What Determines CO2 Solubility in Ionic Liquids ... - ACS Publications

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The Journal of Physical Chemistry

What Determines CO2 Solubility in Ionic Liquids? A Molecular Simulation Study Marco Klähn,1* Abirami Seduraman2 1

Institute of Chemical and Engineering Science, Agency for Science, Technology and Research,

1 Pesek Road, Jurong Island, Singapore 627833 2

Institute of High Performance Computing, Agency for Science, Technology and Research, 1 Fu-

sionopolis Way, #16-16, Connexis, Singapore 138632, Rep. of Singapore

*To whom correspondence should be addressed. Phone: (65) 67963946. E-mail: [email protected]

Abstract Molecular dynamics (MD) simulations of ten different pure and CO2-saturated ionic liquids are performed to identify the factors that determine CO2 solubility. Imidazolium-based cations with varying alkyl chain length and functionalization are paired with anions of different hydrophobicity and size. Simulations are carried out with an empirical force field based on liquid-phase charges. The partial molar volume of CO2 in ionic liquids (ILs) varies from 30 - 40 cm3/mol. This indicates that slight ion displacements are necessary to enable CO2 insertions. However, the absorption of CO2 does not affect the overall distribution of ions in the ILs as demonstrated by almost equal cation – anion radial distribution functions of pure ILs and ILs saturated with CO2. The solubility of CO2 in ILs is not influenced by direct CO2 – ion interactions. Instead, a strong correlation between the ratio of unoccupied space in pure ILs and their ability to absorb CO2 is found. This preformed unoccupied space is regularly dispersed throughout the ILs and needs to be expanded by slight ion displacements to accommodate CO2. The amount of preformed unoccupied space is a good indicator for ion cohesion in ILs. Weak electrostatic cation – anion inter-

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action densities in ILs, i.e. weak ion cohesion, generates more unoccupied space as a result of larger average distances between ions. Weak ion cohesion facilitates ion displacement to enable an expansion of empty space to accommodate CO2. Moreover, it is demonstrated that the strength of ion cohesion is primarily determined by the ion density, which in turn is given by the ion sizes. Ion cohesion is influenced additionally to a smaller extent by local electrostatic interactions among ion moieties between which CO2 is inserted and which do not depend on the ion density. Overall, the factors that determine the solubility of CO2 in ILs are identified consistently across a large variety of constituting ions through MD simulations.

Keywords: Ionic liquid, CO2, carbondioxide, absorptivity, molecular dynamics simulations, MD, force field, partial molar volume, cavities, imidazolium, ion cohesion

1. Introduction Carbondioxide is one of the primary greenhouse gases present in the atmosphere, which is assumed to contribute to an increase of the average temperature of the earth surface. The emission of CO2 gas is rapidly increasing, due to the steadily increasing combustion of fossil fuels, which contributes 86% to greenhouse gases.1 Various methods have been proposed to limit anthropogenic CO2 emissions. Carbon capturing and sequestration has been successfully applied to reduce the CO2 concentration increase in the atmosphere.2 The CCS technology requires chemical and physical absorption and adsorption of CO2. Presently, these processes heavily depend on amine based solvents, typically monoethanolamine (MEA), which is used in CO2 absorption processes.2 However, these amine solvents exhibit various disadvantages, such as causing corrosion problems, as well as being very volatile and not very cost effective. Hence, commercialization of these techniques in large scale power plants is problematic. The identification or design of alternative materials for CO2 capture is therefore of high importance. Ionic liquids (ILs) have been proposed and investigated as an effective alternative for CO2 absorption. ILs are molten salts typically composed of large asymmetric organic cations and smaller inorganic or organic anions. ILs exhibit low melting points below 100 °C and often also below room 2 ACS Paragon Plus Environment

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temperature. IL are characterized, among others, by negligible vapor pressure, a wide liquid range, favorable solvation properties, significant thermal and chemical stability and in this context most importantly the capability of dissolving CO2. Moreover, physicochemical properties can be readily modified over a wide range of values through changing or modifying the anion and cation constituents to optimize CO2 absorption and other relevant properties. Dissolution of CO2 in ILs involves either physical or chemical absorption of the gas. Effective physical absorption of CO2 in ILs is typically accomplished at high pressure and low temperature. These ILs are readily regenerated afterwards by lowering pressure and increasing temperature again to release the captured CO2.3 In chemical absorption, however, CO2 chemically reacts with the IL medium. Suitable ILs for chemical absorption contain amines, acetate and cyanates for instance and have been reported to exhibit a large CO2 solubility that usually exceeds the CO2 solubility found in pure physical absorption processes.4 However, these ILs are less suitable for industrial scale applications because of their high viscosity and especially the high energy requirements to expel captured CO2 to regenerate the IL medium. In the following we will consider only physical CO2 absorption of ILs. Blanchard et. al. were the first to report that CO2 is highly soluble in imidazolium-based ILs, whereas these ILs are insoluble in CO2-rich phases.5 Since this discovery, numerous experimental and computational studies on CO2 solubility in ILs have been reported.4, 6-15 Widely used approaches have been gravimetric microbalance, the synthetic (bubble point) method and isochoric saturation in experimental studies as well as molecular dynamics (MD) and Monte Carlo (MC) simulations in combination with molecular mechanics (MM) or quantum mechanical (QM) molecular models in computational studies. More recently, also thermodynamics models based on COSMO-RS16-20 have been applied to predict CO2 solubility in various ILs.21-23 Imidazolium-based ILs have been considered in the majority of studies because of their high availability and their propensity to absorb CO2. Kazarian et al. used ATR-IR spectroscopy to investigate the mixtures of CO2 and 1-n-butyl-3-methylimidazolium hexafluorophosphate (BMIMPF6) and 1-n-butyl-3-methylimidazolium tetrafluoroborate (BMIM-BF4).24 Their study showed evidence for the existence of favorable interactions between anions and CO2. It was suggested that these interactions could be described as CO2 and anions being Lewis acid and bases, respectively. Cadena et al. confirmed the relevance of anion - CO2 interactions by combining experi3 ACS Paragon Plus Environment


