Solvate Ionic Liquids at Electrified Interfaces - ACS Publications

Subscriber access provided by UNIV OF DURHAM

Energy, Environmental, and Catalysis Applications

Solvate Ionic Liquids at Electrified Interfaces Zhou Yu, Chao Fang, Jingsong Huang, Bobby G. Sumpter, and Rui Qiao ACS Appl. Mater. Interfaces, Just Accepted Manuscript • Publication Date (Web): 29 Aug 2018 Downloaded from http://pubs.acs.org on August 29, 2018

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Applied Materials & Interfaces

Solvate Ionic Liquids at Electrified Interfaces Zhou Yu,1 Chao Fang,1 Jingsong Huang,2,* Bobby G. Sumpter,2 and Rui Qiao1,* 1 2

Department of Mechanical Engineering, Virginia Tech, Blacksburg, VA 24061, USA Center for Nanophase Materials Sciences and Computational Sciences & Engineering Division, Oak Ridge National Laboratory, Oak Ridge, TN 37831, USA

Keywords: solvate ionic liquids, electrical double layer, chelate structure, molecular dynamics, DFT calculation, NBO calculation, nanostructure

*

Address correspondence to [email protected] (R.Q.) and [email protected] (J.H.). 1

ACS Paragon Plus Environment

ACS Applied Materials & Interfaces 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Abstract: Solvate ionic liquids (SILs) are a promising electrolyte for Li-ion batteries; thus, their behavior at electrified interfaces is crucial for the operation of these batteries. We report molecular dynamics simulation results for a prototypical SIL of lithium triglyme bis(trifluoromethanesulfonyl)imide ([Li(G3)][TFSI]) sandwiched between electrified surfaces. At negatively charged as well as neutral electrodes, the electrolyte largely maintains the characteristics of SILs, e.g., a majority of the interfacial Li+ ions remains coordinated by the similar number of oxygen atoms on G3 ligands as the bulk Li+ ions. The persistence of the complex ions is attributed to the 1:1 Li-G3 ratio in bulk SILs and the fact that G3 molecules readily adapt to the interfacial environment, e.g., by aligning themselves with the surface to ensure good solvation of the interfacial Li+ ions. Nevertheless, the interfacial Li+ ions also display changes of solvation from that in bulk SIL by deviating from the molecular plane formed by the oxygen atoms on G3 ligands as electrodes become more negatively charged. Using density functional theory along with natural bond orbital calculations, we examine the effects of such structural distortion on the properties of the complex cation. Both the frontier orbital energies of the complex cation and the donor-acceptor interactions between Li+ ions and G3 ligands are found to be dependent on the deviation of Li+ ions from the molecular plane of the G3 ligands, which suggest that the reduction of Li+ ions should be facilitated by the structural distortion. These results bear important implications for the nanostructures and properties of SILs near electrified interfaces during actual operations of Li-ion batteries and serve to provide guidance toward the rational design of new SILs electrolytes.

2


Page 2 of 32

Page 3 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1. INTRODUCTION Due to their remarkable properties including negligible vapor pressure, non-flammability, and high ionic conductivity,1-5 ionic liquids (ILs) have attracted much attention for applications in diverse fields such as energy storage,3 green solvents,4 and drug delivery.5 ILs can be roughly divided into four categories: aprotic, protic, inorganic, and solvate ILs.6 The novel solvate ionic liquids (SILs) are different from the other members of the ILs family in that they are concentrated mixtures of salts and solvents. In a recent review,7 Watanabe and coworkers proposed a series of criteria for concentrated mixtures to be classified as SILs instead of concentrated solutions. Based on the proposed criteria, many equimolar mixtures of alkali metal salts and oligoethers, or socalled glymes, can be considered as SILs.7 Glymes are often abbreviated as Gn, where n depends on the repeating units in the formula of CH3O(CH2CH2O)nCH3. The multiple alkoxy oxygen atoms in glymes make them excellent polydentate ligands. For instance, in equimolar mixtures of Li[TFSI] salt and Gn (e.g., G3 for triglyme or G4 for tetraglyme), Gn molecules coordinate with Li+ ions to form stable [Li(Gn)]+ complex cations featuring a coplanar chelate structure.8-9 These mixtures are composed entirely of [Li(Gn)]+ complex cations and TFSI- anions, which endow them with key characteristics of ILs and make [Li(Gn)][TFSI] prototypical SILs. Experimental investigations have been carried out to explore the fundamental properties and potential applications for many SILs. SILs with different anions (e.g., TFSI-, BETI-, and OTf-),10 cations (e.g., Li+, Na+, and K+),11 and solvents (e.g., cyclic tetrahydrofuran and acyclic G1 through G4)12-13 have been extensively examined. From these studies, it has been found that [Li(Gn)][TFSI] can serve as a promising electrolyte for electrochemical devices such as Li-ion and Li-S batteries because they offer advantages over conventional RTILs in terms of high Li+ ion concentration, high Li transference number, strong oxidative stability, and ability to suppress the dissolution of

3



lithium polysulfides.7-14 For example, recent experiments suggested that the oxidative stability of Gn molecules is enhanced by their complex formation with alkali metal cations. 9 In addition, it was shown that Li+ ion transport in the electrolyte proceeds mainly through the migration of the [Li(Gn)]+ complex cation, along with a Grotthuss-like ligand exchange mechanism induced by electrochemical reactions at the solid-liquid interface.9 Nevertheless, these SILs still demonstrate good ionic conductivity up to 10-3 S cm-1.15 Further, it was also reported that these SILs can suppress the dissolution of lithium polysulfides and thereby ensure the stable operation of a Li-S battery over 400 cycles with a high discharge capacity of 700 mAh g-1 and a Coulombic efficiency of 98%.14 In parallel with the experimental investigations, simulation studies have been performed to clarify the structural and dynamic characteristics of [Li(Gn)][TFSI] SILs. Molecular dynamics (MD) modeling is an effective tool for the studies of SILs when reliable force fields are available.16-18 MD simulations based on force fields built under the framework of the AMBER and OPLS force fields19 have been shown to be successful in predicting the structure and dynamics of bulk SILs. The structure factor determined in these simulations agrees well with that obtained from diffraction and scattering experiments.16 The coordination environments of Li-Gn and Li-anion in bulk SILs have also been investigated using MD simulations. These simulations confirmed the formation of [Li(Gn)]+ complex cations in bulk SILs.16-18 The computed coordination number of Li+ ion in [Li(G4)]+ agrees well with experimental measurement using the Li isotopic substitution method.18 Furthermore, the trend of the dynamic property of ([Li(Gn)][TFSI]) computed in these simulations is also consistent with the experimental results.16-17 In addition to MD simulations, quantum chemical and density functional theory (DFT) methods have also been used to study SILs, e.g., by optimizing the geometry of the complex

4


Page 4 of 32

Page 5 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


cations and probing the redox mechanism.9,

20

For instance, the crown-ether-like chelate

conformation of the [Li(Gn)]+ complex cation was obtained by geometry optimizations and vibrational analyses.9, 20 Geometry optimizations also indicated that Li+ ions tend to be coordinated by all four oxygen atoms in G3 molecule and two oxygen atoms in the TFSI- ion in the gas phase of [Li(G3)][TFSI].17 The calculations of frontier orbitals of [Li(Gn)][TFSI] showed that the coordination of Li+ ions can considerably lower the highest occupied molecular orbital (HOMO) energy levels of Gn molecules, which is consistent with the experimental observation that the formation of [Li(Gn)]+ complex cation improves the oxidative stability of the Gn molecules.9 On the other hand, the lowest occupied molecular orbital (LUMO) is associated with the electrochemical reduction of reactive species. Although the existing theoretical calculations on SILs provide useful insight, they were performed for structures optimized in gas phase without considering electric double-layer (EDL) effects. These prior studies on the nanostructure, transport, and electrochemical properties of SILs have greatly advanced our understanding of SILs in the bulk. However, the studies of SILs at interfacial layers are still scarce, despite the fact that the application of SILs in electrochemical systems depends critically on their behavior at electrode-SIL interfaces. Although the structure of interfacial electrolytes has been extensively investigated,21-26 limited simulations have been reported for SILs at electrode interfaces.27-28 For example, the structure of [Li(G4)][TFSI] near graphene surfaces have been studied using MD simulations, in which the layering of SILs, adsorption (desorption) of counterions (co-ions), and the coordination environment of Li+ ions at the negative electrodes were reported.28 Overall, however, the behavior of SILs at electrode interfaces is not well understood and many questions remain open. For example, what is the structure of the SILs near electrodes with different surface charge densities? What are the

