Solvation Structure of Sodium Bis(fluorosulfonyl)imide-Glyme Solvate

Article Cite This: J. Phys. Chem. B XXXX, XXX, XXX−XXX

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Solvation Structure of Sodium Bis(fluorosulfonyl)imide-Glyme Solvate Ionic Liquids and Its Influence on Cycling of Na-MNC Cathodes Pieter Geysens,† Vijay Shankar Rangasamy,‡ Savitha Thayumanasundaram,‡ Koen Robeyns,∥ Luc Van Meervelt,† Jean-Pierre Locquet,‡ Jan Fransaer,§ and Koen Binnemans*,† †

Department of Chemistry, KU Leuven, Celestijnenlaan 200F, B-3001 Leuven, Belgium Department of Physics and Astronomy, KU Leuven, Celestijnenlaan 200D, B-3001 Leuven, Belgium § Department of Materials Engineering, KU Leuven, Kasteelpark Arenberg 44, B-3001 Leuven, Belgium ∥ Institute of Condensed Matter and Nanosciences, UC Louvain, Place Louis Pasteur 1, B-1348 Louvain-la-Neuve, Belgium ‡

S Supporting Information *

ABSTRACT: Electrolytes consisting of sodium bis(fluorosulfonyl)imide (NaFSI) dissolved in glymes (monoglyme, diglyme, and triglyme) were characterized by FT-Raman spectroscopy and 13C, 17O, and 23Na NMR spectroscopy. The glyme:NaFSI molar ratio was varied from 50:1 to 1:1, and it was observed that, in the dilute electrolytes, the sodium salt is completely dissociated into solvent separated ion pairs (SSIPs). However, contact ion pairs (CIPs) and aggregates (AGGs) become the predominant species in more concentrated solutions. Some of the electrolytes with the highest concentrations can be classified as solvate ionic liquids (SILs), where all of the solvent molecules are coordinated to sodium cations. Therefore, these electrolytes are fundamentally different from more dilute electrolytes which are typically used in commercially available secondary batteries. The melting point or glass transition temperature, dynamic viscosity, density, sodium concentration, and ionic conductivity of these solvate ionic liquids are reported as well as the crystal structures of [Na(G3)][FSI] and [Na(G3)2][FSI]. Galvanostatic cycling experiments were performed in coin-type cells with a Na2/3[Mn0.55Ni0.30Co0.15]O2 cathode to study the influence of these electrolytes on the electrochemical stability and charge/discharge behavior.

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INTRODUCTION

lytes consisting of various sodium salts dissolved in carbonate and ether solvents have been reported by Palaciń et al.10,11 They found that a 1 M solution of sodium perchlorate in a mixture of ethylene carbonate (EC), propylene carbonate (PC), and dimethyl carbonate (DMC) in the ratio EC0.45:PC0.45:DMC0.1 was the optimum electrolyte for hard carbon electrodes resulting in good rate capability and high capacity upon sustained cycling. These electrolytes are very similar in concentration to the electrolytes used for commercial Li-ion batteries. However, a different type of electrolyte is attracting more and more attention, namely, the highly concentrated electrolytes. One of the first reports describing the use of this kind of electrolyte in Li-ion batteries dates back to 1985, when McKinnon et al. found that increasing the concentration of LiAsF6 inhibits cointercalation of the propylene carbonate solvent into LixZrS2 and other layered electrode materials, thereby preventing their

Due to their high specific energy density and long lifetime, lithium-ion batteries (LIBs) have been the primary power source for small portable electronic devices ever since their commercialization in the early 1990s. More recent research has shown that they can be important tools in the transition to a more sustainable energy policy, for instance, as power sources in electric and hybrid vehicles or as load-leveling devices for renewable energy sources and the power grid.1,2 However, there are concerns that lithium-ion technology is not sustainable enough for such large-scale applications where a low production cost and high sustainability are essential.3,4 In this regard, sodium-ion batteries (NIBs) have become a major contender for large-scale electrochemical energy storage.5,6 The low cost and virtually unlimited availability of sodium salts compared to lithium salts have rekindled research interest in rechargeable sodium-ion cells with intercalation electrodes after they were first conceptualized more than 20 years ago.7 Recent publications have focused mainly on the development of new sodium-ion intercalation compounds to be used as electrodes.5,8,9 Optimization studies of battery electro© XXXX American Chemical Society

Received: October 13, 2017 Revised: November 16, 2017 Published: December 4, 2017 A

DOI: 10.1021/acs.jpcb.7b10158 J. Phys. Chem. B XXXX, XXX, XXX−XXX

Article

The Journal of Physical Chemistry B exfoliation.12 More recently, highly concentrated electrolytes have resurfaced as one of the most promising candidates for application in next-generation fast-charging Li-ion batteries with a cell potential of 5 V and even in rechargeable lithium metal batteries.13,14 Although there is a penalty in conductivity due to the increase in viscosity and the formation of contact ion pairs, these electrolytes have many advantages compared to their more dilute counterparts, such as an enhanced reductive and oxidative stability, lower volatility, enhanced thermal stability, higher carrier density, and faster and more reversible electrode reactions.13−20 Many of these properties are the consequence of their unique coordination structure: as the salt concentration increases, the amount of free solvent molecules decreases until eventually all of the solvent molecules are coordinated to metal cations, which significantly increases their (electro)chemical stability.15,21 Another interesting property of highly concentrated electrolytes, in particular those consisting of equimolar mixtures of a lithium salt and oligo(ethylene glycol) dimethyl ethers (glymes), is that they can behave as ionic liquids, where the cation is a strong complex between a metal cation and a number of neutral ligands.22,23 These so-called solvate ionic liquids (SILs) combine a high metal concentration with the advantages of ionic liquids such as a low volatility, a high (electro)chemical stability, and a high ionic conductivity. The physicochemical properties (conductivity, viscosity, melting point, solvate structure) of lithium-glyme solvate ionic liquids and their use in lithium deposition and cycling experiments on various intercalation electrodes have been extensively reported in the literature.24−32 Solvate ionic liquids and highly concentrated electrolytes have found their way into the realm of sodium-ion batteries. 3 3 − 3 5 Electrolytes based on sodium bis(fluorosulfonyl)imide (NaFSI) are particularly promising because they still have an acceptable conductivity and viscosity at highly concentrated conditions, and because they show highly reversible electrode reactions.36 The high reversibility of the electrode reactions can be attributed to the fact that the FSI− anions form good solid electrolyte interphases (SEI). Recent reports show that electrolytes consisting of NaFSI dissolved in 1,2-dimethoxyethane (DME, monoglyme, G1) at a G1:NaFSI molar ratio of 2:1 or lower enable stable cycling of sodium vanadium phosphate (NVP) cathodes, hard carbon (HC) anodes, and even sodium metal anodes for up to 300 cycles with extremely high Coulombic efficiency.37,38 Although these NaFSI-glyme solutions are promising sodium battery electrolytes, there is to our knowledge only one report on their coordination structure.39 Some FTIR spectra of NaFSI-G1 electrolytes are presented, but the solvation structures are not discussed in detail. The solvation structures of various sodium salt-G4 and -G5 electrolytes were reported by Mandai et al.33 In this paper, we report on the different coordination structures, such as solvent separated ion pairs (SSIPs), contact ion pairs (CIPs), and aggregates (AGGs), that exist in equilibrium in mixtures of NaFSI and 1,2-dimethoxyethane (monoglyme, G1), 1-methoxy-2-(2-methoxyethoxy)ethane (diglyme, G2), and 1,2-bis(2-methoxyethoxy)ethane (triglyme, G3) (Figure 1). While Raman spectroscopy has typically been the technique of choice for such speciation studies in lithium electrolytes,15,18 we show that 13C, 17O, and 23 Na NMR spectroscopy are also reliable and versatile techniques that can be used to gather information on competitive concentration-dependent interactions of solvent

Figure 1. Structural formulas of the anion and ligands used in this work: (a) bis(fluorosulfonyl)imide (FSI−), (b) monoglyme (G1), (c) diglyme (G2), and (d) triglyme (G3).

molecules and anions with the sodium cation. The dynamic viscosity, density, sodium concentration, and melting point or glass transition temperature of the solvate ionic liquids [Na(L)n][FSI] (L = G1, G2, or G3; n = 1, 2, or 3) are also reported. As a proof of concept, we report the charge− discharge cycling results of the monoglyme-based SILs used in coin-type batteries to demonstrate their application in a conventional sodium-ion battery configuration.