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mental and computational methods.11 It was found that ILs that contained BMIM-cations paired with bis(trifluoromethylsulfonyl)-imide anions (BMIM-NTF2) showed the greatest affinity for CO2 and exhibited larger CO2 solubility compared with BMIM+ cations paired with BF4- or PF6anions. Moreover, in was observed in the same study that a substitution of the acidic proton with methyl in the C2 position of the imidazolium ring lowered the solubility of CO2. Also other studies showed that replacing this acidic proton with ether, hydroxyl, nitrile or alkyne groups impeded CO2 solubility.4 Aki et al. investigated the impact of the length of alkyl chains bonded to cations on CO2 solubility and found that an increasing alkyl chain length led to an increase in CO2 solubility.25 However, this trend was found to be only secondary compared with the impact of the anion identity on CO2 solubility in the same study. That this property is indeed primarily determined by the identity of the anions rather than by modifications of imidazolium-based cations has been confirmed in numerous studies.6-15 Regarding anions, it was found that CO2 solubility increased in the order of BF4- < TfO- < TfA- < PF6- < Tf2N- < C7F15CO2- < eFAP- < bFAP-.4 In another study the following trend of increasing CO2 solubility was found: NO3- < SCN- < MeSO4- < BF4- < DCA- < PF6- < Tf2N- < C7F15CO2-.26 Generally, ILs that contained fluorinated anions tended to exhibit higher CO2 solubilities compared with non-fluorinated anions. On the other hand, fluorination of the cations in the IL did not seem to increase CO2 solubility to a larger extent.27 Carvalho et al. found that the solubility in phosphonium-based ILs can be substantially larger than in imidazolium-based ILs, which suggests that the performance of ILs can be improved further also by finding ideal cations.28

Comparing CO2 solubility in BMIM-BF4 with BMIM-PF6, a larger solubility in the IL containing BF4- should be expected, if the direct CO2 – anion interaction strength was the determining factor, because BF4- is a stronger Lewis base than PF6-. This expectation is contradicted, however, by the observation that CO2 is in fact less soluble in BMIM-BF4 than in BMIM-PF6.11 Thus, Kazarian et. al suggested that CO2 – anion interactions alone are unlikely to play a predominant role in determining CO2 solubility and proposed that CO2 occupies preformed cavities in the ILs and that the availability of these cavities are a determining factor.24 Berne et. al carried out MD simulations of BMIM-PF6 and used a Voronoi algorithm to study cavities inside these liquids.12 It has been found that IL structures barely changed after absorption of CO2, which has also been previously confirmed through comparison of simulated radial distribution functions (RDFs) before 4 ACS Paragon Plus Environment

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and after CO2 absorption.11, 29 The sizes of these cavities were found to be smaller than 1.5 Å, which is too small to accommodate a CO2 molecule. Therefore, it was proposed that larger cavities are formed by reorientations of anions in the IL after CO2 absorption. However, changes in cation-anion pair RDFs, which are used to study structural changes in ILs, were not observed. Jiang et al. reported the significance of free volume in ILs that contained nitrile-containing anions paired with imidazolium-based cations.10 MD simulations were used to study 1-ethyl-3methylimidazolium, EMIM+, paired with B(CN)4- and Tf2N-. It was observed that CO2 solubility was influenced by cation – anion interaction strengths, whereas CO2 – anion interactions seemed to play only a secondary role. Moreover, it was proposed that weak cation – anion interactions caused more cavities in ILs and thus the CO2 solubility increased. Nevertheless, it should be noted that this conclusion was only based on the comparison of two different ILs. Also, it is not clear how interactions among cations, anions and CO2 solutes are connected with the availability of free volume in ILs. More recently, ILs that contain nitrile based anions were investigated experimentally30-31 and computationally.32 It was found that CO2 solubility increased with an increasing nitrile (-CN) contents in anions. The largest CO2 solubility was found in BMIMB(CN)4,30 whereas cation variations caused only mild changes in CO2 solubility.31 This trend was reproduced with MD simulations and attributed to a weakening of cation – anion interactions when anions contained additional nitrile.32 The found important connection of cation – anion interaction strength with CO2 solubility in the latter work, however, was limited to ILs with nitrile-based anions. Also, cation – anion interactions in ILs were only estimated from ab-initio calculations of isolated ion pairs.

The aim of the present work is to scrutinize for a broad variety of ILs the importance of preformed cavities for the absorption of CO2, to investigate the main factors that determine the availability of these cavities as well as the role that ion-ion and ion-CO2 interactions play in determining CO2 solubility. Ten different ILs were considered. Cations were based on imidazolium to enable comparison with measured data. Cations with different alkyl chain lengths (n=4, 6 and 8) and functionalization (methoxy and hydroxyl) were considered as well as five different anions ranging from very hydrophobic (Tf2N-) to very hydrophilic (Cl-). Moreover, fluorinated anions and cations were also considered. These ILs were simulated using MD simulations in combination with an empirical force field at the same temperature and pressure, so as to ensure full com5 ACS Paragon Plus Environment


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parability of the ILs with each other and with experimental data. Because the formation of unoccupied space in simulations is very sensitive to the overall IL structure, which in turn is determined by a delicate balance of inter-ionic electrostatic and van der Waals interactions, we applied our previously developed force field.33-35 This force field is based on liquid-phase charges rather than gas-phase charges and was derived with a QM/MM approach. Pure ILs were simulated and the unoccupied space in these ILs was found and quantified. Moreover, the same ILs were simulated together with CO2 at saturation concentrations to assess structural changes induced by CO2 absorption and to determine the partial molar volume of CO2. Furthermore, anion – cation and ion – CO2 interaction strengths were determined. After identifying the factors that determine CO2 solubility in ILs, we compared these factors with those that determine the solubility of H2O in ILs, as described in our preceding work,34 striving to improve a general understanding of the solubility of small gas molecules in ILs.

2. Methodology 2.1 Simulated Ionic Liquids A wide range of ILs was considered to study the influence of anions, changes in cation alkyl chain length and functionalization as well as ion fluorination on CO2 absorption. To study the anion effect, BMIM cations were paired with five different anions, which yielded: 1-butyl-3methylimidazolium hexafluorophosphate (BMIM-PF6), 1-butyl-3-methylimidazolium tetrafluoroborate (BMIM-BF4), 1-butyl-3-methylimidazolium nitrate (BMIM-NO3), 1-butyl-3methylimidazolium

chloride

(BMIM-Cl),

and

1-butyl-3-methylimidazolium

bis(trifluoromethylsulfonyl)imide (BMIM-NTF2). To study the alkyl chain length effect, 1hexyl-3-methylimidazolium bis[(trifluoromethyl)sulfonyl]amide (C6MIM-NTF2) and 1-octyl-3methylimidazolium hexaﬂuorophosphate (C8MIM-PF6) were added and compared with BMIMNTF2 and BMIM-PF6, respectively. To investigate cation fluorination, 1-methyl-3(3,3,4,4,5,5,6,6,6-nonafluorohexyl)imidazolium (C6F9MIM-NTF2) was simulated and compared with C6MIM-NTF2. Also the functionalization of cations was investigated by simulating 1-(2hydroxyethyl)-3-methyl-imidazolium

