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CURATOR
A pinboard by
Ryoichi Tatara

Postdoctoral Associate, Massachusetts Institute of Technology

PINBOARD SUMMARY

To study/understand reaction mechanism and increase capacity/reversibility/lifetime of Li battery

Rechargeable Li-ion batteries are currently used in mobile phones, laptops, electric vehicle etc. and have attracted increasing attention, because we are of course interested in increasing its capacity and lifetime, for example to keep our mobile phone battery much longer! Therefore, not only current Li-ion battery, but also next generation batteries, so called “beyond Li-ion” such as sulfur battery, air battery are promising candidate, because they have much higher energy density than the rechargeable batteries designed thus far. However, a number of challenges exist in the way of realizing reversible and stable operations of beyond-Li-ion batteries. For example, discharge intermediate easily reacts or dissolves with various organic compounds, which severely affects charge-discharge performance. Thus, to avoid undesired side reactions, the electrolyte should be chemically stable.

Conventional electrolytes used for Li-ion battery research (around 1 mol/L) typically have a large proportion of “free” solvent molecules. In solutions containing Li salts, due to electrostatic and induction interactions between Li+ ion and solvent, Li+ is solvated by solvent molecules and forms a solvated [Li(solvent)n]+ cation. With increasing concentration of Li salt, the amount of free solvent that does not participate in the solvation of Li+ is decreased. In the extreme case, all of the solvent molecules are involved in the solvation, and free solvent is not present in solution. Electrolyte solutions with extremely high salt concentrations (≥3 mol/L) have been recently termed “solvent-in-salt” and/or “superconcentrated” electrolytes. Not only do these highly concentrated electrolytes have interesting physicochemical and electrochemical properties, but also solvents complexed with Li+ might have increased chemical stability and electrochemical stability relative to free solvents. Recent use of highly concentrated electrolyte solutions in Li−Sulfur batteries has led to highly efficient discharge/charge characteristics of Li−Sulfur cells, which can be attributed to reduced solubility of lithium polysulfide reaction intermediates with increasing salt concentration.

In the present study, I am focusing on both current Li-ion battery and beyond Li-ion battery, and study the concentration effect of reaction mechanism. The work for oxygen battery in highly concentrated electrolytes has already published (R. Tatara et al., J. Phys. Chem. C, 121, 9162, 2017).

7 ITEMS PINNED

Oxygen Reduction Reaction in Highly Concentrated Electrolyte Solutions of Lithium Bis(trifluoromethanesulfonyl)amide/Dimethyl Sulfoxide

Abstract: The performance of current Li–air batteries is greatly limited by critical obstacles such as electrolyte decomposition, high charging overpotentials, and limited cycle life. Thus, much effort is devoted to fundamental studies to understand the mechanisms of discharge/charge processes and overcome the above-mentioned obstacles. In particular, the search for new stable electrolytes is vital for long-lasting and highly cyclable batteries. The highly reactive lithium superoxide intermediate (LiO2) produced during discharge process can react with the electrolyte and produce a variety of byproducts that will shorten battery life span. To study this degradation mechanism, we investigated oxygen reduction reaction (ORR) in highly concentrated electrolyte solutions of lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA])/dimethyl sulfoxide (DMSO). On the basis of rotating ring disk electrode measurements, we showed that LiO2 dissolution can be limited by increasing lithium salt concentration over 2.3 mol dm–3. Our Raman results suggested that this phenomenon can be related to lack of free DMSO molecules and increasing DMSO–Li+ interactions with higher Li+ concentration. X-ray diffraction measurements for the products of ORR suggested that the side reaction of DMSO with Li2O2 and/or LiO2 could be suppressed by decreasing the solubility of LiO2 in highly concentrated electrolytes.

Pub.: 12 Apr '17, Pinned: 31 Aug '17

Li(+) solvation in glyme-Li salt solvate ionic liquids.

Abstract: Certain molten complexes of Li salts and solvents can be regarded as ionic liquids. In this study, the local structure of Li(+) ions in equimolar mixtures ([Li(glyme)]X) of glymes (G3: triglyme and G4: tetraglyme) and Li salts (LiX: lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA]), lithium bis(pentafluoroethanesulfonyl)amide (Li[BETI]), lithium trifluoromethanesulfonate (Li[OTf]), LiBF4, LiClO4, LiNO3, and lithium trifluoroacetate (Li[TFA])) was investigated to discriminate between solvate ionic liquids and concentrated solutions. Raman spectra and ab initio molecular orbital calculations have shown that the glyme molecules adopt a crown-ether like conformation to form a monomeric [Li(glyme)](+) in the molten state. Further, Raman spectroscopic analysis allowed us to estimate the fraction of the free glyme in [Li(glyme)]X. The amount of free glyme was estimated to be a few percent in [Li(glyme)]X with perfluorosulfonylamide type anions, and thereby could be regarded as solvate ionic liquids. Other equimolar mixtures of [Li(glyme)]X were found to contain a considerable amount of free glyme, and they were categorized as traditional concentrated solutions. The activity of Li(+) in the glyme-Li salt mixtures was also evaluated by measuring the electrode potential of Li/Li(+) as a function of concentration, by using concentration cells against a reference electrode. At a higher concentration of Li salt, the amount of free glyme diminishes and affects the electrode reaction, leading to a drastic increase in the electrode potential. Unlike conventional electrolytes (dilute and concentrated solutions), the significantly high electrode potential found in the solvate ILs indicates that the solvation of Li(+) by the glyme forms stable and discrete solvate ions ([Li(glyme)](+)) in the molten state. This anomalous Li(+) solvation may have a great impact on the electrode reactions in Li batteries.