5



differences of the coordination environments between SILs in the bulk and at electrified interfaces? Would the complex cations undergo structural distortions near electrode surfaces from their ideal coplanar structure and if any, how do the structural distortions affect the frontier orbital characters of [Li(Gn)]+ and the interactions between Li+ ions and G3 ligands? Motivated by these questions, in this work we combined MD simulations and DFT calculations to study the nanostructures and properties of [Li(G3)][TFSI] at the electrified interfaces. 2. SIMULATION SYSTEM AND METHODS MD System. Figure 1 shows a schematic of the simulation system, which consists of two graphene electrodes and a slab of SILs. The electrodes measure 4.68×4.91 nm2 in the xy-plane and their spacing in the z-direction is 8.0 nm, which is wide enough to produce a bulk like SIL in the middle of the system. A large vacuum is introduced beyond the simulation region sandwiched by the two electrodes so that the periodic box measures 24 nm in the z-direction. In different simulations, small partial charges are distributed uniformly on all atoms of the two electrodes to produce a surface charge density (σ) of 0, ±0.12, or ±0.24 C/m2. These surface charge densities are selected to be relevant to the operation of Li-ion batteries. Li-ion batteries have an equilibrium voltage of ~3.7 V. Assuming a capacitance of 0.1 F/m2 for the double layers near the electrodes, which is typical for double layers in conventional RTILs,1 the equilibrium surface charge density of these electrodes can be estimated as ~±0.18 C/m2. Since batteries operate at voltage higher (or lower) than the equilibrium voltage during charging (or discharging), the charge densities of ±0.12 and ±0.24 C/m2 are smaller and greater than the estimated value at equilibrium, and therefore are relevant to the Li-ion battery charging and discharging, respectively. In all simulations, the slab of SILs consists of 323 G3 molecules and 323 pairs of Li+ and TFSI- ions.

6


Page 6 of 32

Page 7 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 1. (a) A schematic of the simulation system composing of a slab of (b) [Li(G3)]+complex cations and (c) [TFSI]- anions sandwiched between two charged electrodes. In (a), the blue, red, green, and gray spheres denote TFSI- ion, G3 molecule, Li+ ion, and electrode atoms, respectively. In (b-c), the magenta, red, gray, white, blue, yellow, and cyan spheres denote the lithium ion, oxygen, carbon, hydrogen, nitrogen, sulfur, and fluorine atoms, respectively. The simulation system, cations, and anions were visualized using the visual molecular dynamics (VMD) package.29

Molecular Models. We adopt [Li(G3)][TFSI] as the model SIL because recent studies have shown that the G3 molecule shows a better oxidative stability than G4 molecule in Li batteries.9 We used all-atom models for [Li(G3)][TFSI] based on the molecular geometry optimized by a quantum chemical method at the HF/6-311G** level.9 We adopted the force fields developed by Shimizu et al.15 for the [Li(G3)][TFSI] molecules.19 MD simulations based on these force fields have been shown to accurately reproduce the nanostructure of bulk [Li(G3)][TFSI] (e.g., coordination of Li+ ions by the G3 molecules) measured experimentally.16 The electrode atoms are modeled as carbon atoms with the same Lennard-Jones parameters as those of the sp2 carbons in the Amber force fields.30 MD Methods. Simulations were performed using the Gromacs code31 in the NVT ensemble. A system temperature of 400 K, which is below the decomposition temperature of

7



[Li(G3)][TFSI],32 was used. This elevated temperature was chosen to ensure effective sampling of the phase space, as common in simulations of ILs. The system was evolved using a time step of 2 fs using the velocity Verlet integrator. The system temperature was maintained by the velocity rescaling thermostat with a time constant of 2ps.33 The non-electrostatic interactions were computed by direct summation with a cut-off length of 1.2 nm. The electrostatic interactions were computed using the Particle Mesh Ewald (PME) method. The real space cutoff and FFT spacing were set to 1.2 and 0.12 nm, respectively. All bonded interactions were computed except that the length of the C-H bonds were constrained using the LINCS algorithm.34 To remove the interactions between periodic images in the z-direction, the slab correction to the PME method was applied.35 All systems were first run for 10 ns, when the density distribution of molecules in the system reached an equilibrium. Then a production run of 40 ns followed. DFT Calculations. Starting from the previously reported molecular geometry,9 the structure of [Li(G3)]+ complex cation was optimized at the DFT level with the B3LYP/6-311G** method, 36-38

using the Gaussian 09 program.39 The energies of delocalized frontier orbitals HOMO and

LUMO were obtained from the optimized geometry. Following the geometry optimization, Natural Bond Orbitals (NBO) calculations were carried out using the NBO 3.1 program40 interfaced to Gaussian 09 program. Within the NBO calculations, the delocalized molecular wave functions were transformed into localized one-center lone-pair (LP) orbitals, and second-order perturbation theory analyses were performed to calculate the donor-acceptor interactions between Li+ ion and G3 ligand. To study the effect of structural distortion of the [Li(G3)]+ complex cation from the ideal coplanar geometry, the frontier HOMO and LUMO orbital energies and donor-acceptor interactions between Li+ ion and G3 ligand were rigidly scanned as a function of Li+-G3 distance Z, defined as the distance from the Li+ cation to the molecular center of G3 ligand. To that end, a

8


Page 8 of 32

Page 9 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


plane is fitted based on the four O atoms of G3 molecule through a singular value decomposition method and then the Li+ ion is pulled out of the G3 ligand along the direction normal to the fitted plane. The atomic coordinates for the [Li(G3)]+ complex cation in a range of z= 0 to 3.0 Å with a step size of 0.25 Å were deposited in the Supporting Information. Both delocalized molecular orbitals and localized LP orbitals were visualized using the VMD 1.9.3 software.29

3. RESULTS AND DISCUSSION Near extended surfaces, conventional ILs are well-known to exhibit distinct structures from their bulk.3, 41-46 Below we first examine the structure of [Li(G3)][TFSI] near electrode surfaces with different surface charges, i.e., the out-of-plane layering of a SIL, the in-plane ordering of the [Li(G3)]+ complex cation, and the coordination of Li+ ions by the G3 molecules. Based on the results from MD simulations, we describe how the structural distortions affect the properties of [Li(Gn)]+ complex cations. Out-of-plane layering. Figure 2 shows the number density profiles of the Li+ ion, the TFSIion, and the G3 molecule in the three systems with electrode charge density of 0, ±0.12, and ±0.24C/m2. The position of the TFSI- ion is based on its nitrogen atom. The position of the G3 molecule is based on only one of its oxygen atoms (atom O2 in Fig. 1), since representing the G3 molecule using other atoms (e.g., atom O1 in Fig. 2) leads to nearly the same density profiles (see Fig. S1 in the Supporting Information). In the neutral system, the distributions of ions and G3 molecules near the two electrodes are identical (see Fig. 2a and 2b). Both ions and G3 molecules show notable density peaks near the electrodes due to the well-known layering effect but their densities become largely uniform about 2 nm from the electrode surface, similar to those reported for [Li(G4)][TFSI] near neutral graphene

9



surfaces.28 Below, we focus on the evolution of the interfacial structure as the electrode charge density changes.