(C2OHMIM-BF4)

and

1-methoxyethyl-36

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methylimidazolium tetrafluoroborate (MOEMIM-BF4), which could be compared with BMIMBF4. The structural formulas of all considered ions are shown in Figure S1. All ILs were simulated without CO2 and saturated with CO2, using measured saturation concentrations for the temperature and pressure used in simulations.25, 27, 36-43 All MD simulations were carried out at the same temperature and pressure to ensure that simulation results could be compared across systems. We chose T = 298 K and P = 10 bar, for which measured CO2 saturation concentrations were available for all considered ILs, except for BMIM-Cl. Measured saturation concentrations of CO2 and the resulting total number of CO2 molecules in simulation boxes are listed in Table 1. In case of MOEMIM-BF4 and BMIM-Cl simulations were only performed without CO2. In the former case of MOEMIM-BF4, the large measured CO2 solubility is known to be caused by chemical absorption, which is not within the scope of this work. Nevertheless, we still considered this IL to study the effect of the cation functionalization on the IL structure and available empty space. BMIM-Cl exhibits a very low CO2 solubility and therefore solubility has only been measured at high pressure. Yet, we preferred to include this IL as a representative of a class of ILs that contains small monoatomic anions to investigate their structural characteristics. The structures of the simulated ions are shown together with their abbreviations in Figure 1.

2.2 Applied Force Field The potential energies of simulated systems were calculated with a force field using a standard functional form as described for instance in ref 33. In all IL simulations we use our own force field parameterization, in which bonded interactions as well as Lennard-Jones parameters were taken from Lopes et al.44-46 Atomistic partial charges for Coulomb potentials, however, were derived as liquid-phase charges by employing a QM/MM approach, in which the classical MD software GROMACS 4.5.547 was coupled with the quantum chemistry package Gaussian 09.48 It is expected that in ILs the charge polarization of ions and charge transfer between anions and cations strongly influence the overall electrostatics within the IL and therefore its structural organization and the formation of empty space between ions. This has already been demonstrated in our previous work.35 For this reason, the average charge transfer and charge polarization in each simulated IL were taken into account.



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For the evaluation of electron charge transfer, a cluster consisting of six or seven coordinated ion pairs were chosen from a liquid structure previously generated with MD as the QM fragment, whereas the remaining ion pairs constitute the MM fragment. The QM fragment was described with DFT using the B3LYP hybrid functional and the 6-31+G(d) basis set.49-50 The MM fragment was described with the force field. The steepest gradient method was used to minimize the potential energy of the QM/MM system until the charges, derived from the electrostatic potential (ESP) using the CHELPG scheme,51 on the atoms converged. This procedure was repeated three times with different QM fragments. The charge transfer was then determined from the average total charge of the ions in the QM fragments. To determine the average charge polarization of ions in the IL, a single ion was chosen for the QM fragment, and all remaining ions were included in the MM fragment. After energy minimization of these systems, the ESP charges of the QM atoms were determined. This process was repeated twelve times for different anions and cations, respectively, after which derived ESP charges were averaged. In the last step, these charges were scaled with the previously determined average total ion charge to ensure that all partial charges were adding up to yield the correct electron charge transfer. More details on this method and the resulting parameters have already been presented for the ILs BMIM-PF6, BMIM-BF4, BMIM-NO3, C8MIM-PF6, and MOEMIM-BF4 in our previous work.34, 52

We used the same approach to derive liquid phase charges for the remaining ILs BMIM-CL,

BMIM-NTF2, C6MIM-NTF2, C6F9MIM-NTF2 and C2OHMIM-BF4, which are displayed in Figure S2. The average total ion charges in the liquid phase of all simulated ILs are summarized in Table 2. For the CO2 solutes the TraPPE force field was used with virtual sites.53

2.3 Specification of MD simulations MD simulations for pure ILs and IL - CO2 mixtures were performed with the GROMACS 4.5.5 software package.47 All simulations contained 500 ion pairs. In simulations that contained CO2, the number of CO2 molecules was determined from measured saturation concentrations, as specified in Table 1. Periodic boundary conditions were applied to the cubic simulation boxes. No de8 ACS Paragon Plus Environment

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grees of freedom were constrained, except for C-O bonds in CO2 due to the use of virtual sites. All systems were simulated in the isothermal-isobaric ensemble (NPT) at T = 298.15 K and P = 10 bar. A velocity rescaling thermostat was used with a coupling constant of 1 ps.54-55 ParrinelloRahman coupling was applied to control the pressure, using a coupling constant of 0.1ps and a compressibility of 4.5 × 10-5 bar-1. Non-bonded interactions were calculated using a cut-off of 15 Å. Long range electrostatics were calculated using fast particle-mesh Ewald (PME) electrostatics together with long-range dispersion corrections.56-58 An integration step size of 1fs was used and translation of the center-of-mass was removed during simulations. For analysis, energy data was written out every 0.2 ps, while coordinates, velocities and forces were written out every 2 ps. Simulations were initialized with arbitrary ion orientations in the simulation box. Initial close contacts between ions were removed by potential energy minimizing with the steepest descent method. The strong electrostatic attraction between counter-ions diminishes ion mobility in simulations and is likely to cause simulations to be trapped in unfavorable high energy states close to their initialization. To mitigate this problem all simulations were started with low density and high temperature to facilitate ion mobility. The temperature of the systems was linearly increased from 0 to 600 K within the first 1 ns, after which a high temperature of 600 K was maintained for 10 ns. The temperature was slowly lowered to the target temperature of 298.15K within the next 1 ns, after which the temperature remained unchanged at 298.15 K during the next 20 ns. The last 10 ns of the simulations were used for analysis. In some cases, MD runs were extended for another 10 ns to ensure complete equilibration of the liquids. Calculated mass densities of all simulated ILs were compared with measured mass densities to validate our force field. Measured mass densities were only available for BMIM-PF6, BMIMBF4, and BMIM-NTF2 for the temperature and pressure that was used in simulations. The densities of the other ILs were compared instead at conditions for which experimental data was available. Only in the case of C6F9MIM-NTf2 it was necessary to compare the CO2 saturated IL density because the density of the pure IL was not available. The comparison is summarized in Table S1. Overall, the rmsd of calculated mass densities from experimental data was only 2.6%, thereby demonstrating that the force field accurately reproduces IL densities.



