Thermodynamics of Ionic Liquid Cosolvent Mixtures Using Molecular

Article Cite This: J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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Thermodynamics of Ionic Liquid Cosolvent Mixtures Using Molecular Dynamics Simulation: 1‑Ethyl-3-methylimidazolium Acetate Santosh Rathan Paul Bandlamudi† and Kenneth M. Benjamin*,† †

Department of Chemical and Biological Engineering, South Dakota School of Mines and Technology, Rapid City, South Dakota 57701, United States S Supporting Information *

ABSTRACT: Molecular dynamics (MD) simulations were conducted to calculate the binary mixture properties (excess molar heats and volumes) of 1-ethyl-3-methylimidazolium acetate [EMIM][Ac] with six polar covalent molecules (acetic acid, acetone, chloroform, dimethyl sulfoxide, isopropyl alcohol, and methanol) over the entire composition range. All of the binary mixtures of [EMIM][Ac] with cosolvents reported negative volumes of mixing over the entire composition range while releasing heat (exothermic heats of mixing), except for the acetone mixture, which had miscibility issues at higher concentrations (xacetone ≥ 0.6). These results were attributed to the nature, size, and structure of the molecules as well as the nature of the interaction between the mixture components. Further, aggregation of the ionic liquid (IL) at higher concentrations of the cosolvent added to the above observed effects, except for the acetic acid mixture. These results have been interpreted in terms of hydrogen-bonding and intermolecular interactions among the binary mixture components via radial distribution functions. The results showed that hydrogen-bonding interactions between the IL and cosolvent play a major role in the volumetric and energetic properties of the mixtures. In the case of the acetic acid mixture, aggregation of acetic acid molecules is reported at lower concentrations of acetic acid (xacetic acid = 0.1), while [EMIM][Ac] molecules were completely solvated at higher concentrations of acetic acid (xacetic acid = 0.9). Mixtures of [EMIM][Ac] with methanol showed isolated methanol molecules at lower concentrations of methanol (xmethanol = 0.1), whereas IL ions aggregated to form pairs at higher concentrations of methanol.

1. INTRODUCTION Ionic liquids (ILs) are a class of salts composed of a larger organic cation and smaller organic or inorganic anion that remain as liquids at ambient temperatures. Apart from being expensive, they have been widely recognized for their unique and versatile physiochemical properties, such as high thermal stability, ionic conductivity, and so on. Since the properties of ILs depend on the specific combination of cation and anion, the availability of a large number of possible combinations of cations and anions results in different properties of ILs,1,2 and hence, they are termed as designer solvents.3 Also, ILs are considered green solvents4 because of their negligible vapor pressures. ILs are used in a diverse number of applications such as extraction of biologically active compounds,5 catalysis,6 chemical processing,7 supercritical fluid reactions8 and separations,9,10 lubricants,11 and petroleum industries.12 Since ILs are highly viscous, their direct use may sometimes hinder the technical and cost efficiencies of the applications. Mixtures of ILs with polar covalent molecules (PCMs) display reduced viscosities, which can sometimes enhance these processing efficiencies. This has been effectively demonstrated especially in the case of extraction of oil from lipid-bearing biomass as well as autopartitioning of extracted lipids into their own separate phase.13,14 Similarly, in liquid−liquid extractions the presence © XXXX American Chemical Society

of a cosolvent in ILs has been proven to enhance the extraction efficiency.15,16 In this study, molecular dynamics (MD) simulations are employed to calculate the heats and volumes of mixing for mixtures of 1-ethyl-3-methylimidazolium acetate [EMIM][Ac] with the PCMs methanol, acetone, acetic acid, propanol, dimethyl sulfoxide (DMSO), and chloroform. [EMIM][Ac] has been effectively used in various applications such as biomass processing17,18 and CO2/N2 and CO2/H2 separations19,20 (because of the high solubility of CO2) and as a cosolvent for bio catalysis.21 Also, it is a good solvent for biomacromolecules such as cellulose,22 uracil,23 and DNA and proteins24 because of the strong hydrogen-bond-accepting ability of acetate anion.25−27 Given the wide range of applications involving ILs and cosolvents, a good understanding of binary mixture properties, such as excess heats and volumes of mixing, along with an understanding of molecular structure and the nature of interactions in solution is both necessary and the focus of this work. These excess thermodynamic mixture Special Issue: Emerging Investigators Received: November 30, 2017 Accepted: April 17, 2018