Pub.: 04 Mar '15, Pinned: 31 Aug '17

Li+ Solvation and Ionic Transport in Lithium Solvate Ionic Liquids Diluted by Molecular Solvents

Abstract: An equimolar mixture of lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA]) and either triglyme (G3) or tetraglyme (G4) yielded stable molten complexes: [Li(G3)][TFSA] and [Li(G4)][TFSA]. These are known as solvate ionic liquids (SILs). Glyme-based SILs have thermal and electrochemical properties favorable for use as lithium-conducting electrolytes in lithium batteries. However, their intrinsically high viscosities and low ionic conductivities prevent practical application. Therefore, we diluted SILs with molecular solvents in order to enhance their ionic conductivities. To determine the stabilities of the complex cations in diluted SILs, their conductivity and viscosity, the self-diffusion coefficients, and Raman spectra were measured. [Li(G3)]+ and [Li(G4)]+ were stable in nonpolar solvents, that is, toluene, diethyl carbonate, and a hydrofluoroether (HFE); however, ligand exchange took place between glyme and solvent when polar solvents, that is, water and propylene carbonate, were used. In acetonitrile (AN) mixed solvent complex cations [Li(G3)(AN)]+ and [Li(G4)(AN)]+ were formed. [Li(G4)][TFSA] was more conductive than [Li(G3)][TFSA] when diluted with nonpolar solvents due to the greater ionic dissociativity in [Li(G4)][TFSA] mixtures. In view of the stability of the Li–glyme complex cations, the enhanced ionic conductivities, and the intrinsic electrochemical stabilities of the diluting solvents, [Li(G4)][TFSA] diluted by toluene or HFE, can be a candidate for an alternative battery electrolyte.

Pub.: 15 Dec '15, Pinned: 31 Aug '17

Oxidative-stability enhancement and charge transport mechanism in glyme-lithium salt equimolar complexes.

Abstract: The oxidative stability of glyme molecules is enhanced by the complex formation with alkali metal cations. Clear liquid can be obtained by simply mixing glyme (triglyme or tetraglyme) with lithium bis(trifluoromethylsulfonyl)amide (Li[TFSA]) in a molar ratio of 1:1. The equimolar complex [Li(triglyme or tetraglyme)(1)][TFSA] maintains a stable liquid state over a wide temperature range and can be regarded as a room-temperature ionic liquid consisting of a [Li(glyme)(1)](+) complex cation and a [TFSA](-) anion, exhibiting high self-dissociativity (ionicity) at room temperature. The electrochemical oxidation of [Li(glyme)(1)][TFSA] takes place at the electrode potential of ~5 V vs Li/Li(+), while the oxidation of solutions containing excess glyme molecules ([Li(glyme)(x)][TFSA], x > 1) occurs at around 4 V vs Li/Li(+). This enhancement of oxidative stability is due to the donation of lone pairs of ether oxygen atoms to the Li(+) cation, resulting in the highest occupied molecular orbital (HOMO) energy level lowering of a glyme molecule, which is confirmed by ab initio molecular orbital calculations. The solvation state of a Li(+) cation and ion conduction mechanism in the [Li(glyme)(x)][TFSA] solutions is elucidated by means of nuclear magnetic resonance (NMR) and electrochemical methods. The experimental results strongly suggest that Li(+) cation conduction in the equimolar complex takes place by the migration of [Li(glyme)(1)](+) cations, whereas the ligand exchange mechanism is overlapped when interfacial electrochemical reactions of [Li(glyme)(1)](+) cations occur. The ligand exchange conduction mode is typically seen in a lithium battery with a configuration of [Li anode|[Li(glyme)(1)][TFSA]|LiCoO(2) cathode] when the discharge reaction of a LiCoO(2) cathode, that is, desolvation of [Li(glyme)(1)](+) and insertion of the resultant Li(+) into the cathode, occurs at the electrode-electrolyte interface. The battery can be operated for more than 200 charge-discharge cycles in the cell voltage range of 3.0-4.2 V, regardless of the use of ether-based electrolyte, because the ligand exchange rate is much faster than the electrode reaction rate.

Pub.: 22 Jul '11, Pinned: 29 Nov '17