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EXPERIMENTAL METHODS Materials and Synthesis. NaFSI (99.7%) was purchased from Solvionic (Toulouse, France) and was used as received. 1,2-Dimethoxyethane (G1, 99.5%, anhydrous) and 1-methoxy2-(2-methoxyethoxy)ethane (G2, 99.5%, anhydrous) were purchased from Sigma-Aldrich (Diegem, Belgium) and were used as received. 1,2-Bis(2-methoxyethoxy)ethane (G3, 99%) was purchased from Sigma-Aldrich (Diegem, Belgium) and dried prior to use on molecular sieves (4 Å, 8−12 mesh), purchased from Acros Organics (Geel, Belgium). The water content of the glymes was measured by a Mettler-Toledo C30S coulometric Karl Fischer titrator and was found to be lower than 50 ppm, as claimed by the manufacturer. Storage of the reagents and all of the manipulations involving contact of the chemicals with the atmosphere were performed in an argonfilled glovebox with an oxygen and water concentration less than 1 ppm. The solutions were prepared by adding a stoichiometric amount of glyme to NaFSI in glass vials and stirring the mixture at room temperature until the salt was completely dissolved. The more concentrated solutions (3:1− 1:1 G1:NaFSI molar ratio) required stirring at 60 °C to accelerate complete dissolution of the salt. Physicochemical Properties. Melting points and glass transition temperatures were determined on a Mettler-Toledo DSC-1 instrument at a heating rate of 10 °C min−1 under a helium atmosphere. Aluminum crucibles were filled with samples of electrolyte (3−12 mg) inside an argon-filled glovebox and sealed to prevent contact with the air. The samples were cycled twice between −70 and 50 °C (90 °C for [Na(G3)2][FSI]). Liquid samples (G1 and G2 complexes) were first cooled down and subsequently heated. For solid samples (G3 complexes), the order was reversed. The samples for which a melting point could not be observed were cycled between −100 and 50 °C in order to determine the glass transition temperature. The values of Tm and Tg observed in the B


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The Journal of Physical Chemistry B

= 0 ppm, δ(23Na) = 0 ppm) for 13C, 17O, and 23Na measurements, respectively. In all of the spectra that are presented below, a negative shift value corresponds to an upfield (shielding) shift vs the reference sample. Calibration of the sample spectra to the external references and peak analysis was performed with the Bruker TopSpin software. Raman Measurements. FT-Raman spectra were recorded between 3500 and 50 cm−1 on a Bruker Vertex 70 spectrometer with a RamII Raman module and a germanium diode detector. Each measurement consisted of 64 scans at a resolution of 2 cm−1 with a 1064.38 nm laser (Nd:YAG) at a power of 250 mW. The samples were the same as those used for the NMR measurements. Deconvolution of the spectra was done by the OriginPro 8 software. Electrochemical Measurements. All of the electrochemical measurements and assembly of the coin cells were performed inside an argon-filled glovebox with water and oxygen concentrations below 1 ppm. Electrochemical windows were determined by linear sweep voltammetry (LSV), using an Autolab PGSTAT302N potentiostat and GPES software. The working electrode was a polished platinum disk electrode (ϕ = 0.5 mm), the reference and counter electrodes were sodium metal, and the scan rate was 10 mV s−1 for all voltammograms. The ionic conductivity of the electrolytes was measured at ambient glovebox temperature (28 °C) by electrochemical impedance spectroscopy (EIS) in a frequency range of 10 Hz−1 MHz, using the same potentiostat and NOVA 1.11 software. A custom-built cell was used with two parallel copper disk electrodes (ϕ = 10 mm) facing each other at a distance of 21 mm. For galvanostatic tests of the monoglyme electrolytes as a case study, 2032-type cells with sodium metal as the reference and counter electrodes were assembled inside an argon-filled glovebox. Pristine Na2/3[Mn0.55Ni0.30Co0.15]O2 was used as the positive electrode, of which the synthesis and processing have been reported elsewhere.45 In summary, mixtures of 80 wt % of the active material and 10 wt % each of polyvinylidene difluoride (PVDF) and acetylene black were used. The active material loading in each cathode disc was about 6 mg. The electrodes were separated from each other by two sheets of Whatman glass microfiber (Grade GF/D). About 50 μL of electrolyte was used in each cell, and all of the electrodes studied herein were immersed in the electrolyte for a period of 12 h prior to cell assembly. Galvanostatic charge− discharge cycles of the assembled coin cells were performed using a Maccor 4000 battery test station in the voltage range 1.5−3.7 V.

second cycle are reported. Thermogravimetric analysis (TGA) was performed on a TA Instruments TGA Q500 at a heating rate of 5 °C min−1 under a nitrogen atmosphere. The sample size was typically between 5 and 10 mg. The viscosity and density of the electrolytes were measured on a Lovis 2000 ME rolling-ball microviscometer and a DMA 4500 M density meter, respectively. In order to avoid contact with air, the samples were syringed out from their sealed argon-filled containers and subsequently injected into the capillaries or density chamber of the device. The temperature during both the viscosity and density measurements was controlled by the internal thermostat of the device. Crystal Structures. Single crystals of [Na(G3)][FSI] were grown by slow cooling of the melt. For [Na(G3)2][FSI], crystals were obtained from a concentrated solution of the complex in triglyme. Once out of the crystallizing solution, the hygroscopic crystals started to absorb moisture from the air and dissolved within a few minutes. X-ray intensity data were collected at −123 °C ([Na(G3)][FSI]) or −73 °C ([Na(G3)2][FSI]) on a Rigaku UltraX 18S generator (Xenocs mirrors, Mo Kα radiation, λ = 0.71073 Å) using a MAR345 image plate. The images were interpreted and integrated with CrysAlisPRO,40 and the implemented absorption correction was applied. The resolution of the diffraction pattern of [Na(G3)2][FSI] was limited; a cutoff of 0.92 Å was used. Using Olex2,41 the structures were solved with the ShelXS42 structure solution program by Direct Methods and refined with the ShelXL43 refinement package using full-matrix least-squares minimization on F2. Non-hydrogen atoms were refined anisotropically and hydrogen atoms in the riding mode with isotropic temperature factors fixed at 1.2 times Ueq of the parent atoms (1.5 for methyl groups). The FSI anion in [Na(G3)2][FSI] is disordered over four positions (occupancies for positions 1 and 2: 0.311(5) and 0.189(5); positions 3 and 4 are generated by a local inversion center) and refined using restraints for distances, angles, and temperature factors. The crystal of [Na(G3)2][FSI] is a two-component twin (BASF 0.28). Twin law was found with the TwinRotMat function of PLATON,44 and a two domain reflection list (HKLF 5) was used for the final refinement. The atomic coordinates and anisotropic displacement factors for both structures have been deposited with the Cambridge Crystallographic Data Centre under the numbers CCDC 1577498 ([Na(G3)][FSI]) and 1577499 ([Na(G3)2][FSI]) and can be received from the following address: https://www.ccdc.cam.ac.uk/structures/. NMR Measurements. The 13C, 17O, and 23Na NMR spectra were recorded on a Bruker Avance II+ 600 MHz spectrometer (operating at 150.903 MHz for 13C, 81.356 MHz for 17O, and 158.746 MHz for 23Na). The samples were prepared by filling oven-dried NMR sample tubes (5 mm cross section) with the solutions inside the glovebox and subsequently capping them with polypropylene caps and sealing with laboratory film to retain a pure argon atmosphere. The external references and samples within one series with the same solvent were all measured in one continuous sequence on the same day within 5 h of each other to minimize the natural drift effects and possible fluctuations of the magnetic field. The drift of the magnetic field was estimated by measuring the reference samples at several time intervals during the sample sequence and was found to be insignificantly small compared to the observed shifts in the samples. The chemical shifts are noted in parts per million (ppm) referenced to DMSO (δ(13C) = 40.45 ppm) or a saturated solution of NaCl in D2O (δ(17O)