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DOI: 10.1021/acs.jced.7b01041 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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properties are required for chemical processing and process design. Although there have been several experimental and simulation studies on other imidazolium-based ILs and cosolvents prior to this study,28−36 studies on these particular [EMIM][Ac] cosolvent systems have been sparse.27 For ideal mixtures, the net enthalpies and volumes of mixing both equal zero, but this does not hold true for real mixtures, where nonzero heats and volumes of mixing can be attributed to the structural changes as a result of intermolecular interactions such as hydrogen bonding and electrostatics. Moreover, these excess thermodynamic properties are useful in predicting the extent of complex (aggregation) formation, or association, in binary mixtures.

+ k 3[1 + cos(3Ø)] + k4[1 − cos(4Ø)]

(4)

kφ(φ − φ0)2

∑

E improper =

(5)

impropers

⎧ ⎡⎛ ⎞12 ⎛ ⎞6 ⎤⎫ σij σij ⎪ ⎪ ∑ ⎨4ϵij⎢⎢⎜⎜ ⎟⎟ − ⎜⎜ ⎟⎟ ⎥⎥⎬ r ⎝ rij ⎠ ⎦⎪ j>i ⎪ ⎭ ⎩ ⎣⎝ ij ⎠

N−1 N

Evdw =

∑ i=1

Eqq =

(6)

qiqj 4ϵ0rij

(7)

∑

Etorsion/dihedral =

k Ø[1 + cos(nØ − δ)]

dihedrals

(8)

All of the symbols in the above equations have their conventional meanings.50 In addition, the net charge on an IL pair is zero (i.e., the pair is electrically neutral). For all of the atoms, nonbonded interactions are given by pairwise-additive Lennard-Jones (LJ) 12−6 potentials (eq 6) and Coulombic interaction of charges (eq 7), where ϵij is the LJ well depth, rij is interatomic separation distance, σij is the distance at which the interatomic particle potential is zero, and qi and qj are partial charges of atom sites i and j, respectively. 2.2. Mixture Thermodynamic Properties. In this study, the excess volume was calculated as ΔVmix = ⟨Vmix⟩ − [x IL⟨VIL⟩ + x PCM⟨VPCM⟩]

(9)

where xIL and xPCM are the mole fractions of the IL and PCM, respectively, ⟨Vmix⟩ is the ensemble-averaged volume of the mixture, and ⟨VIL⟩ and ⟨VPCM⟩ are the ensemble-averaged volumes of the pure IL and PCM, respectively. The enthalpy of mixing was calculated using the formula Δhmix = ⟨Umix ⟩ −

∑ xi⟨Ui⟩ + P(⟨Vmix⟩ − ∑ xi⟨Vi ⟩) (10)

where ⟨Umix⟩ and ⟨Vmix⟩ are the ensemble-averaged total energy and volume of the system, respectively, xi is the mole fraction of component i, and ⟨Ui⟩ and ⟨Vi⟩ are the ensemble-averaged total energy and volume of component i in its pure state, respectively. In the present calculations, the contribution of the last term in eq 10 is very negligible (because of the incompressible nature of liquids and their mixtures), and hence, eq 10 was reduced to Δhmix = ⟨Umix⟩ − [x IL⟨UIL⟩ + x PCM⟨UPCM⟩]

(11)

where ⟨UIL⟩ and ⟨UPCM⟩ are the ensemble-averaged total energies of the IL and PCM in their pure liquid states, respectively. Block averaging50,51 (the block size was 20 000 simulation steps, and 200 blocks were averaged) was used to estimate the standard deviations (σ). The standard deviation of the total energy was calculated as

in which the functional forms of the individual terms are described by eqs 2−7, except for Etorsion for the [EMIM] cation, for which the CHARMM dihedral functional form (eq 8)49 was used instead of the standard OPLS potential (eq 4):48 bonds

k1[1 + cos(Ø)] + k 2[1 − cos(2Ø)]

torsions

(1)