Figure 2. Number density profiles of the Li+ ion, the nitrogen atom of the TFSI- ion, and the O2 atom of the G3 molecule (see Figure 1) in three simulation systems with electrode surface charge densities of 0 (a, b), ±0.12 C/m2 (c, d) and ±0.24 C/m2 (e, f). The surfaces of the left and right electrodes are located at z = 0 and 8 nm, respectively.

As positive charge densities are brought to the electrode surface, the first TFSI- ion density peak near the electrode becomes higher and approaches closer to the electrode while the first Li+ ion peak is pushed away from the electrode (see Fig. 2b, d, and f). At 𝜎 = 0.12 C/m2, the TFSIions and Li+ ions form alternating layers near the electrode, which is one of the hallmarks of EDLs in conventional ILs.3, 41, 47-50 At 𝜎 = 0.24 C/m2, the first layer of TFSI- ions can no longer fully screen the surface charge, and a second layer of TFSI- develops at z = 7.12 nm. Such a phenomenon is similar to the lattice saturation effects reported for EDLs of conventional ILs near highly charged electrodes, where the first, close-packed layer of bulky counterions adsorbed on the electrode is not enough to fully screen the surface charge and a second counterion layer is formed next to it.51-

10


Page 10 of 32

Page 11 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


52

This crowding-induced formation of a second TFSI- layer near the electrode occurs at a charge

density lower than that in conventional RTILs53-55 (e.g., the critical surface charge density has been estimated to be ~0.5 C/m2 in a recent study).54 This difference is attributed to fact that, unlike in prior works, the first layer of TFSI- ions near the electrode are solvated by the G3 molecules in the first G3 peak near the electrode. We also observe that G3 molecules form a distinct layer adjacent to the electrode and the density peak of this layer becomes sharper as 𝜎 increases. At 𝜎 = 0.12 C/m2, the first G3 peak is ~0.36 nm away from the first Li+ peak, indicating that G3 molecules in this peak are no longer coordinated to Li+ ions. The large separation of adsorbed G3 from Li+ ions becomes more distinct at 𝜎 = 0.24 C/m2. This phenomenon implies that the electrolyte near the electrode may lose the characteristics of SILs at the surface charge densities considered here. We caution that these predictions are based on the force fields parameterized to reproduce the behavior of bulk SILs and therefore the performance of these force fields for interfacial SILs remain to be established rigorously. Nevertheless, the above results point to the need to quantify the distribution of interfacial G3 molecules experimentally. As negative surface charge densities are introduced, the distribution of the interfacial Li+ ions also changes greatly. At 𝜎 = 0, the interfacial Li+ ions mostly accumulate in the density peak at z = 0.63 nm and the G3 density is high both above and below this Li+ layer, indicating that these Li+ ions are well solvated by G3 molecules. At 𝜎 = -0.12 and -0.24 C/m2, significant contact adsorption of Li+ ions on the electrode surface occurs. The contact adsorption observed here was not reported in the prior simulations of [Li(G4)][TFSI] at electrified interfaces.28 This is likely caused by two reasons. First, the G4 ligands have more oxygen atoms to coordinate with the Li+ ions, and therefore G4 molecules bind more strongly with the Li+ ions than G3 molecules.9, 16-17 Hence it is more difficult for Li+ ions to weaken its coordination with G4 to become contact adsorbed. Second,

11



the surface charge density in our simulation is higher (-0.12 and -0.24 C/m2 compared to -0.05 and -0.10 C/m2 in the prior studies28), which facilitates the contact adsorption of Li+ ions. The separation between the contact adsorbed Li+ ions and the first G3 density peaks is 0.15~0.20 nm in our systems (cf. Fig. 2c and 2e). This separation is much smaller than that between the first Li+ and G3 peaks near the positively charged electrodes 0.36~0.85 nm (cf. Fig. 2d and 2f). Therefore, although the contact adsorbed Li+ ions are somewhat pulled away from their solvating G3 molecules (see below), the interfacial G3 and Li+ ions near the negatively charged electrodes are better coordinated than those near the positively charged electrodes. The above results clearly show that the coordination in the [Li(G3)]+ complex cations is modified near electrified surfaces. The modification is less distinct near the negatively charged electrodes: the [Li(G3)]+ complex cations likely persist near even highly charged negative electrodes, where the electrolytes largely maintain the characteristics of SILs. Therefore, in the rest of this work, we focus on the interfacial structures near the negatively charged electrodes and compare to those in the bulk and near the neutral electrode. Li+ ion coordination. The out-of-plane layering of SILs near electrode surface shown above is accompanied by the change of coordination of interfacial Li+ ions by G3 molecules. Hence, we analyze the Li+ ion coordination by focusing on the coordination number of the interfacial Li+ ions, their three-dimensional details, and the conformation of the interfacial G3 molecules. We first computed the coordination number of the interfacial Li+ ions by counting the number of G3’s oxygen atoms coordinating each ion. Interfacial Li+ ions are defined as the Li+ ions in the first peak next to the electrode surface (see Fig. 2; near the neutral electrodes, Li+ ions in the broad first peak located within 0.88 nm from the surface are considered to improve statistics). A G3 molecule’s oxygen atom is considered to be coordinating a Li+ ion if its distance to the ion is less 12


Page 12 of 32

Page 13 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


than 0.36 nm, which corresponds to the location of the first minimum of the Li-oxygen radial distribution function in bulk [Li(G3)][TFSI] (see Fig. S2 in the Supporting Information).

Figure 3. (a) The average number of oxygen atoms of G3 molecules coordinating the interfacial Li + ions near neutral and negatively charged electrodes. (b1-b3) The frequency of different number (0, 1, and 2) of G3 molecules within the solvation shell of the interfacial Li+ ions. The corresponding coordination numbers and frequencies for bulk Li+ ions are marked using dashed lines and their values are labeled in red. The Li+ ions in the first density peak near electrodes are taken as the interfacial ions (for 𝜎 = 0: z < 0.88 nm; for 𝜎 = −0.12 C/m2: z < 0.38 nm; for 𝜎 = −0.24 C/m2: z < 0.28 nm).

Figure 3a shows that the interfacial Li+ ions near neutral surface have a coordination number of 4.28, which is slightly higher than that of bulk Li+ ions (3.62). To further examine the coordination of interfacial ions, we evaluated the local solvent structure near a Li+ ion located at the first density peak (z = 0.63 nm, cf. Fig. 2a). Specifically, we computed the density of G3 molecules’ oxygen atoms as a function of their radial distance r from the Li+ ion in the xy-plane and their distance z from the electrode surface. Previous work on the solvation of interfacial ions by RTILs suggested that the cylindrical averages of the three-dimensional solvation structure of interfacial ions can offer valuable insight difficult to obtain in the average solvation number.23, 26 Figure 4a shows that the oxygen atoms of G3 molecules surround the Li+ ions from all directions, i.e., the electrode surface does not impede G3’s oxygen atoms from solvating the ion. This observation, along with fact that G3 molecules are enriched near the electrode surface due to the layering effect, leads to a higher coordination number for the interfacial Li+ ions than the bulk Li+