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RESULTS AND DISCUSSION Synthesis and Physicochemical Properties. Complexes with the general formula [Na(L)n][FSI] (with L = G1, G2, or G3, n = 1, 2, or 3) were prepared by stoichiometric addition of the ligands to NaFSI and subsequent stirring of the mixture at 60 °C until a clear colorless homogeneous liquid was obtained. The detailed synthesis procedure and characterization data of each complex can be found in the Supporting Information. The Watanabe group already reported solvate ionic liquids based on several sodium salts, including NaFSI, and glymes, where they focused primarily on 1:1 complexes with long glyme ligands such as tetraglyme (G4) and pentaglyme (G5).33 Due to the chelate effect, these polydentate ligands result in the formation of complexes with a very high thermal stability and a high ionicity. They also summarized possible criteria to classify compounds as solvate ionic liquids:23 (1) they must be solvate C


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at higher temperatures, the NaFSI-G1 solvate surprisingly becomes the most thermally stable. G1 is lost from 25 °C to approximately 180 °C and a distinct plateau is observed, whereas the complexes with G2 and G3 lost already more of their mass at that temperature. This trend is also reflected in the isothermal TGA curves, which were recorded at 100 °C. For [Na(G1)3][FSI], a fast loss of mass is observed until approximately 57 wt % of the ligands has evaporated after which a thermally stable complex is formed. For the other two complexes containing G2 or G3, the mass of the sample decreases continuously for the entire duration of the experiment without reaching such a stable composition. The dynamic viscosity (ηdyn) and density (ρ) of the NaFSIglyme complexes were determined at 25 °C (Table 2). From

compounds between an ion and a ligand(s) in a certain stoichiometric ratio; (2) they must consist entirely of complex ions (solvates) and their counterions in the molten state; (3) they must show no physicochemical properties based on both pure ligands and precursor salts under using conditions; (4) they must have a melting point below 100 °C to satisfy the criterion for typical ILs; and (5) they must have a negligible vapor pressure under typical application conditions. If these criteria are only partially met, the compounds are classified as “poor solvate ionic liquids”. In order to investigate whether the [Na(L)n][FSI] (L = G1, G2, or G3; n = 1, 2, or 3) complexes can indeed be classified as solvate ionic liquids, they were analyzed with several techniques to see if the above-mentioned criteria are met. Criteria 1 and 2 will be discussed in detail in the NMR and Raman sections of this paper. The melting point was determined by differential scanning calorimetry (DSC), and it was found to be below 100 °C or it was not observed, thereby satisfying criterion 4 (Table 1).

Table 2. Dynamic Viscosity, Density, and Sodium Ion Concentration of Complexes [Na(G1)n][FSI], [Na(G2)n][FSI], and [Na(G3)][FSI], Measured at 25 °C

Table 1. Melting Points and Glass Transition Temperatures of Complexes [Na(G1)n][FSI], [Na(G2)n][FSI], and [Na(G3)n][FSI]

a

complex

Tma (°C)

Tga (°C)

[Na(G1)][FSI] [Na(G1)2][FSI] [Na(G1)3][FSI] [Na(G2)][FSI] [Na(G2)2][FSI] [Na(G3)][FSI] [Na(G3)2][FSI]

n.o. n.o. −21 n.o. 17 38 67

−75 −72 n.o. −81 n.o. n.o. n.o.

complex

ηdyn (10−3 Pa s)

ρ (kg m−3)

Na+ concentration (mol dm−3)

[Na(G1)][FSI] [Na(G1)2][FSI] [Na(G1)3][FSI] [Na(G2)][FSI] [Na(G2)2][FSI] [Na(G3)][FSI]

750 45 11 311 27 497

1510 1300 1200 1440 1260 1380

5.15 3.40 2.54 4.27 2.67 3.62

the obtained density values, the sodium ion concentration was calculated. Even though [Na(G3)][FSI] is a solid at room temperature, once it is in the molten state, it remains stable as an undercooled liquid which allowed the viscosity to be determined. For [Na(G3)2][FSI], this was not the case. Except for the solvate ionic liquids with a 1:1 glyme:NaFSI ratio, the dynamic viscosity of all of the complexes is reasonably low. This in turn translates into a specific ionic conductivity that is sufficiently high to be useful for electrochemical experiments, as was determined by us further down this paper and as was observed by Schafzahl et al.37 The Na+ concentration of these liquid complexes is very high compared to the Li + concentration of a typical commercial Li-ion battery electrolyte (1 M). This is a very advantageous property for achieving a high current density during electrochemical reactions such as electrodeposition. Therefore, these electrolytes should theoretically enable cell charging with a good rate capability and at a high power density. Solvation Study of Glyme-NaFSI Binary Mixtures. When NaFSI is dissolved in a solvent, the sodium cations interact with the solvent molecules and FSI− anions, resulting in a number of possible solvate structures. These solvate structures are typically classified according to the number of anions that are directly coordinated to the cation. The following three types are defined (Figure 2): (1) Solvent separated ion pairs (SSIPs): Only solvent molecules are present in the solvation shell around the cation, separating it from the counteranions. (2) Contact ion pairs (CIPs): Solvent molecules and one anion are present in the solvation shell around the cation, resulting in a neutral complex. A further distinction is made based on whether the anion acts as a monodentate (CIP I) or a bidentate (CIP II) ligand. (3) Aggregates (AGGs): Solvent molecules and multiple anions are present in the solvation shell around the

n.o.: not observed.