k b(r − r0)2

∑

Utorsion =

Etotal = E bond + Eangle + Etorsion + E improper + Evdw + Eqq

∑

(3)

angles

2. METHODS 2.1. Molecular Simulation. Simulations were performed with the MD simulation package LAMMPS (version from June 7, 2013) developed by Sandia National Laboratories.37 All of the methods implemented in this study have been carried on from our previous works.38,39 The cutoff radius for nonbonded interactions was 10 Å, and analytical tail corrections were applied. Long-range electrostatics were computed using the particle mesh Ewald summation method.40,41 Cross-interaction parameters between species were computed from Lorentz− Berthelot combining rules42 and the intramolecular 1−4 interactions were scaled to 0.5. All of the simulations were performed under periodic boundary conditions. Simulations were carried out by equilibrating the system with the isothermal−isobaric (NPT) ensemble for 4 ns followed by production runs at 298 K using the canonical (NVT) ensemble for 5 ns, for a total simulation time of 9 ns with a time step of 1 fs. (Simulations for the [EMIM][Ac]/acetic acid system were also conducted with 0.5 fs time steps, and the results, not shown here, were statistically equivalent to the mixture thermodynamic properties such as excess heats and volumes computed with 1 fs time steps.) Simulation configurations were saved every picosecond (1000 time steps). The Nose−Hoover thermostat and barostat were used to maintain constant temperature and pressure, each with a damping factor of 5 fs.43−46 Initially, simulations for pure components were conducted with 512 molecules and compared to those reported in the literature to ensure accuracy. Mixture simulations contained between 50 and 450 polar covalent molecules and 50−450 ionic liquid molecule pairs based on their composition, for a total of 500 molecules in each simulation box. All of the force fields employed in this study were taken directly from the existing literature and have been successfully implemented in our previous studies.38 For the [EMIM] cation, the fully flexible all-atom model developed by Kelkar et al.47 was used. The OPLS all-atom force field model, developed by Jorgensen and co-workers, was used directly for the [Ac] anion as well as for all of the PCMs (including point charges).48 In all of these instances, the total energy of the system is defined by eq 1,

E bond =

kθ(θ − θ0)2

∑

Eangle =

(2) B


Journal of Chemical & Engineering Data σ2 =

⎡ n 1⎢ ∑ σj 2 + n ⎢⎣ j = 1

n

⎤

j=1

⎦

∑ (uj − u)2 ⎥⎥

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Table 1. Excess Volumes of Mixing for Binary Mixtures of (1 −x)[EMIM][Ac] + xPCM at 25 °C (12)

where σj is the block-averaged standard deviation in each of the n independent MD runs, uj is the average energy of sample run j, and u is the overall average. The standard deviations of the excess properties of mixtures were calculated as σ 2 = σmix 2 + σIL 2x IL 2 + σPCM 2x PCM 2

x

VE (cm3·mol−1)

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.50 −0.49 −0.63 −0.55 −0.87 −1.02 −1.30 −1.54 −1.52

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.25 −0.60 −0.74 −0.96 −0.96 −0.87 −0.79 −0.70 −0.40

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.32 −0.47 −0.68 −0.76 −0.82 −1.04 −1.25 −1.33 −1.45

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.67 −0.97 −1.13 −1.36 −1.71 −1.75 −2.09 −2.16 −1.81

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.35 −0.49 −0.71 −0.79 −1.06 −1.04 −1.14 −1.17 −1.08

uncertainty u (cm3·mol−1)

Acetic Acid

(13)

where σmix is the standard deviation of the mixture and σIL and σPCM are the standard deviations of the energies of the IL and PCM in their pure states, respectively.