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2.4 Analysis of Unoccupied Space in ILs To analyze unoccupied space in simulated ILs we used our previously developed XCav tool as described in ref. 59, where it has been used successfully to analyze the partial molar volume and the structure of unoccupied space in various solutions of organic solvents and small gas solutes, including CO2. XCav was developed specifically to localize and quantify unoccupied space in MD trajectories. Moreover, XCav can be used to determine the shape of cavities and their size distribution. In the present work, however, this tool was only used to determine the percentage of unoccupied space in all simulated ILs. XCav is based on a simple algorithm that has already been proposed in different variants before.60-62 For a given trajectory snapshot, the entire simulation box is divided into cubic cells of user-defined size in the first step, which determines the resolution of the method. We used a cell length of 0.05 nm, which gave these cells a volume of only 0.017·VH, where for comparison VH is the volume of a van-der Waals sphere of a hydrogen atom. In the next step, a sphere with atom type specific radius is defined around each atom in the simulation box. To ensure consistency with the force field used in simulations, these radii were set to 0.56·i, where  is the Lennard-Jones -parameter of atom i (using a Lennard-Jones potential of the form 4·[(/r)12 (/r)6]). The distance of 0.56·i is exactly half of the distance at which the Lennard-Jones potential between two atoms of the same type exhibits its minimum, which therefore provides a reasonable and consistent approximation for the size of atoms. In the next step, all cells were removed that overlapped with at least one of the atom-centered spheres. As a result, all the remaining cells describe the unoccupied space that is contained in the respective trajectory snapshot. The percentage of unoccupied space is then simply calculated by multiplying the number of remaining cells with their cell volume and dividing the result by the volume of the simulation box. This procedure was repeated for 1000 different structures that were taken from MD trajectories with a time interval of 10 ps between two consecutive snapshots, for each simulated IL, respectively. The amount of unoccupied space was then averaged over all these snapshots.


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3. Results and Discussion 3.1 Structural Changes in ILs Induced by CO2 Absorption To study the influence of CO2 absorption on the overall IL structure, the center-of-mass RDFs of cation - anion pairs was determined in ILs before and after CO2 absorption. The results are shown in Figure 2. Consistent with previous simulation studies,12, 29 significant structural changes induced by CO2 were not observed. This result supports previous findings that CO2 absorption does not perturb the basic underlying IL structure significantly. Only in C8MIM-PF6 and C6F9MIM-NTF2 minuscule changes can be noticed. Generally, the intensity of the first maximum of these RDFs is related to the cation – anion interaction strength and corresponds to a coordination of the anions to the positively charged ring of the cations. The comparably small first maximum in C2OHMIM-BF4, despite of strong cation – anion interaction strengths, is a result of the high electronegativity of the cation oxygen, which shifts electron charge from the neighboring CH2-group toward the oxygen. This charge shift leads to a positively charged CH2 group as shown in Figure S2c. As a result the anion did not only coordinate to the cation ring, as indicated by the usual maximum of the RDF at about 5 Å, but also to the CH2 group of the alkyl chain, which caused a second peak of the RDF at about 6 Å. The partial molar volume of solutes, 𝑉̅𝑖 , is a thermodynamic quantity that is defined as the volume change of a solution after addition of one mole of the solute. Its value is a function of temperature, pressure and solute concentration. Values of 𝑉̅𝑖 for CO2 were determined in eight different CO2-saturated ILs and the values are listed together with the average simulation volumes of pure ILs and the IL-CO2 solutions in Table 3. Overall, 𝑉̅𝑖 values of CO2 in ILs at saturation concentration were found between 30-40 cm3/mol, which was substantially smaller than in conventional organic solvents, where 𝑉̅𝑖 ranged from 45-53 cm3/mol in the limit of infinite dilution.59 In other words, CO2 apparently displaces solvent molecules much more in organic solvents than in ILs. This suggests in case of ILs a substantially weaker perturbation of the overall solvent structure caused by CO2 adsorption, which is consistent with the observation of unchanged RDFs above. However, because 𝑉̅𝑖 > 0 it also means that at least a miniscule perturbation is present, even though it may be too small to be detected by a comparison of RDFs. Fur11 ACS Paragon Plus Environment


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thermore, it can be observed that the same 𝑉̅𝑖 value of 36-37 cm3/mol was found for all ILs that contained BMIM cations, which means that anions did not seem to influence the value significantly, despite of the large differences in anion size and polarity used in simulations. In the two ILs that contained C2OHMIM and C6F9MIM cations somewhat lower 𝑉̅𝑖 -values of 32 and 30 cm3/mol were found, whereas in the C8MIM containing IL a larger value of 𝑉̅𝑖 = 40 cm3/mol was observed. The calculated 𝑉̅𝑖 of CO2 in BMIM-PF6 (36 cm3/mol at T = 298 K, P = 10 bar) is comparable to a measured value of 29 cm3/mol (T = 318 K, P = 200 bar), which remained nearly constant up to a CO2-mole fraction of almost 0.5.11-12

3.2 Interactions of CO2 with Ions The average interaction strength of CO2 with the ions in ILs was derived to examine its influence on CO2 saturation concentrations. These average interaction strengths per CO2 molecule with cations and anions, respectively, and with contributions from Coulomb and Lennard-Jones potentials are listed in Table 4. The largest interaction strengths of CO2 with ions were observed in BMIM-NO3 and C2OHMIM-BF4, which involve the most hydrophilic anion and cation, respectively. On the other hand, the weakest interactions were found in C8MIM-PF6, which contains hydrophobic ions. It appears that these interaction strengths were determined by the hydrophilicity, i.e. the polarity, of the ions in the IL. This is not surprising, because the two CO2 bonds are polarized, even though CO2 does not possess an overall dipole moment due to its linearity. Because of this lack of dipole moment, electrostatic interactions of CO2 with ions were smaller than their Lennard-Jones interactions. For this reason, CO2 – IL interaction strengths vary only mildly across ILs, because Lennard-Jones interaction strengths depend mostly only on the solute-solvent contact area size, which should barely change for the CO2 solutes in different ILs. In solutions where solute – solvent Coulomb interactions dominate, it would lead to reorientations and displacements of solvent molecules so as to maximize those interactions. This was not the case in CO2 – IL solutions, which already indicated that direct CO2 – ion interactions could play only a minor role in determining CO2 saturation concentrations.


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To explore the role of CO2 – ion interactions further, these interaction strengths were compared with measured saturation concentrations, which is shown in Figure 3. It can be clearly seen that these two properties do not correlate, which means that direct CO2 – ion interactions do not play a significant role in determining the CO2 saturation concentrations in ILs. For instance, CO2 exhibits the second strongest interactions with ions in C2OHMIM-BF4, however, CO2 saturation in this IL is the lowest in all considered ILs. This result is also consistent with previously reported findings from computational studies.10, 24, 63

