0298 +0.0378 +0.0122 to +0.0250 +0.0162 +0.0256 +0.0106 to +0.0089 +0.0095 to +0.0294 +0.0169 +0.0212

+0.015 +0.025 +0.027 +0.0155 +0.309 +0.026 +0.036 +0.0154 +0.0311 +0.0132 to +0.0029 +0.0166 to +0.0299

+0.0273

∇2ρ (au)

∇2ρ (au)

+0.038 +0.044 +0.0057 to +0.0081 +0.0296 to +0.0439 +0.0226 to +0.0397 +0.0362 to +0.042 +0.0266 +0.0360

0.0085 0.0090 0.0052−0.0081 0.0058−0.0085 0.0041−0.0079 0.0058−0.0083 0.0050 0.0071

2.804−2.999 2.762−3.553 2.695−3.313 2.793−3.392 2.798−3.544

0.0058−0.0097 +0.0304 to +0.0476 +0.0512 to +0.037 0.0104−0.0074 (b) [Hmim][FEP]·n(CO2) 0.009 +0.045

+0.028 to +0.046

2.765 2.778

2.907

2.729−3.274 2.693−3.087

3.207−3.425

0.0056−0.0093

+0.296 +0.0252 +0.018 to +0.042

0.0091 0.0058 0.0039−0.0086

2.683−3.217 2.641−3.315

+0.040 +0.041

0.008 0.008

(a) [ETT][FEP]·n(CO2) 0.012 +0.0552

ρ (au)

2.823 2.804

2.666

r (Å)

C···F

3.043−3.426 2.290−3.483 2.985−3.532

2.977 3.140 2.827 3.065 3.189 3.027−3.298 3.029−3.838

2.958−3.389 3.011−3.407

178.7 178.8 178.7 178.9 178.8

+0.0110 to +0.0262 +0.0161 to +0.0282 +0.0112 to +0.0387 0.0020−0.0057 0.0031−0.0058 0.0020 0.0078

178.4

179.1

178.8 178.9

178.9

+0.020 +0.035 +0.0475 +0.0162 +0.0259 +0.0301 to +0.0041 +0.0160 to +0.0294

+0.019 to +0.031 +0.0146 to +0.029

+0.0217 to +0.0331

178.6

178.4

178.8

177.9

avg OCO bond angle (deg)

0.004 0.007 0.010 0.0032 0.005 0.0057−0.0081 0.0032−0.0067

0.0040−0.0061 0.0029−0.0057

0.0047−0.0084

+0.018 to +0.027

0.0038−0.0053

3.043−3.393 3.233−3.446

+0.0192 +0.0280 +0.021 +0.018 +0.027 +0.0195 to +0.038

∇2ρ (au)

0.004 0.049 0.0046 0.0037 0.0054 0.0042−0.0079

ρbcp (au)

3.049 3.169 3.034 3.133 3.044 3.463−3.056

r (Å)

O···F

Table 3. Intermolecular Distances and Corresponding ρbcp and Laplacian ∇2ρ Parameters in [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) aggregates

The Journal of Physical Chemistry A Article


Article


Figure 5. Optimized structures of (a) [ETT][FEP]·2(CO2) and (b) [Hmim][FEP]·2(CO2) systems. Hydrogen bonding interactions are depicted as broken lines.

Figure 6. Optimized structures of (a, b) [ETT][FEP]·n(CO2) (n = 12, 16) and (c, d) [Hmim][FEP]·n(CO2) (n = 12, 16). Hydrogen bonding interactions are depicted as broken lines.

depicted in Figure 7b,d. The orientation of hexyl chain on the cation renders more “voids” generating large free volumes for absorption that allows the escape of CO2 molecules. The cooperativity between CO2 molecules was not noticed even for n = 30 in the case of [Hmim][FEP] systems (cf. Figure 7d. It is therefore, expected that ILs composed of [Hmim] and [FEP]

consequent to steric repulsions between absorbed molecules, and the onset of cooperative interactions between CO2 can be observed. The ion pair binding sites in the [ETT][FEP]· 20(CO2) aggregates are expended contrary to those in [Hmim]. To examine the efficiency of CO2 absorption and in particular the role of cation, we derive the optimized structures of the absorbed [ETT] and [Hmim] aggregates (n = 30) H


Article


Figure 7. Optimized structures of (a) [ETT][FEP]·n(CO2) (n = 20, 30) and (b) [Hmim][FEP]·n(CO2) (n = 20, 30).

Table 4. Zero Point Corrected Binding Energies (ΔE), Change in Free Energies (ΔG), Enthalpies (ΔH), and Entropies (ΔS) of [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) Complexes ΔE (kJ mol−1)

ΔG (kJ mol−1)

ΔH (kJ mol−1)

ΔS (J mol−1 K−1)

no. of CO2, n

[ETT][FEP]· n(CO2)

[Hmim][FEP]· n(CO2)

[ETT][FEP]· n(CO2)


[ETT][FEP]· n(CO2)


[ETT][FEP]· n(CO2)


1 2 3 4 5 6 8 10

346.0 355.9 385.7 397.1 412.5 429.8 458.1 481.9

369.2 395.4 401.3 418.3 456.8 472.6 488.3 544.9

−249.9 −228.3 −217.6 −197.2 −158.6 −144.1 −97.8 −60.4

−265.4 −247.3 −224.8 −199.8 −183.9 −168.3 −122.9 −62.1

−342.4 −350.1 −378.8 −387.6 −403.0 −418.1 −443.3 −463.6

−357.4 −379.1 −383.5 −396.6 −430.6 −440.9 −449.3 −501.3

−309.9 −408.3 −540.8 −638.7 −819.7 −919.2 −1158.6 −1352.4

−308.5 −441.9 −532.5 −660.1 −827.3 −914.3 −1094.7 −1473.1

pair favors addition of CO2 driven by the interplay of hydrogen and halogen bonding. 3.6. Binding Energies in [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) Systems. Calculated cation−anion binding energies for the [ETT][FEP] ion pair turn out to be 338.7 kJ mol−1, which is ∼15 kJ mol−1 lower than that for [Hmim][FEP]. To examine CO2 capture capacity of the ILs, the binding energies are compared in Table 4. It may be observed that for the [Hmim][FEP] system, the binding energy as a function of absorbed CO2 molecules (up to n = 30) continues to increase linearly, unlike that for [ETT][FEP] based ILs which emerges with a flat region at n ≈ 24 as depicted in Figure 8. Secondly, the slope of the curve clearly shows increment in binding energy and tends to remain nearly constant up to n = 30 for the [Hmim] based ILs. The higher