13



ions. Because all four oxygen atoms of the G3 molecules reside essentially in one plane (see Fig. S3 and S4 in the Supporting Information), the oxygen atom distribution around Li+ ion shown in Fig. 4a suggests that the interfacial G3 molecules solvating the Li+ ion can orient their plane relatively randomly with respect to the electrode surface, which is indeed observed in our simulations (see Fig. S4a in the Supporting Information). When the surface charge density of the electrode changes to -0.12 C/m2, the coordination number of the interfacial Li+ ions decrease slightly to 4.00, which is still 10% higher than that of bulk Li+ ions. This is an interesting observation. Because the interfacial Li+ ions are now contact adsorbed on the electrode surface (cf. Fig. 2c), no oxygen atom can solvate the Li+ ion from the electrode side (see Fig. 4b) and thus one could have expected that the coordination number of Li+ ions to reduce greatly. Indeed, such contact adsorption-induced reduction of ion solvation has been observed for the counterions at organic electrolytes-electrode interfaces.21 In the context of interfacial Li+ ions, their contact adsorption is not favored near neutral or moderately charged electrodes in contact with conventional ILs such as [BMIM][BF4]22-23 or polymer electrolytes.24 However, their contact adsorption has been reported near negatively charged electrodes in contact with mixed-solvent electrolyte tetramethylene sulfone/dimethyl carbonate, and notable reduction of their solvation was also observed.25 The weak reduction of coordination number as Li+ ions in SILs are contact adsorbed on surfaces can be ascribed to the following reason. The coordination number of a Li+ ion is ~3.62 in bulk SILs because of the 1:1 ratio between G3 molecule and Li+ ion. Therefore, when a Li+ ion moves from bulk to the electrode surface, if the G3 molecule coordinating it adopts a co-planar conformation with respect to the surface, then the Li+ ion can maintain a similar coordination number as in bulk. The postulated co-planar conformation of the G3 molecules solvating the Li+ ion is consistent with the calculated oxygen atom distribution near

14


Page 14 of 32

Page 15 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


the ion (see Fig. 4b) and is supported by the calculated orientation of the plane formed by its oxygen atoms (for brevity, this plane is hereafter termed the plane of the G3 molecule) with respect to the electrode surface (see Fig. S3 in the Supporting Information).

Figure 4. Distribution of the oxygen atoms of G3 molecules around interfacial Li+ ions near electrodes with different charge densities: (a) near a Li+ ion positioned 0.63 nm above a neutral electrode; (b) near a Li+ ion positioned 0.23 nm above an electrode with 𝜎 = -0.12 C/m2; (c) near a Li+ ion positioned 0.18 nm above an electrode with 𝜎 = -0.24 C/m2. The three positions of the Li+ ions correspond to their highest peak near the respective electrodes (cf. Figure 2 a, c, and e) and are labeled using a filled white circle. The oxygen atoms within the dashed white circle (radius: 0.36 nm) are considered to be coordinating the Li + ion. The oxygen density is color-coded, and the unit of the color bar is nm-3.

When the surface charge density of the electrode changes to -0.24 C/m2, the coordination number of interfacial Li+ ions drops significantly to 2.34 (65% of the bulk value, see Fig. 3a). As shown in Fig. 4c, the oxygen atoms solvating the Li+ ions reside in a plane slightly higher than the Li+ ion. Hence, the G3 molecules solvating the Li+ ion should adopt a co-planar conformation with respect to the electrode surface (see Fig. S3 in the Supporting Information). The distinct reduction of the Li+ ion’s coordination number in this case is due to the crowding of the [Li(G3)]+ complex cations near the highly charged electrode surface. The [Li(G3)]+ complex cation is bulky and has a diameter of ~0.81 nm in the plane formed by the oxygen atoms of the G3 molecule (see Supporting Information). Using the ion density profile in Fig. 2e, the density of Li+ ions adsorbed on the electrode surface is found to be 2.2 nm-2. Since the complex cations are adsorbed on the

15



surface in a planar conformation (cf. Fig. S3), this density exceeds the close packing limit of [Li(G3)]+ ions on the electrode surface (1.77 nm-2 based on a diameter of 0.81 nm for the complex cation). Therefore, near the highly charged electrode surface, some interfacial Li+ ions must lose their coordination by oxygen atoms and become partially desolvated. Another aspect of the coordination of Li+ ions is how many G3 molecules are coordinated to each Li+ ion. To clarify this aspect, we computed the number of G3 molecules that simultaneously coordinate interfacial and bulk Li+ ions. Following the prior study of bulk [Li(G3)][TFSI],16 a G3 molecule is considered to be coordinating a Li+ ion if at least one of its oxygen atoms is coordinating the ion. Using this criterion, we found that a Li+ ion can be coordinated by up to two G3 molecules, and the frequency of different numbers (0, 1, and 2) of G3 molecules coordinating bulk and interfacial Li+ ions is shown in Fig. 3 (b1-b3). Similar to the bulk Li+ ions, the interfacial Li+ ions near electrodes with 𝜎 = 0 and -0.12 C/m2 are mostly coordinated by one G3 molecule. Occasionally (~16.5% chance), a Li+ ion near neutral electrodes can be coordinated by the oxygen atoms of two G3 molecules. This is likely because the two dense layers of oxygen atoms flanking the interfacial Li+ ions can provide coordination to these ions (see Fig. 2a). As the surface charge density increases to -0.24 C/m2, 17.7% of the interfacial Li+ ions are no longer coordinated by any G3 molecule. The desolvation of these Li+ ions is due to the crowding effect described above. The remaining solvated interfacial Li+ ions are each coordinated by almost all four oxygen atoms of a G3 molecule. This result suggests that, even near a highly charged electrode, a majority of the interfacial Li+ ions remain chelated by G3 molecules. Figure 3 shows that a majority of the interfacial Li+ ions maintains their solvation shell, which mainly consists of the oxygen atoms from a single G3 molecule. While their coordination number could be the same as the bulk Li+ ions, their solvation by the G3 molecules can deviate from those

16


Page 16 of 32

Page 17 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


in bulk. Indeed, the oxygen atom distribution shown in Fig. 4c shows that an interfacial Li+ ion is pulled away from the plane of the G3 molecule solvating it. To quantify this effect, for each interfacial Li+ ion solvated by three or more oxygen atoms of a single G3 molecule, we compute its distance to the plane of that G3 molecule. Figure 5 shows that, near neutral electrodes, the interfacial Li+ ions mostly reside within G3 molecules’ plane, just like the Li+ ions in bulk. Near electrodes with 𝜎 = -0.12 C/m2, on average, Li+ ions are pulled out of the G3 molecule’s plane by 0.08 nm, although some Li+ ions still reside in G3’s plane. As the electrode is electrified even more to 𝜎 = -0.24 C/m2, the interfacial Li+ ions are pulled further out of the G3’s plane by an average distance of 0.14 nm. This observation of structural distortions near charged electrode surfaces is further examined below with DFT calculations.

Figure 5. The distance between an interfacial or bulk Li+ ion and the plane of a G3 molecule solvating it. The plane of the G3 molecule is fitted from its four oxygen atoms.

The structure of interfacial SILs presented above is obtained near relatively ideal electrodes where only the surface charge density of an electrode affects interfacial ions greatly. Nevertheless, the insight gained here can help one construct a reasonable picture on how the molecular details of realistic electrodes affect the solvation and distribution of the interfacial Li+ ions. For example,

17



in presence of hydrophilic functional groups on the electrode, notable contact adsorption of Li + ions should occur before the surface charge density decreases to -0.12 C/m2 because the “solvation” of the Li+ ion by the hydrophilic surface functional groups can compensate partly the penalty caused by reduction of solvation by the G3 molecules. On the other hand, because the functional groups protruding from the electrode surface can compete with G3 molecules for the space near the contact adsorbed Li+ ions, the G3 molecules solvating the Li+ ions may not be able to adopt the ideal co-planar chelate structure shown in Figure 5. In-plane ordering. The in-plane ordering of interfacial ions, i.e., the arrangement of ions in the horizontal plane parallel to the electrode surface, can affect both the thermodynamic properties of the EDLs and the dynamics of these ions.43, 56-57 Here, we probe the in-plane ordering of the first layer of Li+ ions near neutral and negatively charged electrodes by quantifying the uniformity of their packing in the horizontal plane. Specifically, we performed a Voronoi tessellation of the first layer of Li+ ions projected onto the electrode surface. Figure 6a shows a representative Voronoi tessellation of the first layer of Li+ ions near the electrode with 𝜎 = -0.24 C/m2. The dots denote the Li+ ions and each tessellated cell represents the space associated with a Li+ ion. Figure 6b shows the distribution of the areas associated with the Li+ ions near electrodes with three different surface charge densities (note that the area associated with individual ions (𝐴𝑖 ) is normalized by the average area of all ions (𝐴̅ )). The distributions of the area associated with Li+ ions near electrodes with 𝜎 = 0 and -0.12 C/m2 are similar. As the surface charge increases to -0.24 C/m2, the distribution becomes narrower, indicating that the interfacial Li+ ions become more ordered. The transition of the interfacial Li+ ions from a relatively disordered structure at zero and low electrode charge to the more ordered

18


Page 18 of 32

Page 19 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


structure at high electrode charge can be understood from the packing of the [Li(G3)]+ complex cations near the electrode surfaces.