3.3 Analysis of Preformed Unoccupied Space in ILs To quantify the volume of preformed unoccupied space in pure ILs, XCav was used to determine the ratio of empty space present in each IL as described in section 2.4. The results are shown in Table 5. The ratio of empty space present in each pure IL, runocc, decreases in the order of C6MIM-NTF2 > C8MIM-PF6 > BMIM-NTF2 > BMIM-PF6 > C6F9MIM-NTF2 > BMIM-BF4 > BMIM-CL > MOEMIM-BF4 > C2OHMIM-BF4 > BMIM-NO3. Values of runocc varied from 0.175-0.208, which corresponds to an overall variation of less than 20%. To explore the importance of preformed unoccupied space in ILs for CO2 absorption, runocc was compared with measured CO2 saturation concentrations in Figure 4. A strong correlation with a large R2 coefficient of 0.91 was observed, despite of the comparably small variation of runocc across ILs. ILs with larger ratios of preformed empty space are capable of absorbing larger amounts of CO2. We like to emphasize that the statistical error of runocc, calculated with a block averaging method, was in all cases below only 0.001. This correlation strongly suggests that strong CO2 absorption in ILs is accompanied with the presence of larger areas of empty space in these ILs even before CO2 is absorbed, which indicates that direct interactions of the CO2 solutes with surrounding ions cannot influence CO2 absorption significantly. Now the question arises whether CO2 molecules really reside in these preformed cavities found in ILs. The volume of a CO2 molecule can be estimated with the McGowan method (this simple method is outlined in section 3.5) to be roughly 27.7 cm3/mol.64 An estimate using the van der Waals radii of oxygen and carbon results in a somewhat smaller value of 21.2 cm3/mol. C6MIM13 ACS Paragon Plus Environment


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NTf2 for instance exhibited the largest value of runocc, which translates to an area of unoccupied space of 198 cm3/mol that is available for each CO2 molecule at saturation concentration. This means that ample of empty space would be available to accommodate all CO2.molecules. For this IL we searched with XCav for single cavities that are larger than at least half of the volume of CO2 (27.7 cm3/mol) but on the average only one such cavity in this IL was found. This means that preformed cavities in ILs are too small to accommodate CO2. This was also already indicated by calculated values of 𝑉̅𝑖 between 30 – 40 cm3/mol for CO2, which is not far from the expected volume of CO2 that we estimated above. In other words, the volume of the IL increased by the expected volume of CO2 after its absorption instead of adopting a value close to zero as it should have been if preformed empty space had been utilized. From this observation follows that ions in the IL had to be displaced to generate solvent cages for CO2. Nevertheless, the required ion displacements apparently occurred in a manner that minimized a possible perturbation of the overall ion organization in the IL as indicated by virtually unchanged RDFs of cation – anion pairs. It appears that for detecting such miniscule IL structural changes, such RDFs are not sufficiently sensitive. In any case, the preformed empty space in ILs could accommodate CO2 only partially because it was too dispersed throughout the liquid. Direct CO2 – ion interactions were also compared with runocc in Figure 5. A correlation of these two properties suggests that ILs with smaller amounts of empty space led to stronger CO2 interactions with ions. An IL with a lower value of runocc,, i.e. with a more compact ion organization, would lead inevitably to reduced average distances between CO2 solutes and adjacent ions and thereby stronger CO2 – ion Coulomb and Lennard-Jones interactions. Also, more compact ILs were comprised of more polar ions, which enable stronger interactions with CO2. This latter point will be analyzed further in the following section, where the factors that determine the compactness of an IL are discussed. In any event, as demonstrated in Figure 3, direct CO2 – ion interactions did not determine CO2 saturation concentrations, rather the factors that lead to compact ILs and thereby low CO2 saturation concentrations also caused indirectly stronger CO2 – ion interaction strengths. The distribution of empty space, according to the results from XCav, is shown in Figure 6 for pure BMIM-PF6 and when saturated with CO2. The empty space distributions are also shown in Figures S3 and S4 for C2OHMIM-BF4, which exhibits low CO2 solubility, and for C6MIM-NTf2, 14 ACS Paragon Plus Environment

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in which a high CO2 solubility was found. The empty space between the ions is regularly distributed throughout the liquids. It can also be seen that the empty space is only sufficient to accommodate CO2 partially and that small ion displacements are necessary to enable insertion of entire CO2 molecules.

3.4 Ion Interaction Density In this section, the average cation – anion attraction was determined with respective Coulomb and Lennard-Jones contributions. Moreover, a cation – anion interaction density, cat-an, was derived from eq. 1:

𝐂𝐚𝐭−𝐀𝐧 𝐂𝐚𝐭−𝐀𝐧 cat-an = (MD + MD ) / MD

(1)

𝐂𝐚𝐭−𝐀𝐧 𝐂𝐚𝐭−𝐀𝐧 In eq. 1 𝐸Coul and 𝐸LJ are the total Coulomb and Lennard-Jones attraction between cati-

ons and anions in simulations averaged over the entire MD trajectory, whereas V designates the average simulation volume. These energies and energy densities are listed in Table 6 for all ILs. Naturally, electrostatic interactions dominate over Lennard-Jones interactions in such systems. Values of cat-an vary widely from -450 kJ mol-1 nm-3 in C6F9MIM-Tf2N to -1247 kJ mol-1 nm-3 in C2OHMIM-BF4. To examine whether a connection between these interaction densities and the amount of unoccupied space exists, values of cat-an and runocc were compared in Figure 7. Remarkable is a reasonably good correlation between these two quantities (R 2=0.68), where larger amounts of preformed cavities in ILs were observed in ILs with weaker cation – anion interaction densities. Hence, ILs that exhibit a strong cation – anion attraction per volume element tend to be more compact, which allows less empty space between the ions. Despite this general trend we also observed two outliers: BMIM-Cl, which contains the smallest ion pairs of all considered ILs, and C6F9MIM-NTf2, which is comprised of the largest ion pairs. It appears that in BMIM-Cl somewhat more regions of empty space formed than the strong cati15 ACS Paragon Plus Environment


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on – anion interaction density would suggest, whereas in C6F9MIM-NTf2 somewhat less empty space formed despite of very weak cation – anion interactions. A linear regression of cat-an and runocc in which these two ILs are excluded would yield a large R2-value of 0.92 as shown in Figure 7. This finding hints at a secondary effect that influenced the value of runocc in simulations, albeit to a smaller extent. In the next section we will first examine the primary effect on runocc, i.e. cat-an, further. Thereafter, we will discuss possible reasons behind the secondary effects that influenced empty space formation in BMIM-Cl and C6F9MIM-NTf2 in section 3.6.