ions should serve better for CO2 capture than those with [ETT] cation. From above discussion, it may be concluded that the presence of dense σ-holes on fluorine centers and largely delocalized electron-rich surface around the anion results in the complex network of directional O···H, C···F, and O···F interactions on stepwise addition of CO2 in the [X][FEP]· n(CO2) (X= [ETT]+ or [Hmim]+, n = 2−6, 8, 10, 12, 16, 20, and 30) systems and emerge with their signatures as bcps in their MED topography. The ρbcp and corresponding ∇2ρ parameters along with the intermolecular distances are reported in Table 3. Electron-rich as well as electron deficient fluorine atoms in the CO2 absorbed ion pairs can be distinguished. An increase of electron deficient character of fluorines in the ion I


Article


polarizabilities in Figure 10b. Parallel inferences are drawn for the CO2 absorbed [Hmim][FEP] systems. The accompanying noncovalent interactions owing to the hydrogen and σ-holes in the extended network of the [ETT] and [Hmim] based ILs are analyzed. 3.7. Noncovalent Interactions Reduced Density Gradient and Energy Decomposition Analyses. To probe into molecular interactions stabilizing the CO2 absorbed ILs, we incorporate the noncovalent interactions reduced density gradient based on the bonding topologies described by the QTAIM theory. An extended network of halogen as well as hydrogen bonding and other dispersive interactions in [ETT] and [Hmim] based ILs are quantified through the NCI analyses. The NCI is characterized in terms of the electron density (ρ), its Laplacian (∇2ρ), and the reduced gradient of density (s). The ∇2ρ is decomposed into three eigenvalues λ1 + λ2 + λ3 (with λ1 ≤ λ2 ≤ λ3) located in its Hessian. The second component, λ2, provides insights accompanying the reorganization of electron density consequent to noncovalent interactions further distinguishing the bonding (λ2 < 0) features from those of nonbonding (λ2 > 0) ones. Such interactions are characterized in terms of attractive, moderately strong hydrogen bonds, repulsive, steric, and weak dispersion-type of interactions. The elegance of this method lies in elucidating the interactions in real-space and the graphical representation distinguishing various noncovalent interactions. The position, strength, and type of an interaction shown as RDG isosurface are represented on a blue−green−red (BGR) scale according to the values and sign of (λ2)ρ ranging from −0.05 to +0.05 au. The regions with positive (λ2)ρ area are portrayed as red indicating strong repulsive nonbonded overlap, while the attractive interactions are displayed in blue and green regions, which refer to very weak interactions. Figure 11 shows the bonding regions as NCI surfaces corresponding to the same RDG isovalue in [ETT][FEP]·n(CO2) absorbed aggregates (n = 1, 3, 6, 10, and 20). Emergence of NCI-RDG troughs observed in low electron density (−0.002 au) regions between atomic basins in ILs refer to noncovalent bonding sites. The isosurfaces of [ETT][FEP]·1(CO2) shown in green suggest the presence of C···F and O···F interactions between CO2 and anion in addition to the hydrogen bonding from the methyl group of the [ETT] cation (cf. Figure 11a). The region midway between the cation and anion characterizes strong hydrogen bonding depicted in blue. On the other hand, the isosurface shown in red stems from steric repulsions between fluorine atoms. NCI isosurfaces for [ETT][FEP]·n(CO2) are depicted in Figure 11a−e. A comparison of systems with n = 1, 3, and 6 absorbed CO2 molecules reveals similar binding regions originating from hydrogen and halogen bonding interactions. The steric repulsions between the fluorines decrease significantly for [ETT][FEP]·10(CO2), which reveals diminutive red regions as shown in Figure 11d. The relatively large contributions from the electrostatic interactions are reflected in the increase of “spattered” regions as green. The counterbalancing of these interactions results in the formation of “voids” facilitates the diffusion of CO2. With subsequent addition of CO2 molecules (n > 16), cooperative interactions between the absorbed CO2 molecules are observed. This is apparent from the NCI isosurface of [ETT][FEP]·20(CO2) shown in Figure 11e. The weakening of cation−anion interactions generates “free volume” between them and consequently the blue region, observed in [ETT][FEP]· 1(CO2), gradually depletes with the addition of CO2. The

Figure 8. Plot of binding energies as a function of number of CO2 molecules absorbed in [ETT][FEP] and [Hmim][FEP] ion pairs. See text for details.

efficiency of CO2 capture by [Hmim] based ILs agrees well with the conclusion based on the works of Zhang et al.32 The binding energies and thermodynamic parameters for CO2 absorbed ion pairs (up to n = 10) calculated from the vibrational frequency calculations are summarized in Table 4. As may be inferred, the CO2 absorption is enthalpy driven, and large exothermicities of Hmim based ILs suggest higher CO2 capture capacity. The change in ΔG parameters upon sequential addition of CO2 molecules further corroborates these inferences (cf. Figure 9).

Figure 9. Graph showing the change in free energies (−ΔG) versus number of CO2 molecules absorbed in [ETT][FEP] and [Hmim][FEP] ion pairs.