3. RESULTS Initially, the choice and accuracy of the implemented force fields were validated in our previous works38,39 by determining the room-temperature densities of the pure systems using MD simulations and through ion aggregation analysis and comparison to light scattering data. The simulated results agreed well with the experimental values reported in literature, and therefore, further simulations were carried out to determine the excess properties. For pure [EMIM][Ac], the simulated room-temperature density was 1.0855 g·cm−3, which compares well against the experimental value of 1.027 g·cm−3 and represents an improvement over a previously reported simulation value of 1.1131 g·cm−3 at 300 K.52 3.1. Volumetric Properties and Behavior. The volume of mixing, or excess volume (VE), is a property change when two pure species undergo an isothermal, isobaric mixing process, and it reflects how the interactions of the pure species change in the presence of each other. Determining VE via MD simulations is a very sensitive process since small changes in density result in large changes in volume. Figures 1 to 5 and

0.21 0.21 0.23 0.22 0.22 0.20 0.20 0.21 0.18 Chloroform 0.17 0.16 0.18 0.20 0.18 0.21 0.25 0.23 0.25 DMSO 0.20 0.21 0.21 0.20 0.20 0.21 0.16 0.18 0.19 Methanol 0.23 0.21 0.20 0.20 0.18 0.18 0.20 0.20 0.20 Propanol

Figure 1. Excess volumes for IL/acetic acid mixtures.

Table 1 report VE for binary mixtures of [EMIM][Ac] with five different PCMs: acetic acid (Figure 1), chloroform (Figure 2), DMSO (Figure 3), methanol (Figure 4), and propanol (Figure 5). All of the IL/PCM mixtures show negative VE over the entire composition range (except for the IL/acetone mixture, which has miscibility limitations and is discussed later). This implies dominant net attractive forces between the two species, resulting in closer packing compared with when they are in their native states and hence a negative volume change. The observed values of VE for all of the IL/PCM cosolvents are on the order of 1−2 cm3·mol−1. For most of the cosolvent systems (acetic acid, DMSO, methanol, and propanol), the excess molar

0.17 0.16 0.14 0.17 0.17 0.15 0.16 0.16 0.17

volumes attain minimum values at large mole fractions of PCM cosolvent (x ≥ 0.8). In contrast, the mixture with chloroform attains its minimum excess molar volume at a near 0.5/0.5 IL/ chloroform mole fraction composition. (Additional discussion of the competing intermolecular interactions leading to these minima for the cosolvent systems of acetic acid and methanol can be found in section 3.2 on enthalpic properties.) C



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In general, negative volumes of mixing can be attributed to one or more of several effects such as interstitial accommodation of molecules, formation of weak complexes arising from hydrogen-bonding, π−π, dipole−dipole, and ion−dipole interactions, and solvation effects. It should be noted that all of these observed interactions and microstructures within the simulation configurations are a direct result of the strong electrostatic forces present in these systems, dictated by the fixed point charges on the various atoms of the ions and molecules. To that end, for the force fields considered in this study, hydrogen bonding is not a separate term but is represented within the particular electrostatic interaction between a given pair of intermolecular O and H atomswith fixed point chargesbetween two different ions/molecules. For the purposes of our remaining discussion (in this section and the subsequent one on enthalpic properties), we consider and label these particular intermolecular/ion interactions between O and H atoms as “hydrogen bonds”, as their interaction includes the formal, directional hydrogen bonding as a subset (coarse-grained from the force field models because of the use of fixed point charges) of the overall electrostatic interaction. (Moreover, dipole moments are not calculated directly but are rather a consequence of these same fixed point charges.) For all these investigated systems, the negative VE behavior suggests geometrical/interstitial accommodation of molecules enhanced by differences between the molar volumes of the pure species, thereby reducing the free volume of the mixture. Further negative contributions to VE can be due to selfassociation of the IL at higher concentrations of the PCM, where the IL ions aggregate to form small subdomains, as evidenced by radial distribution functions (RDFs) between the anion and cation. In the case of methanol (Figure 6), one finds

Figure 2. Excess volumes for IL/chloroform mixtures.

Figure 3. Excess volumes for IL/DMSO mixtures.

Figure 4. Excess volumes for IL/methanol mixtures.

Figure 6. RDFs between OA of the anion and C2 of the cation in IL/ methanol mixtures as a function of the mole fraction of methanol.

an increased height of the first solvent shell peak for cation− anion interaction with an increase in the concentration of methanol, indicating IL ion aggregation. While similar trends were observed for all of the other investigated PCMs (data not shown here), the extent of aggregation varied at a PCM mole fraction of 0.9 (Figure 7), with acetic acid being an outlier where limited or no IL ion aggregation was reported. Another factor influencing the sign and the magnitude of VE is the formation of cross-complex interactions between IL ions and PCM molecules, resulting from strong electrostatic (and in

Figure 5. Excess volumes for IL/propanol mixtures.