3.5 Ion Density and Size In this section we strive to identify the factors that determine the cation – anion interaction density, cat-an, which in turn has been identified in the previous section as the predominant factor that controls the amount of unoccupied space in pure ILs. It is straightforward to presume a relation between cat-an and the ion density in simulations, which is confirmed in Figure 8. Indeed, a very strong correlation with an R2-value of 0.95 was observed. Large ion densities led to strong cation – anion interaction densities. This confirms that the ion density determines cat-an. For instance, the largest ion density in BMIM-Cl led to the second largest value of cat-an, whereas in the IL with the smallest ion density, C6F9MIM-NTf2, the lowest values of cat-an was found. In the latter case for instance, the IL is composed of large ions and the resulting low ion density led to larger average distances between counter-ions, which reduced their attractive electrostatic interactions and thereby cat-an. These weaker attractive ion interactions allowed a less compact ion organization in the liquid and thus more preformed empty space between ions and the ability to absorb more CO2. Because ion density plays apparently such an important role in determining CO2-absorption of ILs, a simple method to estimate ion densities a-priori would be beneficial. The size of an ion pair can be approximated with the simple McGowan method, in which for each atom in the ion pair an element-specific volume is added and from which a constant volume is subtracted for each present covalent bond.64 The McGowan volumes of all ion pairs considered in simulations are listed in Table 7. These McGowan volumes were converted into McGowan ion densities, 16 ACS Paragon Plus Environment

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which are yielded with the approximation that the entire IL volume is filled with ion pairs of their respective McGowan volumes. McGowan ion densities and actual average ion densities found in simulations are compared in Figure 9. A very strong correlation with an R2-value of 0.98 shows that ion densities are indeed predictable using only the McGowan volumes of the constituting ions. It also confirms that ion densities in ILs simply depend on the volume of the constituting ions.

3.6 Comparison of CO2 absorption in ILs with the absorption of H2O As discussed in the previous sections, to insert a CO2 solute into an IL, ions in the vicinity of the solute need to be slightly separated from each other to generate the solvent cage between these two ions. The ease with which these ions can be separated from each other is determined by the ion density, which in turn determines the strength of electrostatic cation – anion attraction that resists the insertion of CO2. Therefore, ILs with low ion densities and hence weak cation – anion interactions can absorb more CO2. Larger quantities of empty space are formed in such ILs for the same reason because the reduced ion-attraction allows larger average distances between the constituting ions. These unoccupied areas, however, accommodate CO2 only partially because of their dispersal throughout the IL without forming cavities that are sufficiently large to contain single CO2 molecules. These relations are shown schematically in the top row of Figure 10. In section 3.4 it was found that there appeared to be also a secondary effect that determined the amount of unoccupied space in ILs and hence influenced the mutual cohesion of ions in the IL. Ion cohesion should not depend on the global IL property ion density alone but also to a lesser extent on local interactions at the place of the CO2 insertion. To insert a CO2 molecule between two ions they need to be separated, which is generally facilitated where local ion – ion interactions are only weakly attractive or even repulsive. In no other simulated IL is the size difference between cations and anions larger than in BMIM-Cl. An energetically unfavorable direct contact between cations is least avoidable in this IL. The mutual repulsion of cations in these direct contact regions should therefore ease cation displacement and lead to a somewhat higher CO2 absorptivity than the high BMIM-Cl ion density indicates. This could explain the actually observed



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somewhat elevated amount of empty space in this IL. This effect is illustrated in Figure 10 (left hand side, bottom). Similarly, ILs that exhibit a minimum of such repulsive moieties in direct contact, between which CO2 could be inserted, should show a somewhat lower CO2 absorptivity than their ion density suggests. In all considered ILs cations were substantially larger than anions, which is typical for the majority of ILs in general. Hence, direct contact between cations existed in all these ILs to varying extents. In C6F9MIM-NTf2, however, which contains functionalized cations, the fluorinated part of the alkyl chain contains fluorine atoms with negative partial charges on the cation surface as shown in Figure S2c. These negatively charged fluorine atoms strongly interact with the more or less positively charged H-atoms of other cation parts, thereby strengthening ion cohesion. This effect could have been additionally emphasized by the propensity of cations with long alkyl chains to aggregate, as it has been previously observed on numerous occasions (see e.g. ref. 65), so that fluorinated and non-fluorinated parts of the cation alkyl chains could favorably interact. This could explain the observed somewhat lower absorptivity and amount of unoccupied space in this IL, which is shown schematically in Figure 10 (right hand side, bottom). It is instructive to compare the absorption of CO2 with H2O in ILs. Because of the non-linearity of H2O that leads to a strong dipole moment of water, the situation is fundamentally different. In our preceding work we identified the main factors that determine the solubility of water in ILs.34 Water competes with the cations in the IL to coordinate with anions to enable strong favorable electrostatic water – anion interactions. In fact, the water saturation concentration in ILs can well be estimated a-priori from the average interaction strengths between water and anions over a wide range of different ILs, where stronger interactions result in larger water saturation concentrations. Primarily, these interactions are determined by the anion size, where small compact hydrophilic anions with a more localized negative charge are capable of stronger interactions with water. Secondarily, stronger internal anion charge polarization, which leads to larger negative anion charges on the part of the anion surface that is in direct contact with water, strengthens interactions with water further. In contrast, without dipole moment, linear CO2 interacts with the ions too weakly to break the cation-anion coordination. For CO2 absorption, both anions and cations played an important role in determining the solute saturation concentration, whereas in the case of water absorption satu18 ACS Paragon Plus Environment

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ration is mainly determined by anions. Furthermore, for CO2 absorption cation – anion interaction strengths are decisive, whereas for H2O absorption water – anion interactions are crucial. Moreover, large hydrophobic anions weaken cation – anion interactions, which enabled a higher absorptivity of CO2, whereas the same large anions also weaken water – anion interactions and thus impede the absorptivity of H2O. It seems likely that these two absorption mechanisms are also realized for other small solutes, where, depending on the polarity of the solute, absorption undergoes a mechanism more similar to either CO2 or H2O. It remains to be studied weather a transition from a CO2- to an H2O-like absorption mechanism is more continuous or rather more abrupt with increasing solute polarity.

4. Conclusions MD simulations of ten different pure ILs and saturated with CO2 were performed, in which imidazolium-based cations with varying alkyl chain length and functionalization were paired with anions of varying polarity and size. The structures of ILs before and after CO2 absorption were compared with RDFs of cation – anion pairs. The ratio of unoccupied space and cation – anion interaction densities were calculated as well as the PMV of CO2. The CO2 saturation concentration in ILs was not determined by direct CO2 – ion interactions. It was found instead that high CO2 absorptivity was accompanied with a large amount of preformed unoccupied space in ILs. This indicated that a weak ion – ion cohesion, which facilitates ion displacement, enabled the insertion of CO2 between these ions. In ILs with weak interionic interactions, larger areas of unoccupied space formed because the reduced ion-attraction allowed larger distances between ions. Because these areas of unoccupied space were found to be too dispersed throughout the ILs, local expansions of empty space through small ion displacements were required to accommodate CO2 molecules. This finding was in line with calculated PMVs of CO2, which were significantly smaller than in organic solvents but large enough to clearly indicate that ions had to be displaced to form solvent cages. The formation of solvent cages, howev-