Murray et al.41 earlier pointed out that ΔE can be linearly correlated to the strength of interactions through Vs,max for σhole and the Vs,min of Lewis base through ΔE = α2[Vs,maxVs,min] + β2, with α2 and β2 determined from double regression analysis. The directional network of σ-holes and hydrogen bonds herein reveal Vs,max and Vs,min in their anisotropic electronic distributions. A correlation of binding energies to ESP topographical parameters in CO2 absorbed ion pairs hitherto studied is far from straightforward. Binding of CO2 to ILs is facilitated through electrostatic and intermolecular forces from charge assisted hydrogen and halogen bonding. Large Vs,max and Vs,min sites render a certain covalent character to the system as a result of the strong electric field by positive sites and weakly held highly polarizable negative ones. Further addition of CO2 increases the polarizability of the [ETT][FEP]·n(CO2) complexes (cf. Figure 10a) with an increase of its capacity to accommodate dative sharing from fluorine electronic charge. Binding energies are represented in terms of J


Article


Figure 10. A plot of (a, c) polarizabilities as a function of number of CO2 molecules absorbed and (b, d) binding energies of CO2 absorbed complexes versus polarizabilities for (a, b) [ETT][FEP]·n(CO2) complexes and (c, d) their [Hmim] analogues. See text for details.

the many-body effects on structure and binding in absorbed complexes suggests a linear dependence on the electrostatic interaction components with CO2 binding (cf. Figure 13). 3.8. Vibrational Spectra. Structural changes induced by σhole and hydrogen bonding interactions usually are accompanied by a shift in the frequency of the characteristic stretching vibrations in their spectra. The cation−anion binding and the influence of CO2 absorption can be probed through normal vibration analysis. Calculated vibrational spectra portraying the molar absorption coefficient (or molar absorptivity in units of 0.1 m2 mol−1) versus the frequency (in cm−1) of [ETT][FEP] are given in Supporting Information. The methyl vibrations invoke strong coupling from different internal coordinates and are assigned to 3244, 3151, and 3093 cm−1 bands. The P−F stretching emerges with a strong peak at the 823 cm−1. CO2 absorption is facilitated via halogen bonding with the anion and C−H···O hydrogen bonding with methyl group on the [ETT]+ cation. Thus, CO2 binding to the ion pair shifts the C−Ha vibrations (Ha protons refer to those participating in hydrogen bonding) to lower wavenumber (3200 cm−1), whereas the C− N vibration shows a marginal upshift of ∼3 cm−1 (cf. Table 5). It was pointed out earlier that interaction of CO2 with the anion is facilitated through O···F halogen bonding from the CF2 and CF3 functionalities of pentafluoroethyl group on the [FEP] anion. The interacting C−Fa vibration shows an upshift (blue-shift) from 1228 to 1281 cm−1 in the [ETT][FEP]· 1(CO2) complex. The intense 1241 cm−1 vibration of the ion pair was predicted at the 1263 cm−1 on absorption of CO2. As opposed to this the neighboring C−C stretching vibrations exhibit a “red shift”. Absorption of CO2 on the ion pair is

NCI-RDG plots underlying [Hmim][FEP]·n(CO2) systems have been provided in the Supporting Information. As an alternate perspective to real-space NCI isosurface for CO2 absorption, the reduced gradient of density is plotted as a function of electron density multiplied by the sign of λ2 for [ETT][FEP]·n(CO2) aggregates in Figure 12a−e. The O···H, O···F, and C···F interactions contribute to CO2 absorption. As shown in Figure 12a, low reduced gradients of electron density in the [ETT][FEP]·1(CO2) system reveal deep spikes indicating relatively strong electrostatic interactions. Besides steric interactions from [ETT], [FEP] and CO2 constituents spread over the higher electron density regions. A marked separation of the low and high electron density regions (in the neighborhood of 0.00 au) tends to reduce on traversing from n = 3 to 20, which is evident from Figure 12e. The crowding near the very low electron density region and its subsequent narrowing with increasing CO2 molecules can be viewed as interplay of the C···F, O···F, and hydrogen bonding. The σ-hole and hydrogen bonding in the [Hmim][FEP]·n(CO2) systems lead to similar conclusions. Noteworthy enough, the CO2 cooperative interactions were not noticed for the [Hmim][FEP]·20(CO 2) systems contrast to the [ETT][FEP]· 20(CO2). The above inferences are further supported by the energy decomposition analysis (EDA) in which the interaction energies were decomposed as Pauli repulsion, electrostatic, orbital, and dispersion energy terms toward the total interaction energy. Energy contributions were estimated from EDA. The electrostatic and dispersive interactions contribute to halogen and hydrogen bonding for large systems. A delicate balance of K


Article


Figure 11. Color-filled RDG isosurfaces depicting noncovalent interaction (NCI) regions in [ETT][FEP]·n(CO2) complexes (n = 1, 3, 6, 10, and 20). Green regions denote O···H, C···F, and O···F interactions, while the red isosurface refers to steric effects.

splitting of degenerate bending vibration leading to 687 and 698 cm−1 bands as observed from Table 5. Calculated infrared frequencies for the [ETT][FEP]·n(CO2) complexes are compared with those of the free ion pair in Table 5. With subsequent addition of CO2, the stretching frequencies of the interacting C−Fa (facilitating O···F halogen bonding interactions) emerge with their signature as intense bands in the 1280 to 1220 cm−1 region. Relatively pure C−C stretching vibrations occur in the 1380−1356 cm−1 region. The strong

evident from the presence of new intense vibrations. The free CO2 molecule possesses vibrations at 696 cm−1 (doubly degenerate angular deformation), 1413 cm−1 (symmetric stretching), and 2458 cm−1 (the asymmetric stretch band). Stretching vibrations of CO2 are nearly unchanged on interaction with the [ETT][FEP] ion pair. Consequent to such binding, the deviation of CO2 bond angle, which is crucial in governing CO2 capture, concomitantly engenders the L


Article


Figure 12. NCI index plot: plot of function 1 (sign(λ2) ρ values) on the x-axis versus function 2, the reduced density gradient (RDG), on the Y-axis for [ETT][FEP]·n(CO2) complexes (n = 1, 3, 6, 10, and 20).

Figure 13. Binding energies plotted against the electrostatic contributions (Eelect) toward CO2 binding. M


2458, 1413

1380, 1369, 1354 1281, 1263, 1228 823 687, 693

ν CO2

ν C−C

N

a

1648 1359−1385 1220−1282 824 691, 694

See text for details.