Figure 6. (a) Representative Voronoi tessellation of the Li+ ions contact adsorbed on an electrode with 𝜎 = -0.24 C/m2. The dots represent the position of the Li+ ions. (b) The relative frequency of Voronoi cell area associated with individual Li+ ions normalized by the average area. The contact adsorbed Li+ ions are located in the first peak at 0~0.48 nm, 0~0.38 nm, and 0~0.28 nm from electrodes with 𝜎 = 0, -0.12, and 0.24 C/m2.

Our above analysis of the coordination of interfacial Li+ ions show that a majority of these ions remain solvated and their solvating G3 molecules align their plane largely parallel to the electrode surface. Near electrodes with 𝜎 = 0 and -0.12 C/m2, the density of the first layer of Li+ ions (thus the [Li(G3)]+ complex cations) is ~0.3 and 0.9 nm-2, respectively. Because the diameter of a [Li(G3)]+ complex cation is ~0.81 nm, the neighboring [Li(G3)]+ ions do not approach each other closely. Therefore, the distribution of the interfacial [Li(G3)]+ ions (thus the Li+ ions) in the horizontal plane shows weak ordering. Near electrodes with -0.24 C/m2, however, the density of Li+ ions reaches 2.2 nm-2 so that the [Li(G3)]+ ions formed by these ions approach each other closely. The repulsive interactions between these [Li(G3)]+ ions force them to form more ordered structure akin to the two-dimensional Wigner crystals.58

19



DFT Calculations. The MD results presented above reveal that the complex cations of [Li(G3)]+ could survive the negatively charged electrodes by maintaining the SILs characteristics. However, unlike the structure in the bulk and near neutral electrode surfaces, the [Li(G3)]+ complex cation near the negatively charged electrode surfaces displays a structural distortion from its ideal coplanar conformation. To understand how the structural distortion affects the frontier orbital characters of [Li(G3)]+ and the interactions between Li+ ion and G3 ligand, and especially its relevance to the actual operations of Li-ion batteries, in the following we discuss DFT results along with NBO calculations for the [Li(G3)]+ complex cation. Figure 7 shows the variations of the HOMO and LUMO energies of [Li(G3)]+ as a function of the distance Z from Li+ ion to the center of a G3 molecule. Note that the energies of the frontier orbitals are affected by multiple factors including, but not limited to, electrostatic potentials of the added charges on the electrode surface, the electric field strength near EDLs, screening effects from surrounding molecules or ions, and computational methods. Herein we examine only the qualitative trend of the energy change as a [Li(G3)]+ complex cation undergoes a structural distortion by the broadly used hybrid B3LYP functional. It is found that the HOMO energy is lifted while the LUMO energy is lowered with increasing Li+-G3 distance Z. As a result, the HOMOLUMO gap is reduced. However, this does not mean that the electrolyte stability window is narrowed. Near the surface of a negative electrode in a Li-ion battery, oxidation reactions should take place on the HOMO of LiC6 species in the carbon electrode instead of the HOMO of [Li(G3)]+ in the electrolyte during discharge process, and reduction reactions should take place on the LUMO of [Li(G3)]+ in the electrolyte during charge process. In the former process lithium is oxidized and extracted as Li+ from the C6 matrix and solvated by the glyme ligands to form complex ions in the electrolyte, while in the latter process Li+ in the electrolyte is reduced and intercalated into the C6

20


Page 20 of 32

Page 21 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


matrix. Therefore, the lifting of the HOMO energy does not make the complex ion to be oxidized more easily with the structural distortion. On the other hand, the lowering of the LUMO energy implies that during the charge process, the reduction of Li+ ions is facilitated by the structural distortion. The frontier orbitals are further visualized to verify that the LUMO is associated with the Li+ ion instead of the G3 ligand. For the ideal coplanar structure (Z = 0 Å), both HOMO and LUMO are delocalized on the G3 ligand (Fig. S5 in the Supporting Information). In contrast, with the structural distortion (Z = 1.5 Å), the LUMO is no longer delocalized on the G3 ligand but is mainly localized on the Li+ ion, although the HOMO remains so (Fig. S6 in the Supporting Information). This shows that the Li+ ions of complex cations with structural distortion are ready to be reduced and intercalated into negative electrode materials.

Figure 7. Rigid scan of HOMO and LUMO energies as a function of Li+-G3 distance Z, calculated using the B3LYP/6-311G** method for the [Li(G3)]+ complex cation optimized by the same method. The inset shows the top and side views of the [Li(G3)]+ complex cation with Z = 0 Å. The magenta, red, gray, and white spheres denote lithium ion, oxygen, carbon, and hydrogen atoms. The complex cation was visualized using the VMD package.29

We also carried out NBO calculations to gain a deeper insight into the interactions between Li+ ions and G3 ligands as the complex cation undergoes a structural distortion. In line with the

21



Lewis acid-base concept59 and the classification introduced by Pearson,60 the Li+ ion is considered a hard Lewis acid due to its empty orbitals while the G3 ligand molecule is a hard Lewis base due to the presence of lone pair (LP) orbitals on the O atoms. Each of the four O atoms (O1-O4) contributes two LP orbitals, making a total of eight LP orbitals. According to our NBO calculations, all of the LP orbitals display local sigma-pi separations (Fig. S7 in the Supporting Information), just like those observed in a water molecule and other divalent oxygen-containing molecules.61-62 The four sigma-type LP orbitals adopt a sp1.4 hybridization, showing a slightly lower p-component than typical sp2 hybridized orbitals. In comparison, the four pi-type LP orbitals are nearly 100% pure p-orbitals. Due to the presence and absence of the s-components, respectively, the sigma-type LP orbitals (ca. -20 eV in energy) are more stable than the pi-type LP orbitals (ca. -13 eV in energy). The main difference between the sigma- and pi-type LP orbitals manifests in that the former lie in the local C-O-C planes while the latter are perpendicular to the local C-O-C planes. It is also worth pointing out that although the four O atoms are nearly coplanar, the LP orbitals on the four O atoms display different orientations, even within the same sigma- or pi-type orbitals (Fig. S7 in the Supporting Information). Figure 8 shows the donor-acceptor interactions between Li+ ion and G3 ligand as a function of Li+-G3 distance Z, calculated by second-order perturbation theory analysis. For this analysis, we adopted a finer step size of 0.25 Å to avoid missing any details of energy changes. The secondorder interaction energies E(2) characterize the stabilization of the [Li(G3)]+ complex cation as a result of the delocalization from the occupied Lewis LP orbitals on O atoms to the empty nonLewis orbitals on Li+ ion, which is related to the Lewis acid-base binding. The stabilization is dominated by the interactions between Li+ and the sigma-type LP orbitals of O atoms, resulting