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er, perturbed the overall ion organization in the IL only minimally as demonstrated by the lack of change in cation – anion RDFs after CO2 absorption. The cohesion of ions in the ILs was primarily determined by the ion density, which in turn was a consequence of the size of the constituting ions. Low ion densities led to larger average distances between cations and anions and thus weaker mutual attraction. This in turn led to a less compact IL and to a facilitation of CO2 insertion. Methods to estimate cation and anion size, such as for instance the McGowan method, can be readily used to predict ion densities a-priori. To a lesser extent, ion cohesion is also influenced by local interactions between neighboring moieties between which CO2 is inserted. ILs with large cation – cation contact areas that contain very small anions, as it was the case in BMIM-Cl, cation – cation repulsion somewhat facilitated CO2 absorption despite of a large ion density. On the other hand, a cation functionalization that introduced negative surface charge on parts of the cation, as in C6F9MIM-NTf2, reduced such repulsive cation – cation contact areas, thereby somewhat impeding CO2 absorption despite of a low ion density. CO2 absorption was found to be fundamentally different from H2O absorption. While H2O absorption was facilitated by strong H2O – anion interactions, CO2 absorption was instead facilitated by weak cation – anion interactions. From these findings we conclude that to maximize the physical CO2 absorptivity of an IL, the ions should be large to lower the ion density, cations should be considerably larger than anions to provide regions that contain repulsive cations in mutual vicinity, and avoidance of a functionalization that leads to attraction between different cation moieties.

Acknowledgement We acknowledge gratefully the provision of computing facilities by the A*STAR Computational Resource Centre (ACRC) and the financial support from the Agency for Science, Technology and Research (A*STAR) of Singapore.


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Supporting Information Comparison of measured and calculated Mass Densities (Table S1); Structural formulas of all simulated ions (Figure S1); Partial charges used in force field for simulated ions (Figure S2); Visualizations of unoccupied space in C2OHMIM-BF4 and C6MIM-NTf2 (Figures S3-S4). This material is available free of charge via the Internet at http://pubs.acs.org.

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Potoff, J. J.; Siepmann, J. I., Vapor–liquid equilibria of mixtures containing alkanes, carbon dioxide, and nitrogen. AIChE J. 2001, 47 (7), 1676-1682. Berendsen, H. J. C.; Postma, J. P. M.; Gunsteren, W. F. v.; DiNola, A.; Haak, J. R., Molecular dynamics with coupling to an external bath. J. Chem. Phys. 1984, 81, 3684-3690. Bussi, G.; Donadio, D.; Parrinello, M., Canonical sampling through velocity-rescaling. J. Chem. Phys. 2007, 126, 014101. Allen, M. P.; Tildesley, D. J., Computer simulations of liquids. Oxford Science Publications: Oxford, 1987. Darden, T.; York, D.; Pedersen, L., Particle Mesh Ewald - An N.log(N) method for Ewald sums in large systems. J. Chem. Phys. 1993, 98 (12), 10089-10092. Essmann, U.; Perera, L.; Berkowitz, M. L.; Darden, T.; Lee, H.; Pedersen, L. G., A smooth particle mesh Ewald method. J. Chem. Phys. 1995, 103 (19), 8577-8593. Klähn, M.; Martin, A.; Cheong, D. W.; Garland, M. V., Variation and decomposition of the partial molar volume of small gas molecules in different organic solvents derived from molecular dynamics simulations. J. Chem. Phys. 2013, 139 (24), 244506. Rintoul, M. D.; Torquato, S., Algorithm to compute void statistics for random arrays of disks. Phys. Rev. E 1995, 52 (3), 2635-2643. Sastry, S.; Corti, D. S.; Debenedetti, P. G.; Stillinger, F. H., Statistical geometry of particle packings. I. Algorithm for exact determination of connectivity, volume, and surface areas of void space in monodisperse and polydisperse sphere packings. Phys. Rev. E 1997, 56 (5), 5524-5532. Stahl, M.; Bur, D.; Schneider, G., Mapping of proteinase active sites by projection of surfacederived correlation vectors. J. Comput. Chem. 1999, 20 (3), 336-347. Prasad, B. R.; Senapati, S., Explaining the differential solubility of flue gas components in ionic liquids from first-principle calculations. J. Phys. Chem. B 2009, 113 (14), 4739-4743. McGowan, J. C., Estimates of the Properties of Liquids. J. Appl. Chem. Biotech. 1978, 28 (9), 599607. Seduraman, A.; Klähn, M.; Wu, P., Characterization of nano-domains in ionic liquids with molecular simulations. Calphad 2009, 33 (3), 605-613.


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Table 1: Measured CO2 Saturation Concentrations and Number of CO2 Molecules Used in Simulations ILs BMIM-PF6 36-37 BMIM-BF4 36 BMIM-NO3 25 BMIM-NTF2 36, 38 C6MIM-NTF2 42-43 C8MIM-PF6 40 C2OHMIM-BF4 41 C6F9MIM-NTF2 39

[CO2]a 0.166 0.146 0.099 0.231 0.255 0.218 0.093 0.260

# CO2b 100 86 55 150 171 139 51 176

a

Measured CO2 saturation concentration at T = 298 K and P = 10 bar in mole fraction. In some cases where values for the exact temperature and pressure were not available, interpolation of available measured values in the close vicinity of the desired temperature and pressure was used. In cases where more than one value from different measurements was available, averages were used. b

Number of CO2 molecules used in simulations of ILs consisting of 500 ion pairs, according to measured CO2 saturation concentrations.



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Table 2: Average Ion Charges in Different ILsa ILs BMIM-PF6 BMIM-BF4 BMIM-NO3 BMIM-Cl BMIM-NTF2 C6MIM-NTF2 C8MIM-PF6 C2OHMIM-BF4 C6F9MIM-NTF2 MOEMIM-BF4

qb [e] 0.97 0.95 0.90 0.91 0.99 0.99 0.97 0.96 0.95 0.95

a

Average ion charges take into account electron charge transfer from anions to cations as derived from QM/MM simulations of the corresponding ILs. These total charges were used in the applied force field. b

For ILs which differ only in the length of the cation alkyl chain the same average ion charges were used.


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Table 3: Volumes for Pure ILs, CO2-Saturated ILs, and Partial Molar Volume of CO2a,b ILs BMIM-PF6 BMIM-BF4 BMIM-NO3 BMIM-Cl BMIM-NTF2 C6MIM-NTF2 C8MIM-PF6 C2OHMIM-BF4 C6F9MIM-NTF2 MOEMIM-BF4

VIL [nm3] 176.93 161.01 147.11 137.81 237.61 270.91 237.81 139.21 295.61 151.61

VIL+CO2 [nm3] 182.81 166.11 150.41 n/a 246.71 280.91 247.12 141.91 304.41 n/a

̅ i, CO2 𝑽 [cm3/mol] 362 361 362 n/a 371 35.25 401 322 30.15 n/a

a

At T = 298 K and P = 10 bar.

b

Statistical uncertainties of the last digit are given as subscripts.