Except for the complexes with triglyme ligands, the measured melting points were even below room temperature (25 °C). Since the FSI− anion, similarly to bis(trifluoromethylsulfonyl)imide (Tf2N−), is a flexible and highly delocalized anion that inhibits crystallization, the melting point (Tm) is either very low or not observable. In the latter case, the glass transition temperature (Tg) upon cooling is reported. The only complexes for which a well-defined liquid-to-crystal transition was observed were [Na(G1)3][FSI], [Na(G2)2][FSI], [Na(G3)][FSI], and [Na(G3)2][FSI]. There is also clearly an increasing trend of the melting points as a function of the length of the glyme ligands: all of the complexes with monoglyme ligands are liquid below 0 °C, whereas [Na(G2)2][FSI] is liquid just below room temperature and the complexes with triglyme ligands are both solid at room temperature. The thermal stability of [Na(G1)3][FSI], [Na(G2)2][FSI], and [Na(G3)2][FSI] was evaluated by dynamic and isothermal thermogravimetric analysis (TGA) at 100 °C (Figure S1, Supporting Information). The volatility of G1, G2, and G3 is significantly reduced when they are coordinated to sodium cations. Whereas the pure glymes are all completely evaporated before the temperature reaches 120 °C, this temperature is extended to more than the decomposition temperature of the FSI− anion (180 °C) when they are part of a NaFSI-glyme solvate. However, there is still significant loss of ligands for all of the complexes so they do not meet criterion 5, proposed by Mandai et al.23 For the pure glymes, the order of volatility is G1 > G2 > G3, as expected from their respective boiling points (85, 162, and 216 °C). For the complexes, this order is also initially observed, as expected from the chelate effect. However, D


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Figure 2. Schematic representation of the three main solvated species in NaFSI-triglyme binary mixtures: solvent separated ion pairs (SSIP) (left), contact ion pairs (CIP I and CIP II) (middle), and aggregate structures (AGG I and AGG II) (right). CIP I and CIP II are contact ion pairs with respectively a mono- or bidentately coordinated anion, and AGG I and AGG II are aggregates where respectively two or three cations are linked by a single anion.

cation. In the case of long flexible polydentate anions such as FSI− or Tf2N−, this often results in the formation of large polymeric coordination networks.15 AGGs are further distinguished according to the number of cations that are coordinated by the same anion. AGG I and AGG II refer to aggregate structures where two or three cations respectively are linked by a single anion. Henderson et al. have reported crystal structures of many lithium salt-glyme solvates and found examples of all three solvate types.31,32 In solution, these different solvate structures are also present and exist in equilibrium with each other due to the kinetically labile nature of alkali metal ion complexes. The predominant solvate type in solution depends on several parameters that influence the position of this equilibrium, such as the coordination ability of the solvent and anion, the salt concentration, and the temperature. Crystal Structures. Crystal structures of the complexes [Na(G3)][FSI] and [Na(G3)2][FSI] were determined by single crystal X-ray diffraction (XRD). The crystallographic data of both complexes can be found in the Supporting Information (Table S1). In the crystal structure of [Na(G3)][FSI], the asymmetric unit consists of a Na+ cation, a G3 ligand, and a FSI− anion. The Na+ cations are each coordinated by four oxygen atoms from the G3 ligand, one oxygen atom from the FSI− anion, and two oxygen atoms from a symmetry-related FSI− anion, resulting in a total coordination number of seven (Figure 3). Each FSI− anion is coordinated to one Na+ cation as a bidentate ligand and to another Na+ cation as a monodentate ligand, resulting in the formation of chain-like aggregate structures along the c-axis. The fourth sulfonyl oxygen atom in the FSI− anions is not coordinated. Thus, the solvate structure of [Na(G3)][FSI] in the crystalline state is classified as an AGG I-type solvate. The Na−Oglyme distances vary from 2.463 ± 0.002 to 2.556 ± 0.002 Å. The Na−Oanion distances are 2.332 ± 0.002 Å for the FSI− anion that is coordinated via one oxygen atom and 2.430 ± 0.002 and 2.423 ± 0.002 Å for the FSI− anion that is coordinated via two oxygen atoms. The crystal structure of [Na(G3)][FSI] is similar to the structures of [Na(G4)][Tf2N] and [Na(G4)][FSI], which were reported by Mandai et al.33 In these structures, the Na+ cations are coordinated by one single G4 ligand and also linked together by the anions to form chain-like aggregates. However,

Figure 3. View of the packing diagram for [Na(G3)][FSI] along the aaxis. Cyan, Na; red, O; gray, C; blue, N; yellow, S; green, F; white, H.

since G4 has one more oxygen atom available for coordination, the anions coordinate to each Na+ cation via one oxygen atom. In the crystal structure of [Na(G3)2][FSI], the asymmetric unit consists of a half Na+ cation located on a 2-fold axis, two half G3 ligands, and one FSI− anion. The Na+ cations are coordinated by eight oxygen atoms of two perpendicularly oriented G3 ligands (complete ligands are generated by the local 2-fold axis), resulting in a total coordination number of eight (Figure 4). The FSI− anions are not coordinated and are disordered over four positions. Therefore, the solvate structure of [Na(G3)2][FSI] in the crystalline state is classified as a SSIPtype solvate. In the crystal packing, [Na(G3)2]+ cations and FSI− anions are ordered in a parallel arrangement along the bE


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NaFSI and G2-NaFSI mixtures in the range 780−890 cm−1 is shown in Figure 5. In the Raman spectrum of pure monoglyme, two bands are observed at 822 and 850 cm−1 which correspond to the coupled r(CH2)/ν(COC) modes of free G1 molecules in different conformations.51,52 In the spectrum of a 10:1 G1:NaFSI mixture, a third band is observed at 866 cm−1 which indicates the presence of G1 molecules that are coordinated to sodium cations. The two bands that are characteristic for free G1 molecules are also still present, since, at this ratio, G1 is present in a large excess compared to NaFSI. As the concentration of NaFSI in the binary mixtures is increased, the band at 866 cm−1 becomes more intense, whereas the bands at 850 and 822 cm−1 become less intense, which indicates that the amount of free G1 molecules in the mixtures decreases. Eventually, at a G1:NaFSI ratio of 3:1, the peaks at 850 and 822 cm−1 cannot be observed directly, indicating that at these compositions the vast majority of the G1 molecules in the mixture are coordinated to sodium cations. Furthermore, in the spectra of these solvate ionic liquids, two other small bands are observed at approximately 831 and 843 cm−1. Similar bands have been observed in sodium-containing SILs with G4 and G5 ligands, and they were assigned to vibrational modes of coordinated glyme molecules as well.33 The Raman spectrum of pure diglyme is significantly more complex than that of monoglyme and can be deconvoluted into three bands at ca. 807, 828, and 852 cm−1. Johansson et al. have extensively studied the vibrational modes of diglyme and longer glymes and found that the Raman spectrum of liquid diglyme is a combination of a large number of different conformers that exist in equilibrium.53 Thus, deconvolution of this multipeak in only three bands is a very simplified representation. However, they were able to assign the bands at ca. 807 and 852 cm−1, respectively, to the trans and gauche conformations of the C−C bond in diglyme. The band at 807 cm−1 corresponds to the coupled r(CH2)/tw(CH2) modes, and the band at 852 cm−1 corresponds to the coupled r(CH2)/ν(COC) modes of a free G2 molecule; thus, they can be used for solvation studies. In the spectrum of a 10:1 G2-NaFSI mixture, two new bands appear at 840 and 871 cm−1 due to the coordination of G2 molecules to sodium cations. Similarly to the G1-NaFSI mixtures, these bands become more intense as the NaFSI concentration increases, whereas the bands characteristic for free G2 become less intense. For the 2:1 mixture, these bands cannot be observed directly, indicating that virtually all G2 molecules are coordinated to sodium cations. In the 3:1 mixture, bands characteristic of free diglyme are still visible, in contrast to the 3:1 G1-NaFSI mixture. This result can easily be explained, since, in diglyme, one more ether oxygen atom is available for coordination than in monoglyme and hence fewer molecules are required to completely solvate a sodium cation. Binary mixtures of G3 and NaFSI in different ratios were also studied (Figure S6), but for the compositions 5:1−3:1, crystals of [Na(G3)2][FSI] formed in the solution, resulting in an inhomogeneous mixture of a liquid and solid. Therefore, their Raman spectra were not included. Overall, the concentration dependent behavior of G3-NaFSI mixtures is very similar to that of the G2-NaFSI mixtures. However, since the 2:1 mixture [Na(G3)2][FSI] is a crystalline solid at room temperature, some interesting features are observed in its Raman spectrum. The characteristic intense band for solvating glymes is shifted significantly to higher wavenumbers compared to the liquid samples (874 cm−1 vs 871 cm−1), which was also observed for