22


Page 22 of 32

Page 23 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


from their preferential orientations and therefore their better overlap with the empty non-Lewis orbitals on Li+ ion. The stabilizations for the pi-type LP orbitals are on the same order of hydrogen

Figure 8. Rigid scan of donor-acceptor interactions between Li+ ion and G3 ligand as a function of Li+-G3 distance Z, obtained by second-order perturbation theory analysis in Natural Bond Orbital calculations, using the B3LYP/6-311G** method for the [Li(G3)]+ complex cation optimized by the same method. The numberings of O atoms (O1-O4) and the sigma- and pi-type LP orbitals can be seen in Figure S7 in the Supporting Information.

bonds. For the Z = 0 structure, the E(2) energies are identical for O1/O4 and O2/O3 due to the C2 point group symmetry of the complex cation. However, as the complex cation undergoes a structural distortion, the symmetry is broken. The E(2) energies of the sigma-type LP orbitals essentially show monotonic decrease with increasing distance Z, except for O1 and O3 at Z = 0.25 Å where the subtle difference arises from their preferential orientation (Fig. S7a in the Supporting Information). In comparison, all of the E(2) energies of the pi-type LP orbitals display some oscillations in the small Z range until 0.75 Å, after which all decrease monotonically with increasing distance Z. The anomalous variations at small Z range can be rationalized by the different orientations of the pi-type LP orbitals (Fig. S7b in the Supporting Information), while the monotonic variations at large Z range are mainly dictated by the distance Z. Based on these NBO

23



results, it can be seen that the Li+-G3 interactions are significantly weakened as the complex cation undergoes a structural distortion from the ideal coplanar conformation. This effect also contributes to facilitating the reduction of Li+ at the negative electrode during charge process. 4. CONCLUSIONS In summary, we have studied the nanostructure of a prototypical solvate ionic liquid [Li(G3)][TFSI] at electrified interfaces using MD simulations and have calculated the frontier orbital energies and Li+-G3 interactions as a function of the structural distortions for [Li(G3)]+ complex cation using DFT and NBO calculations. Near positive electrodes, alternating layers of Li+ and TFSI- ions are observed, similar to that in conventional ILs. Free G3 molecules not coordinating to any Li+ ions were already observed near electrodes with 𝜎 = 0.12 C/m2, which is aggravated with 𝜎 = 0.24 C/m2. This suggests that the electrolytes near electrodes with moderate or high positive charge may no longer behave as SILs, but such a picture must be scrutinized by further studies. Near negative electrodes, however, the electrolyte largely maintains the characteristics of SILs. For electrodes with 𝜎 = -0.12 C/m2, all contact adsorbed Li+ ions are solvated, with a coordination number slightly higher than bulk Li+ ions; even near electrodes with 𝜎 = -0.24 C/m2, 82.3% of the interfacial Li+ ions are solvated by G3 molecules, with a coordination number slightly lower than that in the bulk. The small change of the coordination number of a Li+ ion as it moves from bulk to negative electrodes is attributed to the low (1:1) Li+-G3 ratio in bulk SIL (thus the coordination number is low in bulk) and the fact that G3 molecules align their geometrical plane with the electrode surface to maximize their solvation of the interfacial Li+ ions. The strong solvation of the interfacial Li+ ions affects the nanostructures of electrolytes near the electrodes, e.g., bulky [Li(G3)]+ complex cations adsorbed

24


Page 24 of 32

Page 25 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


on negative electrodes develop notable in-plane ordering when the surface charge density increases to -0.24 C/m2. While the coordination number of the interfacial Li+ ions is similar to that of bulk Li+ ions, their solvation structure does show some differences. As the electrode becomes more negatively charged, the Li+ ion is gradually pulled away from the geometrical plane formed by the alkoxy oxygen atoms of G3 molecule. DFT molecular orbital calculations show that the HOMO energy is lifted while the LUMO energy is lowered as the [Li(G3)]+ complex cation undergoes structural distortions from the ideal coplanar conformation. The decrease of LUMO energy implies the reduction of Li+ ions near a negative electrode in a Li-ion battery is facilitated by the structural distortion during the charge process, which is verified by the localization of the LUMO orbital on the Li+ ion that is pulled out of the G3 molecular plane. NBO calculations further show that the interactions between Li+ ion and G3 ligand are significantly weakened by the structural distortions. This effect may also facilitate the reduction of Li+ at the negative electrode during the charge process. Overall, our MD simulations have revealed some unique features for the nanostructures of the [Li(G3)][TFSI] including in-plane ordering, solvation, and structural distortion of [Li(G3)]+ complex cations at electrified interfaces compared with bulk SILs or classic interfacial characteristics of conventional ILs. DFT molecular orbital and NBO calculations have uncovered new results for the effects of structural distortions on the reduction of the [Li(G3)]+ complex cation near the negative electrode of a Li-ion battery. These observations and understandings are helpful in clarifying the changes of nanostructures and properties of SILs near electrified interfaces during actual operations of Li-ion batteries and in providing guidance toward the rational design of new SILs electrolytes.

25



Supporting Information: The Supporting Information is available free of charge on the ACS Publications website. Number density profiles of O1 and O2 atoms in G3 molecules, radial distribution function of the oxygen atoms of G3 molecules with respect to Li+ ions in bulk [Li(G3)][TFSI], orientation of the plane formed by the oxygen atoms of interfacial G3 molecules, visualizations of HOMO and LUMO orbitals for the [Li(G3)]+ complex cation with a Li+G3 distance of 0 and 1.5 Å, visualizations of natural bond orbitals for the [Li(G3)]+ complex cation with a Li+-G3 distance of 0 Å, the coordinates for the [Li(G3)]+ complex cation in a range of Z=0 to 3.0 Å with a step size of 0.25 Å (PDF)

Acknowledgments: We thank the ARC at Virginia Tech for allocations of computer time. R.Q. was partially supported by an appointment to the HERE program for faculty at the Oak Ridge National Laboratory (ORNL) administered by Oak Ridge Institute of Science and Education. J.H. and B.G.S. acknowledge work performed at the Center for Nanophase Materials Sciences, which is a DOE Office of Science User Facility, and the computing resources provided by the National Energy Research Scientific Computing Center. VMD was developed by the Theoretical and Computational Biophysics Group in the Beckman Institute for Advanced Science and Technology at the University of Illinois at Urbana-Champaign.