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Table 4: Average Interaction Strength of CO2 with Cations and Anions and with Contributions from Coulomb and Lennard-Jones Potentialsa,b,c

Cation ILs BMIM-PF6 BMIM-BF4 BMIM-NO3 BMIM-NTF2 C6MIM-NTF2 C8MIM-PF6 C2OHMIM-BF4 C6F9MIM-NTF2

CO2−Cat ECoul

Anion

CO2−Cat ELJ

-9.207 -8.402 -8.103 -7.905 -7.102 -7.002 -11.601 -8.901

-16.604 -18.504 -20.302 -11.802 -12.905 -17.803 -17.907 -13.404

CO2−An ECoul

-5.608 -8.31 -10.72 -4.803 -4.704 -4.809 -8.50.3 -4.507

a

At T = 298 K and P = 10 bar.

b

All energies are given in units of kJ/mol.

c


Total

CO2−An ELJ

-7.003 -4.503 -7.402 -14.005 -12.505 -5.301 -5.601 -12.305

CO2 ETot

-38.51 -39.91 -46.42 -38.51 -37.31 -34.81 -43.63 -39.11


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Table 5: Ratio of Unoccupied Space in Pure ILsa

ILs BMIM-PF6 BMIM-BF4 BMIM-NO3 BMIM-Cl BMIM-NTF2 C6MIM-NTF2 C8MIM-PF6 C2OHMIM-BF4 C6F9MIM-NTF2 MOEMIM-BF4 a

runocc 0.19403 0.19102 0.1751 0.19103 0.19901 0.20801 0.20501 0.17601 0.19301 0.17804

At T = 298 K and P = 10 bar.



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Table 6: Average Cation – Anion Interaction Strength with Contributions from Coulomb and Lennard-Jones Potential and Interaction Density a,b

ILs BMIM-PF6 BMIM-BF4 BMIM-NO3 BMIM-Cl BMIM-NTF2 C6MIM-NTF2 C8MIM-PF6 C2OHMIM-BF4 C6F9MIM-NTF2 MOEMIM-BF4

𝐂𝐚𝐭−𝐀𝐧 𝐄𝐂𝐨𝐮𝐥 [kJ/mol] -249.73 -262.25 -263.44 -299.46 -226.04 -207.51 -206.25 -329.33 -185.82 -284.56

𝐂𝐚𝐭−𝐀𝐧 𝐄𝐋𝐉 [kJ/mol] -37.43 -24.68 -41.12 -2.62 -79.62 -82.52 -38.28 -17.81 -80.51 -23.99

𝐂𝐚𝐭−𝐀𝐧 𝐄𝐓𝐨𝐭 [kJ/mol] -287.24 -286.99 -304.54 -302.06 -305.54 -290.02 -244.39 -347.13 -266.32 -3081

a

At T = 298 K and P = 10 bar.

b


ɛcat-an [kJ/(mol·nm3)] -8122 -8913 -10351 -10962 -6431 -535.24 -5142 -12471 -450.54 -10183


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Table 7: McGowan Volume of Ion Pairs a,b

ILs BMIM-PF6 BMIM-BF4 BMIM-NO3 BMIM-Cl BMIM-NTF2 C6MIM-NTF2 C8MIM-PF6 C2OHMIM-BF4 C6F9MIM-NTF2 MOEMIM-BF4

VMcG [ml/mol] 174.6 160.2 158.2 147.2 239.8 268.0 230.9 137.9 283.9 152.0

a

At T = 298 K and P = 10 bar.

b

Total volume calculated as the sum of cation and anion volume, using the empirical McGowan method.64



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Figure 1: Structures and abbreviations of all ions considered in simulations. Full chemical names are given in section 2.1. Color code: hydrogen – white, carbon – cyan, nitrogen – blue, oxygen – red, fluorine – pink, chlorine – green, boron - turquois, phosphorus – brown, sulfur – yellow.


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Figure 2: Center-of-mass radial distribution functions, gRDF, of cation – anion pairs in ILs without CO2 (straight line) and with CO2 in saturation concentration (broken line) at T = 298 K and P = 10 bar. The RDFs of ILs before and after CO2 absorption are essentially indistinguishable, thereby demonstrating the negligible influence of CO2 on the overall IL structure.



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Figure 3: Comparison of experimentally determined CO2 saturation concentrations and calculated CO2 - IL interaction strengths in different simulated ILs at T = 298 K and P = 10 bar. A correlation between these two properties was not observed.


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Figure 4: Comparison of experimentally determined CO2 saturation concentrations and calculated ratio of preformed unoccupied space in different simulated pure ILs at T = 298 K and P = 10 bar. A linear regression curve together with R2-value is shown. A large amount of unoccupied space in an IL allows a larger concentration of CO2.



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Figure 5: Comparison of calculated ratio of unoccupied space, runocc, in ILs before CO2 absorption and average CO2 – ion interaction strengths after CO2 absorption at T = 298 K and P = 10 bar. A linear regression curve together with R2-value is shown. Weaker CO2 interactions with ions were observed in ILs with larger amounts of empty space.


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Figure 6: Distribution of empty space in pure BMIM-PF6 (top) and when saturated with CO2 (bottom) in a cross section of a sample structure of the equilibrated IL. Regions of empty space were located with XCav and are shown in orange. CO2 is displayed in green. Ions and CO2 are shown in the van der Waals representation.



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Figure 7: Comparison of calculated ratio of unoccupied space, runocc, and cation – anion interaction density, cat-an, of pure ILs at T = 298 K and P = 10 bar. A linear regression curve together with R2-value that considers all ILs is shown in black, whereas for the curve and R2-value in red BMIM-Cl and C6F9MIM-NTf2 were excluded as explained in the text. ILs with large interaction density tended to exhibit less preformed unoccupied space.


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Figure 8: Comparison of ion densities in simulations and cation – anion interaction densities, cat-an, of pure ILs at T = 298 K and P = 10 bar. A linear regression curve together with R2-value is shown. Values of cat-an are determined by ion density, where large ion densities lead to high interaction densities.



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Figure 9: Comparison of ion densities in simulations with ion densities that were estimated with the McGowan method at T = 298 K and P = 10 bar. A linear regression curve together with R2value is shown. The ion density in simulations can be readily predicted from McGowan volumes of the constituting ions.


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Figure 10: Absorptivity of CO2 is determined in ILs primarily by electrostatic cation – anion interaction strengths (top row). The larger the ions, the lower the ion density and the weaker the cation – anion attraction, which increases the amount of empty space and ease with which ions can be displaced to accommodate CO2 between them. Secondarily, negatively charged parts of cation surfaces due to cation functionalization in the vicinity of other cations lead to additional ion attraction that reduces CO2 absorptivity (bottom left corner). Very small anions result in unavoidable direct cation – cation contacts, whose repulsion ease absorption of CO2 between those cations (bottom right corner).



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Table of contents graphics


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