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Figure 9. Enthalpies of mixing for IL/chloroform mixtures.

Figure 7. RDFs between OA of the anion and C2 of the cation in IL/ PCM mixtures at a PCM mole fraction of 0.9, indicating different levels of ion aggregation.

some cases hydrogen-bonding) anion−cation, cation−PCM, and anion−PCM interactions. These interactions in turn result in more closely packed structures in the presence of each other than in their native pure-component states. More details regarding hydrogen bonding are discussed in later sections of this article. This aggregation behavior of ILs in PCM cosolvents53−56 as well as in aqueous environments32,47,57−60 has been studied and reported extensively. These volume changes are consistent with the experimental results for other ILs with short cation chain lengths (methyl to butyl) and other cosolvents.29−31,33−36,61 It can be noted that for all of the cases studied, the overall changes in volume are very small. This indicates that the IL/PCM mixtures are nearly ideal volumetrically (VE ≈ 0). 3.2. Enthalpic Properties and Behavior. Similar to the volume, addition of PCMs to ILs alters their enthalpic state, which can be determined by the equations given in section 2.2 (via MD simulation results). All of the IL mixtures with PCMs (except acetone) investigated in this study display exothermic mixing, as shown in Figures 8 to 12 and Table 2. The heat of mixing curves for chloroform and DMSO are weakly exothermic (nearly athermal), largely because those molecules do not interact appreciably with the IL ions. This limited degree of interaction is suggested by the respective RDF curves in Figure 7, which indicate large degrees of IL ion aggregation

Figure 10. Enthalpies of mixing for IL/DMSO mixtures.

Figure 11. Enthalpies of mixing for IL/methanol mixtures.

Figure 12. Enthalpies of mixing for IL/propanol mixtures. Figure 8. Enthalpies of mixing for IL/acetic acid mixtures. E



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Table 2. Excess Heats of Mixing for Binary Mixtures of (1 − x)[EMIM][Ac] + xPCM at 25 °C x

HE (kcal/mol)

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.82 −1.66 −2.63 −3.27 −4.03 −4.73 −5.40 −5.31 −3.08

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.04 −0.13 −0.19 −0.31 −0.33 −0.36 −0.31 −0.28 −0.16

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.02 −0.09 −0.10 −0.10 −0.21 −0.25 −0.31 −0.42 −0.48

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.31 −0.65 −1.23 −1.54 −1.97 −2.29 −2.67 −2.31 −1.63