C−N C−C C−Fa P−Fa CO2

ν ν ν ν δ

3239, 3212, 3132 3213, 3157, 3138 2458, 2460, 1413, 1415 1648 1359−1384 1221−1288 831 687−697

3236, 3212 3132 3265, 3136

2458, 1414

3345

2462, 2464, 1415, 1415 1380, 1371, 1356 1287, 1265, 1240 820 690−697

3244, 3190, 3089

2CO2

3393

ν C−Ha ring proton ν C−Ha methyl ν C−H2a alkyl ν CO2

ν P−Fa δ CO2

ν C−Fa

3255, 3207, 3182

1CO2

ν C−Ha

assignments

1654 1348 1222−1279 812 684−698

2456−2469, 1413−1418

3214, 3203, 3116, 3150, 3139, 3073 3176, 3157, 3112, 3070

3367

813 673−691

1224−1280

1365−1383

2452−2459, 1411−1415

3212, 3190, 3154

3CO2

6CO2

1653 1350−1377 1223−1278 815 689−697

1220−1276

1362, 1385

1651 1357−1384 1222−1281 823 676−693

2454−2464, 1414−1415

3249, 3227, 3141, 3157, 3148, 3084 3202, 3135

800 681−698 (b) [Hmim][FEP] 3391

2452−2467, 1414−1417

3135, 3061

3213, 3191, 3108

3363

823 688−694

1240−1281

1371, 1381

(a) [ETT][FEP] 3225, 3182, 3179, 3168, 3224, 3178, 3085, 3233, 3200, 3161, 3091 3163, 3166, 3150, 3086, 3224, 3144, 3081, 2452−2465, 1413−1415 2457−2468, 1413−1419

4CO2

1646 1359−1384 1222−1289 821 686−696

2451−2465, 1413−1416

3132, 3064, 3113, 3071

3244, 3232, 3145, 3163, 3147, 3084

3381

808 683−691

1216−1261

1363, 1377, 1384

3258, 3189, 3096, 3242, 3145, 3082, 3207, 3174, 3102, 3239, 3167, 3152, 3164, 3086 2455−2470, 1413−1419

8CO2

1648 1349−1379 1224−1281 815 683−699

2451−2472, 1413−1424

3196, 3111, 3052, 3122

3230, 3211, 3124, 3150, 3134, 3068

3363

808 679−698

1215−1272

1363, 1375, 1384

3213, 3153, 3088, 3259, 3206, 3103, 3240, 3135, 3074, 3217, 3179, 3107, 3254, 3181, 3169, 3091 2451−2467, 1411−1417

10CO2

Table 5. Comparison of Selected Vibrational Frequencies (ν = Stretching and δ = Bending) in [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) Complexesa

The Journal of Physical Chemistry A Article


Article


Figure 14. A plot of binding energies for (a) [ETT][FEP]·n(CO2) and (b) [Hmim][FEP]·n(CO2) complexes as a function of the separation (Δν in cm−1) of degenerate CO2 bending vibrations.

Figure 15. Difference MED maps for the (a) [ETT][FEP]·1(CO2) and (b) [Hmim][FEP]·1(CO2) complexes in the plane passing through CO2 molecule and interacting C−Ha and C−Fa bonds. Contours in the range of ±0.001 to ±0.009 au are shown. The bcps in the MED topography are also given.

intense band ∼820 cm−1 assigned to the P−F stretching observed in the [ETT][FEP] ion pair and their CO2 absorbed complexes remains unchanged. The O···F halogen bonding interactions reflect distinct vibrational features from those arising as a result of cation hydrogen bonding. The hydrogen bonding interactions lead to the downshift in frequency of C−Ha stretching bringing about the concomitant upshift of vicinal C−N vibration. On the other hand, the O···F interactions engender frequency “blue shift” for C−F a stretching accompanied by a downshift of its corresponding C−C covalent counterpart. Thus, hydrogen and halogen bonding reveal complementary vibrational characteristics. As noticed for [ETT][EFP]·1(CO2) complex, the symmetric and asymmetric stretching vibrations of the isolated CO2 are nearly unchanged on subsequent addition of CO2 to the ion pair. Noteworthy enough, the interaction of CO2 brings about a deviation of the COC bond angle accompanied by splitting of the degenerate bending vibration. A separation (Δν) of these vibrations further can be correlated to binding energies of complexes until CO2 cooperative interactions occur. A plot of binding energies of the CO2 absorbed [ETT][EFP]·n(CO2) complex as a function of Δν is

shown in Figure 14. Similar observations can be made for the [Hmim] based ILs. The reorganization of molecular electron density engenders accumulation or depletion of electron density on CO 2 absorption by the ion pair. The difference MED plot employed herein sheds light on the strengthening (or weakening) of the bond and explains the direction of frequency shift of corresponding vibrations in the calculated infrared spectra. To envisage this, difference MED contours in the plane passing through the ion pair containing interacting C−Ha and C−Fa bonds (those participating in O···H and O···F or C···F interactions) and the CO2 molecule were considered. The Δρ contours from ±0.001 to ±0.009 au for [ETT][FEP]· 1(CO2) and [Hmim][FEP]·1(CO2) complexes are displayed in Figure 15a,b. Here blue lines represent positive valued contour, while the negative valued contours appear in red. A zero valued contour is depicted as green. As may be noticed from Figure 15a, the bcp of the C−Ha bond with methyl proton interacting with the CO2 molecule appears in the region (red) where the electron density is depleted as a result of hydrogen bonding. The bond weakening is also apparent from the difference MED contour plots. The frequency down-shift of the C−H a stretching vibrations in [ETT][FEP]·1(CO2) complex thus is O