Figure 4. View of the crystal packing diagram of [Na(G3)2][FSI] along the a-axis. Disorder of the FSI− anions is shown. Cyan, Na; red, O; gray, C; blue, N; yellow, S; green, F; white, H.

axis. The Na−Oglyme distances vary from 2.525 ± 0.005 to 2.751 ± 0.006 Å. A summary of all of the Na−O distances of both complexes can be found in the Supporting Information (Table S2), as well as the thermal ellipsoid drawings of the solved structures and close-up views of the Na+ coordination (Figures S2−S5). The other complexes that could be isolated in the crystalline state, i.e., [Na(G1)3][FSI] and [Na(G2)2][FSI], were unsuitable for single crystal XRD due to their low melting point. However, their crystal structure can be estimated on the basis of existing literature. Crystal structures of various sodium salt complexes with tetrahydrofuran (THF), glymes, and cryptands have been reported by Bock et al.46,47 They found that, with monoglyme and diglyme, the Na+ ion forms the six-coordinate complex cations [Na(G1)3]+ and [Na(G2)2]+ in the crystalline phase, as these shorter glymes have a nearly ideal geometry to form a compact octahedral arrangement around Na+. Thus, it is very likely that [Na(G1)3][FSI] and [Na(G2)2][FSI] have a SSIP-type solvate structure in the crystalline state. One single triglyme molecule would be too strained in a tetrahedral conformation around Na+, so when available, Na+ will be coordinated by a second triglyme molecule.48 Bock et al. also reported the only other crystal structure containing a [Na(G3)2]+ cation. However, the triglyme ligands in this structure are coordinated to Na+ in a parallel fashion, resulting in a sandwich-type complex. Raman Spectroscopy. Raman spectroscopy has been used intensively to study interactions of alkali metal cations with crown ethers as well as glymes. In particular, the bands in the range 780−900 cm−1, which correspond to the coupled CH2 rocking (r(CH2)) and COC stretching (ν(COC)) modes, are useful in this regard.33,49,50 When a 1:1 complex is formed between an alkali metal cation and a crown ether or glyme ligand, a characteristic intense Raman band appears at ca. 870 cm−1, which corresponds to the coupled r(CH2) and ν(COC) modes or so-called “ring breathing mode” of coordinated ether molecules.49 This band is shifted significantly to higher wavenumbers compared to the analogous bands of uncoordinated crown ethers or glymes and can therefore be used as one of the fingerprint modes for the study of solvated alkali metal cations in glymes. In order to study the coordination states of the glyme molecules, Raman spectra were recorded for glyme-NaFSI mixtures with various glyme:NaFSI molar ratios. The concentration dependence of the Raman spectra for G1F


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the analogue SIL [Na(G3)2][Tf2N], and it is most likely because glymes are more tightly wrapped around the sodium cations in the ordered crystalline state vs the more disordered liquid state.54 The band at 828 cm−1 which is indicative for free glymes is also present in the spectrum, whereas the bands at 852 and 807 cm−1 are completely absent. While the Raman bands of the complex cations provide useful information on the coordination state of the glymes, it is also important to consider the bands assignable to the anions, since they hold information on the type of solvate structures that are formed, i.e., SSIPs, CIPs, or AGGs. For the FSI− anion, the range 700−770 cm−1, where the bands assignable to the S− N stretching modes (ν(SN)) are located, is particularly interesting, since these bands are highly sensitive to the coordination environment of the anion.49,55 The concentration dependence of the Raman spectra for G1-NaFSI, G2-NaFSI, and G3-NaFSI binary mixtures in the range 690−770 cm−1 is shown in Figure 6. For all three systems, a strong dependence is observed between the ν(SN) band and the NaFSI concentration. For the free, uncoordinated FSI− anion, this band is located at ca. 720 cm−1, as is observed in the spectra of the most dilute mixtures (20:1 molar ratio).56 Therefore, it can be concluded that at these compositions NaFSI is fully dissociated in SSIP solvates. As the NaFSI concentration increases, the ν(SN) band becomes significantly broader and shifted to higher wavenumbers as the S−N stretching vibrations are increasingly restricted due to coordination with sodium cations. Thus, these results indicate that, at higher concentrations, CIP- and AGGtype solvates are formed. When the three systems are compared to each other, an interesting trend is observed. The broadening and shifting to higher wavenumbers of the ν(SN) band occurs already at fairly dilute compositions for the G1-NaFSI system (10:1 molar ratio), whereas it occurs at more concentrated compositions for the G2-NaFSI (3:1 molar ratio) and even more so for the G3-NaFSI system (2:1 molar ratio). This is a strong indication that, for the G1-NaFSI system, CIPs and AGGs are formed even though G1 is present in a sufficient amount to completely solvate sodium cations. Due to the chelate effect, G2 and G3 have stronger solvating properties and thus competition with the FSI− anions only occurs at higher glyme:NaFSI ratios when there is an insufficient number of ether oxygen atoms available to completely solvate the sodium cations. In particular, for the G3-NaFSI system, no shift of the ν(SN) band is observed until there are less than two G3 ligands (eight ether oxygen atoms) available per sodium cation, which further confirms the SSIP-type solvate structure for [Na(G3)2][FSI] that was determined by XRD. The formation of CIPs in the 3:1 G1-NaFSI and 2:1 G2-NaFSI mixtures also implies that a small amount of free glymes is present. Since the solvate ionic liquid [Na(G3)][FSI] is solid at room temperature but is also stable as an undercooled liquid, both the crystalline and liquid phases were studied by Raman spectroscopy in order to investigate possible changes in solvation structure upon melting (Figure 7). Other than the small feature at 820 cm−1 in the spectrum of [Na(G3)][FSI] in the liquid state, no significant differences are observed in the range 780−890 cm−1 when compared to the crystalline state. This indicates that the coordination of triglyme to the sodium cation remains unchanged when the complex is melted. On the other hand, in the range 690−770 cm−1, the spectra of the crystalline and liquid complex are significantly different. From the XRD data, it was determined that the FSI−

Figure 5. Concentration dependence of the Raman spectra in the range 780−890 cm−1 for (a) G1-NaFSI and (b) G2-NaFSI mixtures. The Raman spectra of the pure G1 and G2 are also included in the respective figures. The blue and red curves are peak fits produced by the deconvolution of the multipeaks, where the blue curves are fitted to bands characteristic of free glymes and the red curves to bands characteristic of solvating glymes. G


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Figure 7. Raman spectra of [Na(G3)][FSI] in the crystalline and the undercooled liquid state in the range 690−890 cm−1, measured at room temperature.

the spectrum of the liquid, this band is significantly broader and shifted to lower wavenumbers, indicating that the FSI− anion is less restricted and probably occurs in a number of different environments that are in equilibrium with each other. NMR Results. The idea to use nuclear magnetic resonance (NMR) spectroscopy to study the solvation of sodium cations in solution is not new.57,58 Since the resonance frequency of a nucleus in a magnetic field highly depends on the local environment of the nucleus, NMR is a useful technique to investigate solvent−anion competition in NaFSI-glyme equimolar mixtures. Furthermore, many nuclei are NMR-sensitive and can be used to collect information on all of the components of the mixture. In this case, 23Na was used for the cation, 17O for the anion and solvent, and 13C for the solvent. 23 Na is a medium-NMR-sensitive nucleus with 100% abundance that yields resonances over a moderate chemical shift range. This nucleus has spin 3/2 and is therefore a quadrupolar nucleus. As a result, the lines in 23Na spectra are moderately broad and the line width increases with the asymmetry of the environment. Due to its medium sensitivity and high abundance, the resonances are intense and the signalto-noise ratio is good with a limited number of scans. The 23Na chemical shift (δ(23Na)) of a series of equimolar mixtures with the glyme:NaFSI ratio ranging from 50:1 to 1:1 was measured (Figure 8). To ensure that the G3-NaFSI samples were homogeneous liquids, the samples were measured at 80 °C. The 23Na NMR spectra of the 10:1 mixtures are included in the Supporting Information (Figure S7). The average standard deviation of δ(23Na) was 0.02 ppm.