26


Page 26 of 32

Page 27 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


References (1) Ohno, H., Electrochemical Aspects of Ionic Liquids. John Wiley & Sons: 2005; Chapter 1, pp 1-4; Chapter 17, pp 213. (2) Koel, M., Ionic Liquids in Chemical Analysis. Crit. Rev. Anal. Chem. 2005, 35, 177-192. (3) Fedorov, M. V.; Kornyshev, A. A., Ionic Liquids at Electrified Interfaces. Chem. Rev. 2014, 114, 2978-3036. (4) Plechkova, N. V.; Seddon, K. R., Applications of Ionic Liquids in the Chemical Industry. Chem. Soc. Rev. 2008, 37, 123-150. (5) Greaves, T. L.; Drummond, C. J., Protic Ionic Liquids: Properties and Applications. Chem. Rev. 2008, 108, 206-237. (6) Angell, C. A.; Ansari, Y.; Zhao, Z., Ionic Liquids: Past, Present and Future. Faraday Discuss. 2012, 154, 9-27. (7) Mandai, T.; Yoshida, K.; Ueno, K.; Dokko, K.; Watanabe, M., Criteria for Solvate Ionic Liquids. Phys. Chem. Chem. Phys. 2014, 16, 8761-8772. (8) Yoshida, K.; Tsuchiya, M.; Tachikawa, N.; Dokko, K.; Watanabe, M., Change from Glyme Solutions to Quasi-Ionic Liquids for Binary Mixtures Consisting of Lithium Bis (Trifluoromethanesulfonyl) Amide and Glymes. J. Phys. Chem. C 2011, 115, 18384-18394. (9) Yoshida, K.; Nakamura, M.; Kazue, Y.; Tachikawa, N.; Tsuzuki, S.; Seki, S.; Dokko, K.; Watanabe, M., Oxidative-Stability Enhancement and Charge Transport Mechanism in Glyme–Lithium Salt Equimolar Complexes. J. Am. Chem. Soc. 2011, 133, 13121-13129. (10) Ueno, K.; Park, J.-W.; Yamazaki, A.; Mandai, T.; Tachikawa, N.; Dokko, K.; Watanabe, M., Anionic Effects on Solvate Ionic Liquid Electrolytes in Rechargeable Lithium–Sulfur Batteries. J. Phys. Chem. C 2013, 117, 20509-20516. (11) Mandai, T.; Yoshida, K.; Tsuzuki, S.; Nozawa, R.; Masu, H.; Ueno, K.; Dokko, K.; Watanabe, M., Effect of Ionic Size on Solvate Stability of Glyme-Based Solvate Ionic Liquids. J. Phys. Chem. B 2015, 119, 1523-1534. (12) Zhang, C.; Ueno, K.; Yamazaki, A.; Yoshida, K.; Moon, H.; Mandai, T.; Umebayashi, Y.; Dokko, K.; Watanabe, M., Chelate Effects in Glyme/Lithium Bis (Trifluoromethanesulfonyl) Amide Solvate Ionic Liquids. I. Stability of Solvate Cations and Correlation with Electrolyte Properties. J. Phys. Chem. B 2014, 118, 5144-5153. (13) Zhang, C.; Yamazaki, A.; Murai, J.; Park, J.-W.; Mandai, T.; Ueno, K.; Dokko, K.; Watanabe, M., Chelate Effects in Glyme/Lithium Bis (Trifluoromethanesulfonyl) Amide Solvate Ionic Liquids, Part 2: Importance of Solvate-Structure Stability for Electrolytes of Lithium Batteries. J. Phys. Chem. C 2014, 118, 17362-17373.

27



(14) Dokko, K.; Tachikawa, N.; Yamauchi, K.; Tsuchiya, M.; Yamazaki, A.; Takashima, E.; Park, J.-W.; Ueno, K.; Seki, S.; Serizawa, N., Solvate Ionic Liquid Electrolyte for Li–S Batteries. J. Electrochem. Soc. 2013, 160, A1304-A1310. (15) Tamura, T.; Yoshida, K.; Hachida, T.; Tsuchiya, M.; Nakamura, M.; Kazue, Y.; Tachikawa, N.; Dokko, K.; Watanabe, M., Physicochemical Properties of Glyme-Li Salt Complexes as a New Family of Room-Temperature Ionic Liquids. Chem. Lett. 2010, 39, 753-755. (16) Shimizu, K.; Freitas, A. A.; Atkin, R.; Warr, G. G.; FitzGerald, P. A.; Doi, H.; Saito, S.; Ueno, K.; Umebayashi, Y.; Watanabe, M., Structural and Aggregate Analyses of (Li Salt+ Glyme) Mixtures: The Complex Nature of Solvate Ionic Liquids. Phys. Chem. Chem. Phys. 2015, 17, 22321-22335. (17) Tsuzuki, S.; Shinoda, W.; Matsugami, M.; Umebayashi, Y.; Ueno, K.; Mandai, T.; Seki, S.; Dokko, K.; Watanabe, M., Structures of [Li (Glyme)]+ Complexes and Their Interactions with Anions in Equimolar Mixtures of Glymes and Li [Tfsa]: Analysis by Molecular Dynamics Simulations. Phys. Chem. Chem. Phys. 2015, 17, 126-129. (18) Saito, S.; Watanabe, H.; Hayashi, Y.; Matsugami, M.; Tsuzuki, S.; Seki, S.; Canongia Lopes, J. N.; Atkin, R.; Ueno, K.; Dokko, K., Li+ Local Structure in Li–Tetraglyme Solvate Ionic Liquid Revealed by Neutron Total Scattering Experiments with the 6/7li Isotopic Substitution Technique. J. Phys. Chem. Lett. 2016, 7, 2832-2837. (19) Jorgensen, W. L.; Maxwell, D. S.; Tirado-Rives, J., Development and Testing of the Opls All-Atom Force Field on Conformational Energetics and Properties of Organic Liquids. J. Am. Chem. Soc. 1996, 118, 11225-11236. (20) Ueno, K.; Tatara, R.; Tsuzuki, S.; Saito, S.; Doi, H.; Yoshida, K.; Mandai, T.; Matsugami, M.; Umebayashi, Y.; Dokko, K., Li+ Solvation in Glyme–Li Salt Solvate Ionic Liquids. Phys. Chem. Chem. Phys. 2015, 17, 8248-8257. (21) Feng, G.; Huang, J.; Sumpter, B. G.; Meunier, V.; Qiao, R., Structure and Dynamics of Electrical Double Layers in Organic Electrolytes. Phys. Chem. Chem. Phys. 2010, 12, 5468-5479. (22) Haskins, J. B.; Wu, J. J.; Lawson, J. W., Computational and Experimental Study of Li-Doped Ionic Liquids at Electrified Interfaces. J. Phys. Chem. C 2016, 120, 11993-12011. (23) Ivaništšev, V.; Méndez-Morales, T.; Lynden-Bell, R. M.; Cabeza, O.; Gallego, L. J.; Varela, L. M.; Fedorov, M. V., Molecular Origin of High Free Energy Barriers for Alkali Metal Ion Transfer through Ionic Liquid–Graphene Electrode Interfaces. Phys. Chem. Chem. Phys. 2016, 18, 1302-1310. (24) Ebadi, M.; Costa, L. T.; Araujo, C. M.; Brandell, D., Modelling the Polymer Electrolyte/Li-Metal Interface by Molecular Dynamics Simulations. Electrochim. Acta 2017, 234, 43-51.

28


Page 28 of 32

Page 29 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(25) Xing, L.; Vatamanu, J.; Borodin, O.; Smith, G. D.; Bedrov, D., Electrode/Electrolyte Interface in Sulfolane-Based Electrolytes for Li Ion Batteries: A Molecular Dynamics Simulation Study. J. Phys. Chem. C 2012, 116, 23871-23881. (26) Ivaništšev, V.; Fedorov, M.; Lynden-Bell, R., Screening of Ion–Graphene Electrode Interactions by Ionic Liquids: The Effects of Liquid Structure. J. Phys. Chem. C 2014, 118, 5841-5847. (27) McLean, B.; Li, H.; Stefanovic, R.; Wood, R. J.; Webber, G. B.; Ueno, K.; Watanabe, M.; Warr, G. G.; Page, A.; Atkin, R., Nanostructure of [Li (G4)] Tfsi and [Li (G4)] No 3 Solvate Ionic Liquids at Hopg and Au (111) Electrode Interfaces as a Function of Potential. Phys. Chem. Chem. Phys. 2015, 17, 325-333. (28) Coles, S. W.; Mishin, M.; Perkin, S.; Fedorov, M. V.; Ivaništšev, V. B., The Nanostructure of a Lithium Glyme Solvate Ionic Liquid at Electrified Interfaces. Phys. Chem. Chem. Phys. 2017, 19, 11004-11010. (29) Humphrey, W.; Dalke, A.; Schulten, K., "VMD: Visual Molecular Dynamics". J. Mol. Graph. 1996, 14, 33-38; software can be found at http://www.ks.uiuc.edu/Research/vmd/ (accessed August 22, 2018). (30) Cornell, W. D.; Cieplak, P.; Bayly, C. I.; Gould, I. R.; Merz, K. M.; Ferguson, D. M.; Spellmeyer, D. C.; Fox, T.; Caldwell, J. W.; Kollman, P. A., A Second Generation Force Field for the Simulation of Proteins, Nucleic Acids, and Organic Molecules. J. Am. Chem. Soc. 1995, 117, 5179-5197. (31) Hess, B.; Kutzner, C.; Van Der Spoel, D.; Lindahl, E., Gromacs 4: Algorithms for Highly Efficient, Load-Balanced, and Scalable Molecular Simulation. J. Chem. Theory Comput. 2008, 4, 435-447. (32) Ueno, K.; Yoshida, K.; Tsuchiya, M.; Tachikawa, N.; Dokko, K.; Watanabe, M., Glyme–Lithium Salt Equimolar Molten Mixtures: Concentrated Solutions or Solvate Ionic Liquids? J. Phys. Chem. B 2012, 116, 11323-11331. (33) Bussi, G.; Donadio, D.; Parrinello, M., Canonical Sampling Through Velocity Rescaling. J. Chem. Phys. 2007, 126, 014101. (34) Hess, B.; Bekker, H.; Berendsen, H. J.; Fraaije, J. G., Lincs: A Linear Constraint Solver for Molecular Simulations. J. Comput. Chem. 1997, 18, 1463-1472. (35) Yeh, I.-C.; Berkowitz, M. L., Ewald Summation for Systems with Slab Geometry. J. Chem. Phys. 1999, 111, 3155-3162. (36) Becke, A., Density-Functional Thermochemistry. Iii. The Role of Exact Exchange. J. Chem. Phys. 1993, 98, 5648-5652. (37) Lee, C.; Yang, W.; Parr, R. G., Development of the Colle-Salvetti Correlation-Energy Formula into a Functional of the Electron Density. Phys. Rev. B 1988, 37, 785.