Article

The Journal of Physical Chemistry A evident. As opposed to this, the C−Fa bonds emerge with bcps in enhanced electron density (blue) region implying increased bond strength for the vibrations which is attributed to halogen bonding. The corresponding vibration shifts to higher wavenumber (1281 cm−1) compared with the free ion pair (1228 cm−1). The [ETT][FEP]·1(CO2) complex further reveals depletion of electron density in the vicinity of CF2/CF3 groups, and explains frequency down-shift of C−C stretching vibration of the pentafluoroethyl substituent. These arguments further can be extended to rationalize the direction of frequency shifts in the infrared spectra of [Hmim][FEP]·1(CO2) complex. Besides the occupancies of σ* (C−C), σ*(C−F), and σ* (C−H) antibonding natural orbitals reported in Table 2 further corroborate the inferences on the bond strengthening (or weakening) in [ETT][FEP]·1(CO2) and [Hmim][FEP]· 1(CO2) complexes.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Fax no.: +91-20225691728. Telephone no.: +91 020 25601225. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS S.P.G. acknowledges support from the Research Project (37(2)/14/11/2015-BRNS) from the Board of Research in Nuclear Sciences (BRNS), India. S.S.R. is grateful to UGC for award of the meritorious research fellowship. Authors thank the Center for Development of Advanced Computing (CDAC), Pune, for providing National Param Supercomputing Facility.

■

CONCLUSIONS Electronic structure, binding energies, CO2 capture capacity, and vibrational characteristics of [ETT][FEP]·n(CO2) and [Hmim][FEP]·n(CO2) based ILs have been carried out employing the dispersion corrected M05-2x/6-311++G(d,p) density functional theory. The fluorine atoms engendering anisotropic distribution of ESP around the ion pair dictate the CO2 binding as a result of combined effect from head-on O···F halogen bonding interactions, electrostatic C···F interactions from the anion, and O···H hydrogen bonding with the cation. The CO2 absorption is enthalpy driven. Binding energies, free energies, and polarizabilities correlate well with the absorbed CO2 molecules. The σ-hole interactions reveal their signature as the intense blue-shifted C−F vibrations (1280 to 1220 cm−1) relative to the free ion pair. The C−H vibrations show a shift in the opposite direction on absorption of CO2 owing to hydrogen bonding, and thus σ-hole and hydrogen bonding interactions exhibit complementary characteristics for the frequency shifts. Moreover, the binding of CO2 to the ion pair lifts the degeneracy of bending vibration (at 677 cm−1) and consequent separation (Δν in cm−1) in the absorbed complexes can be correlated to their binding energies. Electron-rich regions of the anion, σ-holes, and the absorbed molecules render free volumes for further binding with increasing CO2 molecules. This is evident from the spattered electrostatic basins extending over the ion pair in NCI plots. Further the cooperativity between absorbed CO2 molecules was not observed until n = 30 for the [Hmim][FEP] system, whereas the [ETT][FEP] complex reveal emergence of cooperative interactions with n = 16 CO2 molecules. The change in binding energy with absorbed CO2 molecules demonstrates the efficiency of [Hmim] based ILs over [ETT], as also evidenced from their large exothermicities. Thus, interactions from the cation become crucial in the assessment of the overall performance of ILs with increasing CO2 absorption. To summarize, the present endeavor underlines the nature and manifestations of hydrogen and σ-hole interactions in structure, binding energies, and spectral characteristics, which should prove useful in designing TSILs for CO2 capture.

■

Optimized geometries, structural parameters, vibrational frequencies, and NBO parameters in ion pair conformers (PDF)

■

REFERENCES

(1) Herzog, H. F. Scaling up Carbon dioxide Capture and Storage: From Megatons to Gigatons. Energy Econ. 2011, 33, 597−604. (2) Gough, C. State of the Art in Carbon Dioxide Capture and Storage in the UK: An experts’ review. Int. J. Greenhouse Gas Control 2008, 2, 155−168. (3) Figueroa, J. D.; Fout, T.; Plasynski, S.; McIlvried, H.; Srivastava, R. D. Advances in CO2 Capture Technology-The U.S. Department of Energy’s Carbon Sequestration Program. Int. J. Greenhouse Gas Control 2008, 2, 9−20. (4) Jones, C. W. CO(2) Capture from Dilute Gases as a Component of Modern Global Carbon Management. Annu. Rev. Chem. Biomol. Eng. 2011, 2, 31−52. (5) MacDowell, N.; Florin, N.; Buchard, A.; Hallett, J.; Galindo, A.; Jackson, G.; Adjiman, C. S.; Williams, C.; Shah, N.; Fennell, P. An Overview of CO2 Capture Technologies. Energy Environ. Sci. 2010, 3 (11), 1645−1669. (6) D’Alessandro, D. M.; Smit, B.; Long, J. R. Carbon Dioxide Capture: Prospects for New Materials. Angew. Chem., Int. Ed. 2010, 49, 6058−6082. (7) Rochelle, G. T. Amine Scrubbing for CO2 Capture. Science 2009, 325, 1652−1654. (8) Vaidya, P. D.; Kenig, E. Y. CO2-Alkanolamine Reaction Kinetics: A Review of Recent Studies. Chem. Eng. Technol. 2007, 30, 1467− 1474. (9) Working group III of the IPCC. IPCC Special Report on Carbon Capture and Storage Inter Governmental Panel on Climate Change, Ganeva, Switzerland, 2005. (10) Blanchard, L. A.; Hancu, D.; Beckman, E. J.; Brennecke, J. F. Green Processing Using Ionic Liquids and CO2. Nature 1999, 399, 28−29. (11) Anthony, J. L.; Maginn, E. J.; Brennecke, J. F. Solubilities and Thermodynamic Properties of Gases in the Ionic Liquid 1-n-Butyl-3Methylimidazolium Hexafluorophosphate. J. Phys. Chem. B 2002, 106, 7315−7320. (12) Ramdin, M.; de Loos, T. W.; Vlugt, T. J. H. State-of-the-Art of CO2 Capture with Ionic Liquids. Ind. Eng. Chem. Res. 2012, 51, 8149− 8177. (13) Cadena, C.; Anthony, J. L.; Shah, J. K.; Morrow, T. I.; Brennecke, J. F.; Maginn, E. J. Why Is CO2 so Soluble in ImidazoliumBased Ionic Liquids? J. Am. Chem. Soc. 2004, 126, 5300−5308. (14) Blanchard, L. A.; Gu, Z.; Brennecke, J. F. High-Pressure Phase Behaviour of Ionic Liquid/CO2 Systems. J. Phys. Chem. B 2001, 105 (12), 2437−2444. (15) Freemantle, M. An Introduction to Ionic Liquids; RSC publishing: London, 2010.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.5b12161. P