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

−0.15 −0.51 −0.82 −1.20 −1.54 −1.92 −2.10 −1.87 −1.22

atom and one hydrogen-bonding H atom in each molecule, for potential cross-interactions with the charged IL ions. In contrast, the heat of mixing curve for acetic acid reaches a minimum value roughly twice that of methanol/propanol (Δhmix = −6 kcal/mol vs −3 kcal/mol, respectively), because acetic acid has twice the number of hydrogen-bonding oxygen atoms available (carbonyl and carboxylic oxygen atoms) relative to methanol/propanol. When ILs and PCMs mix, there can be net destruction (endothermic) and net creation (exothermic) of interactions simultaneously, and the one that dominates the other defines the nature and magnitude of the HE curve. In the IL/PCM mixtures studied, the latter one is more dominant, and hence, net exothermic behavior is observed. This also indicates attractive interactions between ILs and PCMs that are more stable than those found in their respective pure states. These relative energetically favorable states (compared with their pure states) likely can be a result of short-range ordering of molecules in the solution arising from cross-complex interactions between the species, formation of weak hydrogen-bonding complexes, and solvation effects. (As in the section on volumetric properties above, it should be noted again that all of these observed interactions and microstructures within the simulation configurations are a direct result of the electrostatic forces dictated by the fixed point charges on the various atoms of the ions and molecules.) Since ILs are complex molecules with multiple sites that are capable of participating in hydrogen bonding with themselves as well as with the PCMs, this can play a major role in the nature of the HE curves. It has been reported that in pure [EMIM][Ac] each oxygen atom in an acetate ion (from here on denoted as OA) is capable of forming a hydrogen bond with a hydrogen atom bonded to the carbon between the two nitrogen atoms in the ring of the cation (denoted as atom type H2).25,26 The addition of the PCM alters this network, resulting in the formation of new interactions between the components. The nature and the strength of these interactions as well as their effect on the enthalpic state for each IL/PCM mixture are discussed below in terms of radial distribution functions. The enthalpy of mixing for the IL/acetic acid binary mixture, shown in Figure 8, indicates that the mixing is highly exothermic. There are two distinct regions in the curve: (1) between acetic acid mole fractions of 0.1 and 0.7 there is an increase in the exothermicity of the binary mixture, and (2) beyond an acetic acid mole fraction of 0.7 the exothermic behavior of the mixture decreases. The nature of the HE curve passing through minimum can be due to several factors. The heights of the first peak in the RDFs (Figure 13) generated between the anion and cation (IL aggregation), between OA of the acetate anion and H2 of the cation (IL self-hydrogen bonding), and between OA of the acetate ion and the hydroxyl H of acetic acid (anion−acetic acid hydrogen bonding) indicate that between acetic acid mole fractions of 0.1 and 0.7 there is an increase in the anion−cation, anion−acetic acid, and IL selfhydrogen bonding interactions, leading to the increase in the exothermicity of the binary mixture up to an acetic acid mole fraction of 0.7. Beyond this composition, all of the abovementioned interactions decreased, resulting in a decrease in the exothermicity of the mixture. Hence, the IL−acetic acid HE curve has maximum exothermicity at an acetic acid mole fraction of 0.7. With increases in acetic acid mole fraction from 0.1 to 0.9, the interactions between acetic acid and the cation increased while the self-hydrogen bonding within acetic acid

uncertainty u (kcal/mol)

Acetic Acid 0.11 0.27 0.22 0.18 0.16 0.13 0.11 0.15 0.11 Chloroform 0.11 0.09 0.08 0.08 0.07 0.08 0.07 0.08 0.08 DMSO 0.12 0.11 0.10 0.12 0.13 0.15 0.19 0.22 0.27 Methanol 0.12 0.10 0.10 0.09 0.09 0.10 0.10 0.12 0.14 Propanol 0.11 0.11 0.10 0.11 0.12 0.14 0.16 0.20 0.24

in the presence of chloroform and DMSO (relative to the other PCM cosolvents studied here). This large extent of IL ion aggregation and limited mixing with chloroform and DMSO PCM cosolvents are a result of diminished attractive electrostatic forces based on the smaller point charges on the Cl, H, S, and O atoms of those PCM molecules. The heat of mixing curves for methanol and propanol are essentially identical, largely dominated by the presence of one hydrogen-bonding O F



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water mixtures, too).27 However, as the acetic acid mole fraction is increased to 0.7, the acetic acid−cation, acetic acid− anion, and anion−cation interactions keep on increasing while acetic acid aggregation is less pronounced. Beyond an acetic acid mole fraction of 0.7, all of these interactions are reduced since hydrogen bonding in acetic acid is reestablished and reverts back to its local pure-component structure. However, the effect of the acetic acid self-hydrogen bonding is less compared with the other interactions, leading to the observed result. Similarly, for IL/methanol mixtures, the generated RDFs (and specifically the first peak heights) between (I) the hydroxyl oxygen and hydrogen in methanol (hydrogen bonding) (Figure 15), (II) the anion and cation (Figure 6), Figure 13. RDFs for IL/acetic acid mixtures for the following interactions in the system: (1) between the most negatively charged O atom on the anion [O(Anion)] and the carboxylic H atom of acetic acid [H(AA)]; (2) between O(Anion) and the most positively charged H atom in the imidazolium ring of the cation [H2(Cation)]; (3) between the double-bonded oxygen atom in acetic acid [AA(=O)] and H2(Cation); (4) between the IL anion and IL cation [Anion−Cation, represented by the interaction between O(Anion) and H2(Cation)], and (5) hydrogen bonding within acetic acid [Hbond(AA), represented by the interaction between AA(=O) and H(AA)].

decreased, eventually reaching its pure limit. One other important aspect is that at lower concentrations of acetic acid the self-hydrogen bonding within acetic acid is very high, indicating aggregation of acetic acid molecules (Figure 13). A simulation snapshot is provided in Figure 14 to further support the evidence of acetic acid aggregation at an acetic acid mole fraction of 0.1.