Figure 6. Concentration dependence of the Raman spectra in the range 690−770 cm−1 for (a) G1−NaFSI, (b) G2−NaFSI, and (c) G3−NaFSI. G3−NaFSI 5:1−3:1 were solid/liquid mixtures and were omitted from the plot. G3−NaFSI 2:1 was completely solid, and 1:1 was measured as an undercooled liquid.

anion only occurs in AGG I-type solvate structures in the crystalline state. Therefore, the ν(SN) mode is quite restricted, resulting in a narrow, slightly shouldered peak at 740 cm−1. In H


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Å for an eight-coordinate complex.46 The latter value is in good agreement with the average sodium−oxygen bond distance found for [Na(G3)2][FSI] (2.622 ± 0.005 Å). Since the value of δ(23Na) depends on the orbital overlap of Na+ and the coordinating atoms in its solvation shell, it is evident that the coordination bond distances also have a significant influence.62,63 In solid state magic angle spinning (MAS) NMR literature, the same relationship between δ(23Na) and the sodium−oxygen bond distance is often observed and used to gain information on the sodium coordination environment in minerals and glasses.64,65 Therefore, it can be concluded that the low δ(23Na) value of the G3-NaFSI system is caused by the larger sodium−oxygen bond distance in the solvated species. Furthermore, the small increase of δ(23Na) upon transition from a 2:1 to a 1:1 G3:NaFSI ratio can also be explained. The XRD and Raman measurements both show that FSI− anions are present in the solvation shell around Na+ in the 1:1 mixture ([Na(G3)][FSI]), as opposed to the 2:1 mixture ([Na(G3)2][FSI]), where Na+ is completely solvated by G3 ligands. However, the coordination number of sodium in [Na(G3)][FSI] is seven, as opposed to eight in [Na(G3)2][FSI]. As a result, the average sodium−oxygen bond distance in [Na(G3)][FSI] (2.453 ± 0.002 Å) is significantly shorter than that in [Na(G3)2][FSI] (2.622 ± 0.005 Å), explaining the increase of δ(23Na). Thus, 23Na NMR can also be used to reveal information about the coordination number of complexes. Another interesting trend is observed when the three systems are compared: the change of δ(23Na), i.e., the formation of CIPs and AGGs, starts at much lower NaFSI concentrations for the G1-NaFSI system than for the G2-NaFSI and G3-NaFSI systems, which is in complete accordance with the trends observed in the Raman spectra. 17 O is a nucleus with a low NMR sensitivity bandwidth and a low abundance (0.038%), displaying lines over a very wide chemical shift range. This quadrupolar nucleus has spin 5/2, and therefore, the lines are rather broad, even for small molecules. This combined with the high cost of 17O-enriched samples explains why almost no 17O NMR studies can be found in the older literature. However, due to advances in NMR technology, it has become possible to acquire reliable results with nonenriched samples, and therefore, the 17O NMR technique has gained popularity in recent reports.66,67 We found that 17O NMR is a useful technique for this study, since coordination of the glymes as well as the FSI− anions to sodium cations occurs via oxygen atoms and, hence, significant shifts can be expected for the 17O lines. The 17O chemical shift (δ( 17 O)) of a series of equimolar mixtures with the glyme:NaFSI ratio ranging from 50:1 to 1:1 was measured. The concentration dependence of the 17O resonances of the sulfonyl oxygen atoms in the FSI− anion is shown in Figure 9. Some data points are missing either because the signal-to-noise ratio was too small (low concentrations) or because the lines were very broad and not reproducible due to the high viscosity of some of the samples (high concentrations). The 17O NMR spectra of the 10:1 mixtures are included in the Supporting Information (Figure S8). The average standard deviation of δ(17O) of the sulfonyl signals was 0.04 ppm. The lines of the FSI− sulfonyl oxygen atoms can be found downfield in the 17O spectrum at about 170 ppm. Overall, the 17 O chemical shifts assignable to the FSI− anion follow a trend very similar to the 23Na shifts of the sodium cation, which is expected, since this trend is caused by the coordination of these ions to each other. In the dilute regime, all of the anions are

Figure 8. Concentration dependence of the 23Na chemical shift. G1 and G2 samples were measured at RT, and G3 samples at 80 °C.

Since the exchange of Na+ is fast compared with the NMR time scale, all of the observed signals are weighed averages of Na+ in different solvation structures. According to Erlich et al., the 23Na chemical shift of a solvated sodium ion is directly proportional to the Gutmann donor number (DN) of the solvating ligands, where the lower the DN value is, the more upfield (negative) the shift is.59 The DN value is a measure for the coordinating abilities of a ligand. Therefore, δ(23Na) is higher (downfield shifted) for a complex with strongly coordinating ligands (high DN) compared to a complex with weaker coordinating ligands (low DN). For the G1-NaFSI and G2-NaFSI systems, this holds true: under dilute conditions, δ(23Na) does not change significantly with the concentration, since the glymes are present in a large excess and can completely solvate the sodium cations to form SSIPs. As the NaFSI concentration is increased, a strong decrease of the 23Na chemical shift over several ppm is observed as the stronger coordinating glyme ligands around the sodium cations (DN = 20 kcal mol−1 for G1)60 are gradually exchanged by weaker coordinating FSI− anions (DN ≈ 11 kcal mol−1 for [emim][Tf2N] as a reference),61 resulting in the formation of CIPs and AGGs. On the other hand, for the G3-NaFSI system, δ(23Na) is already very low under dilute conditions and a small increase is observed in the transition from a 2:1 to a 1:1 G3:NaFSI ratio. Also, when the three systems are compared, the following trend is observed for the 23Na chemical shifts: δ[Na(G3)n] ≪ δ[Na(G2)n] < δ[Na(G1)n], whereas the reverse trend would be expected on the basis of the respective coordinating abilities of the glymes due to the chelate effect. To ensure that this observation was not a result of the elevated measuring temperature of the G3 samples, the 10:1 sample was measured both at 25 and 80 °C and the difference in δ(23Na) was only 1 ppm, which is considerably lower than the difference in δ(23Na) between the G2 and G3 series. When comparing these three systems, the differences in coordination number and the sodium−oxygen bond distance also have to be considered. As mentioned in the XRD section, Na+ forms six-coordinate complexes with monoglyme and diglyme and eight-coordinate complexes with triglyme. Due to the increased steric hindrance associated with the coordination of two triglyme ligands around Na+ as opposed to two smaller diglyme or three monoglyme ligands, the average sodium−oxygen bond distance increases from 2.39 ± 0.04 Å for a six-coordinate complex to 2.62 ± 0.13 I


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Figure 9. Concentration dependence of the 17O chemical shifts of the FSI− anion. G1 and G2 samples were measured at RT, and G3 samples at 80 °C.