Article

The Journal of Physical Chemistry A (16) Rogers, R.; Seddon, K. Chemistry. Ionic liquids–Solvents of the Future? Science 2003, 302, 792−793. (17) Davis, J. H., Jr. Task-Specific Ionic Liquids. Chem. Lett. 2004, 33, 1072−1077. (18) Bates, E. D.; Mayton, R. D.; Ntai, I.; Davis, J. H. CO2 Capture by a Task-Specific Ionic Liquid. J. Am. Chem. Soc. 2002, 124, 926−927. (19) Lei, Z.; Dai, C.; Chen, B. Gas Solubility in Ionic Liquids. Chem. Rev. 2014, 114, 1289−1326. (20) Vaidya, P. D.; Kenig, E. Y. CO2-Alkanolamine Reaction Kinetics: A Review of Recent Studies. Chem. Eng. Technol. 2007, 30, 1467−1474. (21) Camper, D.; Bara, J.; Gin, D. L.; Noble, R. Room-Temperature Ionic Liquid-Amine Solutions: Tunable Solvents for Efficient and Reversible Capture of CO2. Ind. Eng. Chem. Res. 2008, 47, 8496−8498. (22) Goodrich, B. F.; de la Fuente, J. C.; Gurkan, B. E.; Zadigian, D. J.; Price, E. A.; Huang, Y.; Brennecke, J. F. Experimental Measurements of Amine-Functionalized Anion-Tethered Ionic Liquids with Carbon Dioxide. Ind. Eng. Chem. Res. 2011, 50, 111−118. (23) Zhang, J. M.; Zhang, S. J.; Dong, K.; Shen, Y. Q.; Lv, X. M.; Zhang, Y. Supported Absorption of CO2 by Tetrabutylphosphonium Amino Acid Ionic Liquids. Chem. - Eur. J. 2006, 12, 4021−4026. (24) Zhao, Y.; Zhang, X.; Dong, H.; Zhen, Y.; Li, G.; Zeng, S.; Zhang, S. Solubilities of Gases in Novel Alcamines Ionic Liquid 2-[2hydroxyethyl (methyl) amino] Ethanol Chloride. Fluid Phase Equilib. 2011, 302, 60−64. (25) Kanakubo, M.; Umecky, T.; Hiejima, Y.; Aizawa, T.; Nanjo, H.; Kameda, Y. Solution Structures of 1-butyl-3-methylimidazolium hexafluorophosphate Ionic Liquid Saturated With CO2: Experimental Evidence of Specific Anion-CO2 Interaction. J. Phys. Chem. B 2005, 109, 13847. (26) Aki, S. N. V. K.; Mellein, B. R.; Saurer, E. M.; Brennecke, J. F. High-Pressure Phase Behavior of Carbon Dioxide with ImidazoliumBased Ionic Liquids. J. Phys. Chem. B 2004, 108, 20355. (27) Wilhelm, E.; Battino, R. Thermodynamic Functions of the Solubilities of Gases in Liquids at 25. deg. Chem. Rev. 1973, 73, 1−9. (28) Zhang, X.; Huo, F.; Liu, Z.; Wang, W.; Shi, W.; Maginn, E. J. Absorption of CO2 in the Ionic Liquid 1-n-Hexyl-3-methylimidazolium Tris(pentafluoroethyl) trifluorophosphate ([hmim][FEP]): A Molecular View by Computer Simulations. J. Phys. Chem. B 2009, 113, 7591−7598. (29) Pringle, J. M.; Golding, J.; Baranyai, K.; Forsyth, C.; Deacon, G. B.; Scott, J. L.; MacFarlane, D. R. The Effect of Anion Fluorination in Ionic LiquidsPhysical Properties of a Range of Bis (methanesulfonyl) Amide Salts. New J. Chem. 2003, 27, 1504. (30) Yokozeki, A.; Shiflett, M. B.; Junk, C. P.; Grieco, L. M.; Foo, T. Physical and Chemical Absorptions of Carbon Dioxide in RoomTemperature Ionic Liquids. J. Phys. Chem. B 2008, 112, 16654. (31) Muldoon, M. J.; Aki, S. N. V. K.; Anderson, J. L.; Dixon, J. K.; Brennecke, J. F. Improving Carbon Dioxide Solubility in Ionic Liquids. J. Phys. Chem. B 2007, 111, 9001. (32) Zhang, X.; Liu, Z.; Wang, W. Screening of Ionic Liquids to Capture CO2 by COSMO-RS and Experiments. AIChE J. 2008, 54, 2717. (33) Raveendran, P.; Wallen, S. L. Exploring CO2-Philicity: Effects of Stepwise Fluorination. J. Phys. Chem. B 2003, 107, 1473. (34) Ignat’ev, N. V.; Biermann, U. W.; Kucheryna, A.; Bissky, G.; Willner, H. New Ionic Liquids With Tris(perfluoroalkyl)trifluorophosphate (FAP) Anions. J. Fluorine Chem. 2005, 126, 1150−1159. (35) Zhu, X.; Lu, Y.; Peng, C.; Hu, J.; Liu, H.; Hu, Y. Halogen Bonding Interactions Between Brominated Ion Pairs and CO2Molecules: Implications for Design of New and Efficient Ionic Liquids for CO2 Absorption. J. Phys. Chem. B 2011, 115, 3949−3958. (36) Murray-Rust, P.; Stallings, W. C.; Monti, C. T.; Preston, R. K.; Glusker, J. P. Intermolecular Interactions of the Carbon-Fluorine Bond: the Crystallographic Environment of Fluorinated Carboxylic Acids and Related Structures. J. Am. Chem. Soc. 1983, 105, 3206−3214.