Figure 15. RDFs between the hydroxyl oxygen and hydrogen of methanol depicting hydrogen bonding within methanol in IL/ methanol mixtures as a function of the mole fraction of methanol (methanol mole fractions are shown in the legend).

and (III) the anion and methanol (hydroxyl hydrogen and OA of acetate ion) (Figure 16) indicate that as the concentration of methanol is increased, interactions I and II increase while interaction III increases until a methanol mole fraction of 0.8 and decreases beyond that. At lower concentrations of

Figure 14. Snapshot of the IL/acetic acid mixture, showing only acetic acid molecules and depicting acetic acid aggregation, at an acetic acid mole fraction of 0.1.

On the basis of the above discussion it is evident that at lower concentrations of acetic acid in the IL, acetic acid molecules form aggregates, but as the concentration of acetic acid increases, the presence of acetic acid weakens the anion− cation hydrogen bonding by introducing a hydrogen bond with the cation as well as anion (as seen in the case of [EMIM][Ac]/

Figure 16. RDFs between the hydroxyl hydrogen of methanol and the oxygen of the anion (OA) depicting hydrogen bonding between methanol and acetate anion in IL/methanol mixtures as a function of the mole fraction of methanol (methanol mole fractions are shown in the legend). G



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methanol (xM = 0.1), methanol molecules are found isolated, as shown in Figure 17, in agreement with the RDF in Figure 15,

Figure 18. RDFs between OA of the anion and H2 of the cation in IL/ acetic acid mixtures as a function of the mole fraction of acetic acid (acetic acid mole fractions are shown in the legend).

interactions to the observed net negative volumes and enthalpies, we note the following: 1. There is a precedent in the literature for this type of interaction being prevalent in IL/PCM mixtures.28,30,31,73 2. The negative heats of mixing (from the MD simulations) suggest that cross-interactions (unlike interactions) are more favorable than like interactions. This opens the possibility of cross-complexes (either chemically equilibrated complexes or physical aggregates) between IL ions and PCM molecules. 3. The large RDF first peak values in Figures 13 and 16 suggest extremely high degrees of cross-interaction, well beyond just typical bulk molecular liquid values. In fact, some of the cross (IL ion/PCM) RDF values are 3−5 times larger than typical values, indicating a 3−5 times greater probability of finding a cross species as a nearest neighbor than in the respective pure-component liquid environments. With such a high degree of association, it is unlikely that that is due solely to entropic forces (or packing). It is highly likely that some of those aggregates are energetically stable cross-complexes, which are likely stabilized via solvation in the condensed phase. In the case of IL/acetone mixtures, at higher concentrations of acetone (xacetone ≥ 0.6), the IL and acetone are completely immiscible, as shown in Figure 19. Because of this fact, mixture properties for IL/acetone are not reported. [It should be noted that a mixture with a 10:1 acetone:[EMIM][Ac] molar ratio was prepared in the laboratory according to the experimental methods in our previous study38 and confirmed the lack of miscibility and formation of two phases as predicted by our MD simulations (picture not shown).] HE is an important property for understanding the nature of interactions in the solution thermodynamics of fluid mixtures. HE via experiments for systems with ILs and molecular solvents have been reported extensively.62−73 While both exothermic and endothermic behaviors have been reported, the nature of the counterions and cosolvent present in the system and their hydrogen-bond formation capability define the nature of the HE curve. [EMIM][Ac] is a simpler and less aggressive IL with a hydrophilic anion and therefore should exhibit mild exothermic behavior in the presence of PCMs, as is the case with our previous study with 1-ethyl 3-methylimidazolium methylsulfate, [EMIM][MeSO4], with the same PCMs as studied here39

Figure 17. Snapshot of the IL/methanol mixture, showing only methanol molecules and indicating isolated (nonaggregated) methanol molecules, at a methanol mole fraction of 0.1.