free in solution, resulting in a constant 17O chemical shift, while, in the concentrated regime, CIPs and AGGs are formed, resulting in a gradual decrease of the shift. This decrease starts at lower concentrations for the G1-NaFSI system compared to the other two which is again in agreement with the 23Na NMR and Raman spectroscopy results. Since glyme ligands coordinate to metal cations via their ether oxygen atoms, the 17O NMR spectra are also useful to study their coordination state. Ether oxygen atoms yield lines that are found more upfield in the 17O NMR spectra compared to the sulfonyl oxygen signals. For pure glymes, one (G1) or two (G2, G3) lines are observed. For G2 and G3, the line corresponding to the central ether oxygen atom is shifted more downfield (∼0 ppm) compared to the lines of the terminal ether oxygen atoms (∼ −21 ppm). The concentration dependence of these lines is shown in Figure 10. The average standard deviation of δ(17O) of the ether signals was 0.05 ppm. For all of the ether 17O signals, a similar gradual decrease of the chemical shift is observed as the NaFSI concentration is increased. This observation can easily be rationalized by considering that the measured signals are all weighed averages of coordinated and free ether oxygen atoms. It is already observed by Peng et al. that the dissolution of alkali metal salts in glymes causes an upfield shift for the ether 17O signals.67 Thus, as the NaFSI concentration is increased, a higher fraction of glymes in the mixture is coordinated to sodium cations, resulting in an upfield shifted average signal. This conclusion is also in agreement with our observations for the vibrational modes of the glymes in the Raman study. Besides the ether oxygen atoms in glymes, the neighboring carbon atoms were also used to study solvation effects. Since they are located right next to the coordinating oxygen atoms, significant shifts in the 13C spectra are expected. Furthermore, since the 13C nucleus has spin −1/2 and is therefore not quadrupolar, the resulting lines are very narrow and precise and do not broaden significantly due to viscosity, in contrast to 23 Na and 17O lines. The concentration dependence of the lines corresponding to the most central CH2 groups can be found in the Supporting Information (Figure S9). The chemical shift of the 13C lines follows the same trend as that of the 17O lines of the neighboring oxygen atoms. However, since 13C is not a quadrupolar nucleus, shifts in a wider concentration range

Figure 10. Concentration dependence of the 17O chemical shifts of (a) terminal and (b) central ether oxygen atoms in G1, G2, and G3. G1 and G2 samples were measured at RT, and G3 samples at 80 °C.

could be determined more precisely. The concentration dependent behavior of the other methylene signals was determined as well, and similar trends were observed. Conductivity of Glyme-NaFSI Binary Mixtures. The conductivity of an electrolyte is an important property to consider when designing a battery, since it can have a significant influence on the cycling performance. Two important factors that determine the ionic conductivity of an electrolyte are the total concentration of charge carriers and the mobility of those charge carriers. The mobility of charge carriers, in turn, depends on the viscosity of the electrolyte and the nature of the solvated species that are present in solution. A neutral solvate such as a contact ion pair is not attracted to the electrodes, as opposed to an ionic solvate such as a solvent separated ion pair, and is therefore less mobile. In order to investigate the influence of these factors on the performance of our glyme-NaFSI electrolytes, their specific ionic conductivity was measured by impedance spectroscopy (Figure 11). A summary of the specific ionic conductivity, dynamic viscosity, density, and sodium concentration can be found in Table 3. The G3-NaFSI electrolytes were not included in this study, since many of the more concentrated samples are either solid or partially solid at the measuring temperature. For both the G1-NaFSI and G2-NaFSI electrolytes, an initial increase of conductivity is observed; then, a maximum is reached and the conductivity rapidly drops as the fraction of J


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that of the G2-NaFSI electrolytes with the same ratio. This is a consequence of their lower viscosity on one hand and their higher sodium ion concentration at the same glyme:NaFSI ratio on the other hand. The latter is a result of the lower molar volume of G1 (104 cm3 mol−1) compared to G2 (143 cm3 mol−1). However, after reaching the maximum conductivity, the viscosity of the G1-NaFSI electrolytes increases faster with the glyme:NaFSI ratio compared to the G2-NaFSI electrolytes, resulting in a sharper drop in conductivity. Second, the maximum conductivity is reached at a 4:1 glyme:NaFSI ratio for the G2-NaFSI electrolytes as opposed to a 6:1 ratio for the G1-NaFSI electrolytes, which corresponds to a sodium ion concentration of approximately 1.53 and 1.43 mol dm−3, respectively. This cannot be explained by only considering the viscosity: the G1-NaFSI electrolytes with 5:1 and 4:1 ratio are significantly less viscous than their G2-NaFSI analogues, yet their conductivity already decreases at this point. However, the conductivity of the G2-NaFSI analogues still increases at these ratios. Thus, the higher degree of CIP and AGG formation for the G1-NaFSI electrolytes is the most probable cause of this behavior. In conclusion, sodium ion concentration and viscosity seem to be the two most important factors that determine the conductivity of an electrolyte, while CIP and AGG formation has a smaller influence. This is particularly well illustrated by comparing the 2:1 electrolytes of both solvents, i.e., the solvate ionic liquids [Na(G1)2][FSI] and [Na(G2)2][FSI]. Despite having a significantly higher viscosity, [Na(G1)2][FSI] is equally conductive as [Na(G2)2][FSI], since its sodium ion concentration is higher as well. Electrochemical Results. In order to be useful as a battery electrolyte, the electrochemical window of the electrolyte should be wider than the operating potential window of the battery. Therefore, the electrochemical window of all of the solvate ionic liquids that were liquid at room temperature was determined by linear sweep voltammetry with a platinum working electrode (Figure S10). For all of the SILs, the cathodic limit corresponds to the deposition of sodium metal which occurs at a slight overpotential of −0.1 V vs Na+/Na. The anodic limit of the SILs is around 4.6 V vs Na+/Na and corresponds to the decomposition of the ligands. Overall, all of the SILs have an electrochemical window of approximately 4.7 V, which is similar to the electrochemical window of the SIL [Na(G5)][Tf2N], reported by Terada et al.34 Since the electrodes used in the cycling experiments can also have a significant influence on the limiting potentials of the electrochemical window, the voltage range was chosen well within the determined boundaries (1.5−3.7 V vs Na+/Na). Charge−discharge cycling is the most common experiment to evaluate the performance of the electrodes and the electrolyte in a practical battery configuration. Even when an electrolyte shows excellent redox behavior and stability, its performance may vary significantly when used under typical battery conditions, depending on factors such as the compatibility with the electrodes, possible side reactions with the electrode, and the consequent formation of surface layers. Therefore, electrochemical studies in a battery setup are indispensable for the evaluation of practical application of any electrolyte. Here, we report, as a proof-of-concept, the charge− discharge profiles of the battery cells containing the G1-based SILs. The G1-based SILs were selected for electrochemical testing because they are all liquid at room temperature which makes the fabrication of coin cells easier, and since their dynamic viscosity and ionic conductivity varies over a large

Figure 11. Specific ionic conductivity (σ) of electrolytes vs the glyme:NaFSI ratio, measured at ambient glovebox temperature (28 °C).