(37) Brinck, T.; Murray, J. S.; Politzer, P. Surface Electrostatic Potentials of Halogenated Methanes as Indicators of Directional Intermolecular Interactions. Int. J. Quantum Chem. 1992, 44, 57−64. (38) Auffinger, P.; Hays, F. A.; Westhof, E.; Ho, P. S. Halogen Bonds in Biological Molecules. Proc. Natl. Acad. Sci. U. S. A. 2004, 101, 16789−16794. (39) Politzer, P.; Lane, P.; Concha, M. C.; Ma, Y.; Murray, J. S. An Overview of Halogen Bonding. J. Mol. Model. 2007, 13, 305−311. (40) Politzer, P.; Murray, J. S.; Clark, T. Halogen Bonding: An Electrostatically-Driven Highly Directional Noncovalent Interaction. Phys. Chem. Chem. Phys. 2010, 12, 7748−7757. (41) Politzer, P.; Murray, J. S.; Clark, T. Halogen Bonding and Other σ-Hole Interactions: A Perspective. Phys. Chem. Chem. Phys. 2013, 15, 11178. (42) Johnson, E. R.; Keinan, S.; Mori-Sanchez, P.; Contreras-García, J.; Cohen, A. J.; Yang, W. Revealing Noncovalent Interactions. J. Am. Chem. Soc. 2010, 132, 6498−6506. (43) Contreras- García, J.; Johnson, E. R.; Keinan, S.; Chaudret, R.; Piquemal, J.-P.; Beratan, D. N.; Yang, W. NCIPLOT: A Program for Plotting Noncovalent Interaction Regions. J. Chem. Theory Comput. 2011, 7, 625−632. (44) Tomasi, J.; Persico, M. Molecular Interactions in Solution: An Overview of Methods Based on Continuous Distributions of the Solvent. Chem. Rev. 1994, 94, 2027−2094. (45) Naray-Szabo, G.; Ferenczy, G. G. Molecular Electrostatics. Chem. Rev. 1995, 95, 829−847. (46) Gadre, S. R.; Shirsat, R. N. Electrostatics of Atoms and Molecules; Universities Press: Hyderabad, 2000. (47) Becke, A. D. Density-Functional Exchange-Energy Approximation with Correct Asymptotic Behavior. Phys. Rev. A: At., Mol., Opt. Phys. 1988, 38, 3098. (48) Lee, C.; Yang, W.; Parr, R. G. Development of the Colle-Salvetti Correlation-Energy Formula into a Functional of the Electron Density. Phys. Rev. B: Condens. Matter Mater. Phys. 1988, 37, 785. (49) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, O.; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09; Gaussian, Inc.: Wallingford, CT, 2009. (50) Zhao, Y.; Truhlar, D. G. Density Functionals with Broad Applicability in Chemistry. Acc. Chem. Res. 2008, 41, 157. (51) Zhao, Y.; Truhlar, D. G. Benchmark Databases for Nonbonded Interactions and their Use to Test Density Functional Theory. J. Chem. Theory Comput. 2005, 1, 415−432. (52) Zhao, Y.; Schultz, N. E.; Truhlar, D. G. Design of Density Functionals by Combining the Method of Constraint Satisfaction with Parametrization for Thermochemistry, Thermochemical Kinetics, and Noncovalent Interactions. J. Chem. Theory Comput. 2006, 2, 364. (53) Zhao, Y.; Truhlar, D. G. A New Local Density Functional for Main-Group Thermochemistry, Transition Metal Bonding, Thermochemical Kinetics and Noncovalent Interactions. J. Chem. Phys. 2006, 125, 194101. (54) Izgorodina, E. I.; Bernard, U. L.; MacFarlane, D. R. Ion-Pair Binding Energies of Ionic Liquids: can DFT Compete with Ab InitioBased Methods? J. Phys. Chem. A 2009, 113, 7064. (55) Zahn, S.; MacFarlane, D. R.; Izgorodina, E. I. Assessment of Kohn-Sham Density Functional Theory and M?ller-Plesset PerturbaQ


Article


(77) Murray, J. S.; Lane, P.; Politzer, P. Expansion of the σ-Hole Concept. J. Mol. Model. 2009, 15, 723−729. (78) Riley, K. E.; Murray, J. S.; Fanfrlík, J.; Rézác,́ J.; Solá, R. J.; Concha, M. C.; Ramos, F. M.; Politzer, P. Halogen Bond Tunability I: The Effects of Aromatic Fluorine Substitution on the Strengths of Halogen-Bonding Interactions Involving Chlorine, Bromine, And Iodine. J. Mol. Model. 2011, 17, 3309−3318. (79) Metrangolo, P.; Murray, J. S.; Pilati, T.; Politzer, P.; Resnati, G.; Terraneo, G. Fluorine-Centered Halogen Bonding: A Factor in Recognition Phenomena and Reactivity. Cryst. Growth Des. 2011, 11, 4238−4246. (80) Pal, R.; Nagendra, G.; Reddy, S. M.; Sureshbabu, V. V.; Guru Row, T. N. Observation of a Reversible Isomorphous Phase Transition and An Interplay of “σ-Holes” and “π-Holes” in Fmoc-Leu-ψ [CH 2NCS]. Chem. Commun. 2015, 51, 933−936. (81) Alabugin, I. V.; Manoharan, M.; Peabody, S.; Weinhold, F. Electronic Basis of Improper Hydrogen Bonding: A Subtle Balance of Hyperconjugation and Rehybridization. J. Am. Chem. Soc. 2003, 125, 5973. (82) Gupta, K. M.; Jiang, J. Systematic Investigation of Nitrile Based Ionic Liquids for CO2 Capture: A Combination of Molecular Simulation and Ab Initio Calculation. J. Phys. Chem. C 2014, 118, 3110−3118. (83) Koch, U.; Popelier, P. L. A. Characterization of C-H-O Hydrogen Bonds on the Basis of the Charge Density. J. Phys. Chem. 1995, 99 (24), 9747−9754. (84) Grabowski, S. J. Halogen Bond and its Counterparts: Bent’s Rule Explains the Formation of Nonbonding Interactions. J. Phys. Chem. A 2011, 115, 12340−12347. (85) Bent, H. A. An Appraisal of Valence-Bond Structures and Hybridization in Compounds of the First-Row Elements. Chem. Rev. 1961, 61, 275.