which completely lacks a first solvent shell peak. This is in contrast to acetic acid molecules, which are found to be aggregated (Figure 14). As the concentration of methanol increases, the hydrogen bonding within methanol is reestablished, and at the same time ion aggregation between the anion and cation and anion−methanol interactions increases until xM = 0.8 and then decreases beyond that. These collective interactions/reasons contribute to the formation of more energetically more favorable states, and hence, the enthalpy of mixing for the IL/methanol mixture is negative (exothermic mixing), with a minimum around xM = 0.7 (Figure 11). However, the exothermic behavior for IL/ methanol (Figure 11) is not as pronounced for the IL/acetic acid mixture (Figure 8). Similar results were observed for other the PCMs (chloroform, DMSO, and propanol). While the behavior of propanol was similar to that of methanol, ion aggregation within ILs was more predominant in chloroform and DMSO with respect to the PCM concentration since they are all hydrogen-bond acceptors and are sterically hindered to form a hydrogen bond with the H2 atom type of the cation. In contrast, the acetate anion itself is a hydrogen-bond acceptor with two electronegative OA atoms, and therefore, the exothermic behavior observed for the IL with these two PCMs is less pronounced than in mixtures with alcohols and acetic acid. For IL/acetic acid mixtures at higher mole fractions of acetic acid (xAA = 0.9), significantly less ion aggregation is reported (Figures 7 and 18). Because acetic acid and acetate anion are similar structurally and there is strong crossinteraction (attraction) between O and H atoms on both the molecule and the ion, the acetate anion associates readily with acetic acid, thereby disrupting the acetic acid aggregation extent/network. As stated previously, we define cross-complex interactions as those between an IL ion and a PCM molecule. In regard to the physical justification for the contribution of cross-complex H



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Article

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b01041.

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Tables S1−S7 (PDF)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: (+1) 605 3942636. Fax: (+1) 605 394-1232. ORCID

Kenneth M. Benjamin: 0000-0002-7200-9579 Funding

S.R.P.B. and K.M.B. thank Tyndall AFB (AFRL Award IIP0832549 and FA4819-11-C-0004) and the National Science Foundation Industry/University Collaborative Research Center, the Center for BioEnergy Research and Development, for their financial support of this work.

Figure 19. Snapshot of the IL/acetone mixture at xacetone = 0.9 showing immiscibility between the IL and acetone (only IL molecules are rendered).

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors thank Drs. Michael Cooney and Catherine Rong of the University of Hawaii, Manoa for experimentally confirming the lack of miscibility in [EMIM][Ac]/acetone mixtures with large acetone mole fractions.

(since all of the PCMs participate in hydrogen bonding). As a result, mixtures of these PCMs with the IL result in more efficient packing than in the pure state. Since the species are closer together, dispersion forces come more into effect and result in lower energies. The negative HE values for all of the IL/PCM mixtures likely are accompanied by a net decrease in entropy, leading to more complex, ordered structures and increasing the affinity of the IL toward the PCM. Finally, the current pairwise-additive force fields likely tend to overestimate the excess properties, and it should be noted that the accuracy of predicting excess thermodynamic properties via molecular simulation depends on the accuracy of the force field.74 While a complete description regarding these issues can be found elsewhere,39 these limitations can be overcome by the inclusion of polarizable force fields.74,75

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REFERENCES

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4. CONCLUSIONS This article presents the excess thermodynamic properties of [EMIM][Ac]/PCM binary mixtures obtained via MD simulations. Overall, the mixtures of the IL with PCMs show negative volumes and negative heats of mixing over the entire composition range, except for acetone, which is immiscible with the IL at higher concentrations of acetone. The observed negative VE and HE values are due to one or more of the following reasons: (I) a net increase in attractive intermolecular interactions compared with their native pure-component states; (II) interstitial rearrangement of molecules in the mixtures (not proven here absolutely but strongly suggested in the literature, especially for mixtures with methanol); (III) hydrogen-bond formation and formation of IL−PCM complex interactions; (IV) ion aggregation; and (V) ion solvation effects. The interplay among these interactions, including analyses of IL ion and cosolvent molecule aggregation behaviors, has been demonstrated effectively using radial distribution functions depending on the mixture composition. The resultant exothermic behavior of IL/PCM mixtures is desired in many industrial applications, especially extractions. I



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