Table 3. Specific Ionic Conductivity (σ), Dynamic Viscosity (ηdyn), Density (ρ), and Sodium Ion Concentration of Glyme-NaFSI Electrolytes, Measured at 28 °C electrolyte

σ (Ω−1 m−1)

40:1 30:1 20:1 15:1 10:1 8:1 6:1 5:1 4:1 3:1 2:1 1:1

0.47 0.65 1.00 1.34 1.84 1.94 2.14 2.11 1.71 1.32 0.66 0.07

40:1 30:1 20:1 15:1 10:1 8:1 6:1 5:1 4:1 3:1 2:1 1:1

0.21 0.30 0.47 0.67 1.00 1.04 1.11 1.19 1.20 0.97 0.67 0.16

ηdyn (10−3 Pa s−1)

ρ (kg m−3)

G1-NaFSI 0.6 890 0.6 910 0.7 930 0.8 950 1.0 990 1.4 1020 2.0 1060 2.7 1100 4.5 1140 10 1190 39 1280 561 1480 G2-NaFSI 1.2 0.96 1.3 0.97 1.4 0.98 1.6 1.00 2.0 1.03 2.3 1.04 2.9 1.08 3.7 1.10 5.5 1.13 9.4 1.18 23 1.24 241 1.35

Na+ concentration (mol dm−3) 0.23 0.31 0.46 0.61 0.90 1.10 1.43 1.68 2.02 2.51 3.34 5.05 0.17 0.23 0.34 0.45 0.67 0.82 1.07 1.26 1.53 1.95 2.63 4.00

NaFSI in the solution is further increased. This indicates that the increase in concentration of charge carriers is initially dominant over the decrease of their mobility due to the viscosity. Eventually, the increase in viscosity becomes too large and the roles are reversed, leading to the sharp decrease in conductivity. Although the overall trend in conductivity vs the glyme:NaFSI ratio is very similar for the monoglyme and diglyme electrolytes, some significant differences can be observed. First, for the glyme:NaFSI ratios 40:1−6:1, the conductivity of the G1-NaFSI electrolytes is roughly double K


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The Journal of Physical Chemistry B range, it is interesting to investigate the influence on the cycling performance. The SILs were subjected to galvanostatic cycling for 20 cycles on electrodes containing P2-phase Na2/3[Mn0.55Ni0.30Co0.15]O2 which proved to be a promising MNC-type active material for Na+ intercalation in a previous study, showing excellent structural stability during cycling (Figure 12).45 Furthermore, the relatively lower cobalt composition is highly favorable in the context of cost, since cobalt is an expensive and critical raw material.

Figure 13. (a) Cycling performance at 1 C rate of the cell with [Na(G1)3][FSI] electrolyte, Na2/3[Mn0.55Ni0.35Co0.15]O2 cathode and Na anode. (b) Rate capability of the cell with [Na(G1)3][FSI] electrolyte, Na2/3[Mn0.55Ni0.35Co0.15]O2 cathode and Na anode.

the initial capacity values after successive cycles at different current densities (0.1, 1, and 2 C). This can be attributed to the combination of the high conductivity and sodium concentration of [Na(G1)3][FSI]. These results are encouraging, considering that the studied systems show reasonable sodium-ion diffusion and their charge−discharge cycling performance is comparable to the cells with ethylene carbonate based electrolytes.

Figure 12. (a) Comparison of specific capacity of cells with different G1-based electrolytes, Na2/3[Mn0.55Ni0.35Co0.15]O2 cathode and Na anode cycled at 0.1 C rate. (b) Charge−discharge profile of cell with [Na(G1)3][FSI] electrolyte, Na2/3[Mn0.55Ni0.35Co0.15]O2.

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CONCLUSIONS Several complexes with the general formula [Na(L)n][FSI] (with L = G1, G2, or G3, n = 1, 2, or 3) were prepared and characterized. Many of the prepared SILs are liquid at room temperature and have a reasonably low viscosity, making them promising electrolytes for sodium-based secondary batteries. Their volatility was significantly reduced compared to dilute electrolytes, but they still had a considerable vapor pressure. Therefore, these complexes can be classified as poor solvate ionic liquids, according to the criteria proposed by Mandai et al.23 The SILs and similar dilute electrolytes were thoroughly studied by single crystal XRD, Raman spectroscopy, and multinuclear NMR spectroscopy in order to reveal information on their structure in the liquid state and solid state when possible. The results of the three methods were in good

For all three electrolytes, the cells demonstrate stable cycling behavior, but the best performance in terms of specific capacity, reversibility, and stability is observed for the cell with the [Na(G1)3][FSI] electrolyte. A discharge capacity of 120 mA h g−1 is delivered at 0.1 C rate, which corresponds to the reversible extraction of 0.5 Na per unit cell of the cathode material. The charge−discharge efficiency of all of the cells is almost 100%. Further tests on the cell with [Na(G1)3][FSI] electrolyte demonstrate reasonable cycling performance for 100 cycles at 1 C rate (Figure 13). At this rate, an initial charging capacity of 84 mA h g−1 was obtained. About 88% of the initial discharge capacity is retained after 100 cycles at 1 C rate. In addition, as shown in Figure 13, the cell with this electrolyte exhibits a good rate capability, as indicated by the retention of L


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transition temperature; RT, room temperature; ηdyn, dynamic viscosity; ρ, density; XRD, X-ray diffraction; δ, chemical shift; DN, donor number; MAS, magic angle spinning; σ, specific ionic conductivity

agreement with each other and showed that a significant amount of contact ion pairs and aggregate solvates was present in the G1-based electrolytes, even under dilute conditions. Solvent separated ion pairs were the main species in the G2and G3-based electrolytes, even under very concentrated conditions. However, the presence of CIPs and AGGs only had a minor influence on the conductivity, since the G1-based electrolytes were by far the most conductive mainly due to their low viscosity. In a proof-of-concept study, the G1-based SILs were found to be promising electrolytes in half-cells with Na2/3[Mn0.55Ni0.35Co0.15]O2 cathodes, showing stable cycling over 20 cycles. [Na(G1)3][FSI] was a particularly successful electrolyte, as it was cycled over 100 cycles and achieved good rate capability.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcb.7b10158. Synthesis of complexes, crystallographic data of [Na(G3)][FSI] and [Na(G3)2][FSI] (Table S1), sodium− oxygen bond distances of [Na(G3)][FSI] and [Na(G3)2][FSI] (Table S2), dynamic and isothermal TGA curves of [Na(G1)3][FSI], [Na(G2)2][FSI], and [Na(G3)2][FSI] (Figure S1), thermal ellipsoid models and additional images of the crystal structures of [Na(G3)][FSI] and [Na(G3)2][FSI] (Figures S2−S5), Raman spectra of G3:NaFSI mixtures in the range 690−780 cm−1 (Figure S6), 23Na and 17O NMR spectra of 10:1 glyme:NaFSI mixtures (Figures S7 and S8), concentration dependence of δ(13C) for glyme:NaFSI mixtures (Figure S9), and linear sweep voltammograms for G1 and G2 solvate ionic liquids (Figure S10) (PDF)

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REFERENCES

AUTHOR INFORMATION

Corresponding Author

*Phone: +32 16 32 74 46. E-mail: koen.binnemans@kuleuven. be. ORCID

Pieter Geysens: 0000-0002-8846-7777 Koen Binnemans: 0000-0003-4768-3606 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge financial support by the KU Leuven (project KP/14/005). The authors thank Karel Duerinckx for his support with the NMR measurements.

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ABBREVIATIONS FSI, bis(fluorosulfonyl)imide; NMR, nuclear magnetic resonance; SSIP, solvent separated ion pair; CIP, contact ion pair; AGG, aggregate; SIL, solvate ionic liquid; LIB, lithium ion battery; NIB, sodium ion battery; EC, ethylene carbonate; PC, propylene carbonate; DMC, dimethyl carbonate; G1, DME, monoglyme; G2, diglyme; G3, triglyme; NVP, sodium vanadium phosphate; HC, hard carbon; DSC, differential scanning calorimetry; TGA, thermogravimetric analysis; LSV, linear sweep voltammetry; PVDF, polyvinylidene difluoride; G4, tetraglyme; G5, pentaglyme; Tf2N, bis(trifluoromethylsulfonyl)imide; Tm, melting point; Tg, glass M


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Solvation Structure of Sodium Bis(fluorosulfonyl)imide-Glyme Solvate

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