tion Theory for Ionic Liquids. Phys. Chem. Chem. Phys. 2013, 15, 13664. (56) Dennington, R.; Keith, T.; Milliam, J. GaussView, version 5; Semichem Inc.: Shawnee Mission, KS, 2009. (57) Matta, C. F.; Boyd, R. J. The Quantum theory of Atoms in Molecules. In An Introduction to the Quantum Theory of Atoms in Molecules; Matta, C. F., Boyd, R. J., Eds.; Wiley-VCH: Weinheim, Germany, 2007. (58) Keith, T. A. AIMAll, version 14.11.23; TK Gristmill Software: Overland Park, KS, 2014 (aim.tkgristmill.com). (59) Foster, J. P.; Weinhold, F. Natural Hybrid Orbitals. J. Am. Chem. Soc. 1980, 102, 7211−18. (60) Kirchner, B. In Ionic Liquids; Kirchner, B., Ed.; Springer-Verlag: Berlin/Heidelberg, 2010; Vol. 290, pp 213−262. (61) Koßmann, S.; Thar, J.; Kirchner, B.; Hunt, P. A.; Welton, T. Cooperativity in Ionic Liquids. J. Chem. Phys. 2006, 124, 174506. (62) Hunt, P. A.; Kirchner, B.; Welton, T. Characterizing the Electronic Structure of Ionic Liquids: an Examination of the 1-Butyl-3methylimidazolium Chloride Ion Pair. Chem. - Eur. J. 2006, 12, 6762− 6675. (63) Yan, X. Q.; Zhao, X. R.; Wang, H.; Jin, W. J. The Competition of σ-Hole···Cl− And Π-Hole···Cl− Bonds Between C6F5X (X = F, Cl, Br, I) and the Chloride Anion and its Potential Application in Separation Science. J. J. Phys. Chem. B 2014, 118, 1080−1087. (64) Hardegger, L. A.; Kuhn, B.; Spinnler, B.; Anselm, L.; Ecabert, R.; Stihle, M.; Gsell, B.; Thoma, R.; Diez, J.; Benz, J.; Plancher, J. M.; Hartmann, G.; Banner, D. W.; Haap, W.; Diederich, F. Systematic Investigation of Halogen Bonding in Protein-Ligand Interactions. Angew. Chem., Int. Ed. 2011, 50, 314−318. (65) Wilcken, R.; Zimmermann, M. O.; Lange, A.; Joerger, A. C.; Boeckler, F. M. Principles and Applications of Halogen Bonding in Medicinal Chemistry and Chemical Biology. J. Med. Chem. 2013, 56, 1363−1388. (66) Metrangolo, P.; Neukirch, H.; Pilati, T.; Resnati, G. Halogen Bonding Based Recognition Processes: A World Parallel to Hydrogen Bonding. Acc. Chem. Res. 2005, 38, 386−395. (67) Aakeröy, C. B.; Fasulo, M.; Schultheiss, N.; Desper, J.; Moore, C. Structural Competition Between Hydrogen Bonds and Halogen Bonds. J. Am. Chem. Soc. 2007, 129, 13772−13773. (68) Metrangolo, P.; Meyer, F.; Pilati, T.; Resnati, G.; Terraneo, G. Halogen Bonding in Supramolecular Chemistry. Angew. Chem., Int. Ed. 2008, 47, 6114−6127. (69) Metrangolo, P.; Carcenac, Y.; Lahtinen, M.; Pilati, T.; Rissanen, K.; Vij, A.; Resnati, G. Nonporous Organic Solids Capable of Dynamically Resolving Mixtures of Diiodoperfluoroalkanes. Science 2009, 323, 1461−1464. (70) Lu, Y.; Zou, J.; Wang, H.; Yu, Q.; Zhang, H.; Jiang, Y. Triangular halogen trimers. A DFT Study of the Structure, Cooperativity and Vibrational Properties. J. Phys. Chem. A 2005, 109, 11956−11961. (71) Lu, Y.-X.; Zou, J.-W.; Wang, Y.-H.; Jiang, Y.-J.; Yu, Q.-S. Ab Initio Investigation of the Complexes Between Bromobenzene And Several Electron Donors: Some Insights Into the Magnitude and Nature of Halogen Bonding Interactions. J. Phys. Chem. A 2007, 111, 10781−10788. (72) Lu, Y.; Wang, Y.; Xu, Z.; Yan, X.; Luo, X.; Jiang, H.; Zhu, W. C− X··· H Contacts in Biomolecular Systems: How They Contribute to Protein− Ligand Binding Affinity. J. Phys. Chem. B 2009, 113, 12615− 12621. (73) Lu, Y.; Shi, T.; Wang, Y.; Yang, H.; Yan, X.; Luo, X.; Jiang, H.; Zhu, W. Halogen Bonding-A Novel Interaction for Rational Drug Design? J. Med. Chem. 2009, 52, 2854−2862. (74) Lu, Y.; Wang, Y.; Zhu, W. Nonbonding Interactions Of Organic Halogens In Biological Systems: Implications For Drug Discovery And Biomolecular Design. Phys. Chem. Chem. Phys. 2010, 12, 4543−4551. (75) Clark, T.; Henneman, M.; Murray, J. S.; Politzer, P. Halogen Bonding: The σ-hole. J. Mol. Model. 2007, 13, 291−296. (76) Murray, J. S.; Lane, P.; Clark, T.; Riley, K. E.; Politzer, P. σHoles, π-Holes and Electrostatically-Driven Interactions. J. Mol. Model. 2012, 18, 541−548. R


CO2 Absorption Using Fluorine Functionalized Ionic Liquids: Interplay

Recommend Documents