29



(38) McLean, A.; Chandler, G., Contracted Gaussian Basis Sets for Molecular Calculations. I. Second Row Atoms, Z= 11–18. J. Chem. Phys. 1980, 72, 5639-5648. (39) Frisch, M.; Trucks, G.; Schlegel, H. B.; Scuseria, G.; Robb, M.; Cheeseman, J.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G., Gaussian 09, Revision D.01, Gaussian. Inc., Wallingford, CT, 2009. (40) Glendening, E.; Reed, A.; Carpenter, J.; Weinhold, F., Nbo, Version 3.1. 1998. (41) Fedorov, M. V.; Kornyshev, A. A., Towards Understanding the Structure and Capacitance of Electrical Double Layer in Ionic Liquids. Electrochim. Acta 2008, 53, 6835-6840. (42) Feng, G.; Zhang, J.; Qiao, R., Microstructure and Capacitance of the Electrical Double Layers at the Interface of Ionic Liquids and Planar Electrodes. J. Phys. Chem. C 2009, 113, 4549-4559. (43) Kornyshev, A. A.; Qiao, R., Three-Dimensional Double Layers. J. Phys. Chem. C 2014, 118, 18285– 18290. (44) Gómez-González, V.; Docampo-Álvarez, B.; Méndez-Morales, T.; Cabeza, O.; Ivaništšev, V. B.; Fedorov, M. V.; Gallego, L. J.; Varela, L. M., Molecular Dynamics Simulation of the Structure and Interfacial Free Energy Barriers of Mixtures of Ionic Liquids and Divalent Salts near a Graphene Wall. Phys. Chem. Chem. Phys. 2017, 19, 846-853. (45) Borodin, O.; Ren, X.; Vatamanu, J.; von Wald Cresce, A.; Knap, J.; Xu, K., Modeling Insight into Battery Electrolyte Electrochemical Stability and Interfacial Structure. Acc. Chem. Res. 2017, 50, 2886-2894. (46) Rotenberg, B.; Salanne, M., Structural Transitions at Ionic Liquid Interfaces. J. Phys. Chem. Lett. 2015, 6, 4978-4985. (47) Merlet, C.; Salanne, M.; Rotenberg, B.; Madden, P. A., Imidazolium Ionic Liquid Interfaces with Vapor and Graphite: Interfacial Tension and Capacitance from Coarse-Grained Molecular Simulations. J. Phys. Chem. C 2011, 115, 16613-16618. (48) Vatamanu, J.; Borodin, O.; Smith, G. D., Molecular Simulations of the Electric Double Layer Structure, Differential Capacitance, and Charging Kinetics for N-Methyl-N-Propylpyrrolidinium Bis (Fluorosulfonyl) Imide at Graphite Electrodes. J. Phys. Chem. B 2011, 115, 3073-3084. (49) Kislenko, S. A.; Samoylov, I. S.; Amirov, R. H., Molecular Dynamics Simulation of the Electrochemical Interface between a Graphite Surface and the Ionic Liquid [Bmim][Pf 6]. Phys. Chem. Chem. Phys. 2009, 11, 5584-5590. (50) Pinilla, C.; Del Pópolo, M.; Kohanoff, J.; Lynden-Bell, R., Polarization Relaxation in an Ionic Liquid Confined between Electrified Walls. J. Phys. Chem. B 2007, 111, 4877-4884. (51) Bazant, M. Z.; Storey, B. D.; Kornyshev, A. A., Double Layer in Ionic Liquids: Overscreening Versus Crowding. Phys. Rev. Lett. 2011, 106, 046102.

30


Page 30 of 32

Page 31 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(52) Kornyshev, A. A., Double-Layer in Ionic Liquids: Paradigm Change? J. Phys. Chem. B 2007, 111, 5545-5557. (53) Ivaništšev, V.; O’Connor, S.; Fedorov, M., Poly (a) Morphic Portrait of the Electrical Double Layer in Ionic Liquids. Electrochem. Commun. 2014, 48, 61-64. (54) Kislenko, S.; Moroz, Y. O.; Karu, K.; Ivaništšev, V.; Fedorov, M., Calculating the Maximum Density of the Surface Packing of Ions in Ionic Liquids. Russ. J. Phys. Chem. A 2018, 92, 999-1005. (55) Coles, S. W.; Smith, A. M.; Fedorov, M. V.; Hausen, F.; Perkin, S., Interfacial Structure and Structural Forces in Mixtures of Ionic Liquid with a Polar Solvent. Faraday Discuss. 2018, 206, 427-442. (56) Merlet, C.; Limmer, D. T.; Salanne, M.; Van Roij, R.; Madden, P. A.; Chandler, D.; Rotenberg, B., The Electric Double Layer Has a Life of Its Own. J. Phys. Chem. C 2014, 118, 18291-18298. (57) Pounds, M.; Tazi, S.; Salanne, M.; Madden, P. A., Ion Adsorption at a Metallic Electrode: An Ab Initio Based Simulation Study. J. Phys.: Condens. Matter 2009, 21, 424109. (58) Dean, D. S.; Dobnikar, J.; Naji, A.; Podgornik, R., Electrostatics of Soft and Disordered Matter. CRC Press: Boca Raton, FL, 2014, Chapter 8, pp 93-104. (59) Lewis, G. N., Acids and Bases. J Franklin Inst. 1938, 226, 293-313. (60) Pearson, R. G., Hard and Soft Acids and Bases. J. Am. Chem. Soc. 1963, 85, 3533-3539. (61) Laing, M., No Rabbit Ears on Water. The Structure of the Water Molecule: What Should We Tell the Students? J. Chem. Educ. 1987, 64, 124. (62) Clauss, A. D.; Nelsen, S. F.; Ayoub, M.; Moore, J. W.; Landis, C. R.; Weinhold, F., Rabbit-Ears Hybrids, Vsepr Sterics, and Other Orbital Anachronisms. Chem. Educ. Res. Prac. 2014, 15, 417-434.

31



TOC Image

32


Page 32 of 32

Solvate Ionic Liquids at Electrified Interfaces - ACS Publications

Recommend Documents