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Oxygen Reduction Reaction in Highly Concentrated Electrolyte Solutions of Lithium Bis(trifluoromethanesulfonyl)amide/Dimethyl Sulfoxide Ryoichi Tatara, David G. Kwabi, Thomas P. Batcho, Michal Tulodziecki, Kenta Watanabe, Hoi-Min Kwon, Morgan L. Thomas, Kazuhide Ueno, Carl V. Thompson, Kaoru Dokko, Yang Shao-Horn, and Masayoshi Watanabe J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.7b01738 • Publication Date (Web): 12 Apr 2017 Downloaded from http://pubs.acs.org on April 13, 2017

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The Journal of Physical Chemistry C 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.

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The Journal of Physical Chemistry

J. Phys. Chem. C

Oxygen Reduction Reaction in Highly Concentrated Electrolyte Solutions of Lithium Bis(trifluoromethanesulfonyl)amide/Dimethyl Sulfoxide Ryoichi Tatara, a,b,c David G. Kwabi, b Thomas P. Batcho, d Michal Tulodziecki, c Kenta Watanabe,a Hoi-Min Kwon, a Morgan L. Thomas,a Kazuhide Ueno, a Carl V. Thompson, d Kaoru Dokko,a,* Yang Shao-Horn,b,c,d,* Masayoshi Watanabe a [a] Department of Chemistry and Biotechnology, Yokohama National University,79-5 Tokiwadai, Hodogaya-ku, Yokohama 240-8501, Japan [b] Department of Mechanical Engineering, [c] Research Laboratory of Electronics, [d] Department of Materials Science and Engineering, Massachusetts Institute of Technology, 77 Massachusetts Ave., Cambridge, MA 02139, USA CORRESPONDING AUTHOR FOOTNOTE: To whom correspondence should be addressed. Telephone/Fax: +81-45-339-3942. E-mail: [email protected] (K. D.) Telephone/Fax: +1-617-253-2259. E-mail: [email protected] (Y. S.-H)

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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 in order 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). Based on 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. Xray 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.


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1. Introduction The oxygen reduction reaction (ORR) has been widely used for electrochemical energy storage applications such as fuel cells and rechargeable metal-air batteries. Li-air batteries have attracted much attention owing to the high theoretical energy density. 1-3 The cathode reaction of an Li-air cell using aprotic electrolyte solutions containing Li salts is notably different from ORR in aqueous solutions, in that the cathode reaction product is solid lithium peroxide (2Li+ + O2 + 2e− → Li2O2). There have been numerous reports on the ORR mechanism, 4-10 the possible role for catalysts for ORR and oxygen evolution reaction (OER),11-13 and the use of redox mediators during OER.14-15 It is widely believed that ORR in aprotic organic solutions containing Li salt proceeds as follows: 2-4 O2 + Li+ + e− → LiO2

(Electrochemical reduction of O2)

(1)

LiO2 + Li+ + e− → Li2O2

(Electrochemical reduction of LiO2)

(2)

2LiO2 → Li2O2 + O2

(Chemical disproportionation of LiO2)

(3)

A number of reports have shown that discharge at higher overpotentials/current densities tends to favor electrochemical reduction of LiO2 (eq 2) over disproportionation (eq 3), and that these two pathways influence Li2O2 morphology.16 It is also well known that the superoxide radical (O2•–) produced right after electrochemical O2 reduction (eq 1) can dissolve in aprotic electrolyte solutions.16-23 The chemical reactivity of superoxide with electrolytes24-26 and carbon22 is high, resulting in parasitic reactions, which greatly influence the charge and discharge performance of Li-air cells. For example, carbonate-based electrolyte solutions such as LiPF6 dissolved in a mixed solvent of ethylene carbonate and diethyl carbonate, which is used in commercial Li-ion batteries, readily decompose due to the nucleophilic attack of superoxide species and are not suitable for Li-air cells.25 On the other hand, ethers, such as 1,2dimethoxyethane (DME)18 and tetraglyme,27 and dimethyl sulfoxide (DMSO)17 are relatively stable against superoxide. Although Bruce et al. have shown reversible Li-O2 chemistry in a DMSO-based electrolyte using a porous gold electrode,28 recent work using carbon nanotube electrodes revealed


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instability of DMSO as indicated by the conversion of discharged Li2O2 to LiOH over time in the electrolyte.29 Conventional electrolytes used for Li-O2 battery research (around 1 M) typically have a large number 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.30-33 With increasing concentration of Li salt, the amount of free solvent which 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.34 Electrolyte solutions with extremely high salt concentrations (≥ 3 mol dm−3) have been recently termed “solvent-in-salt” and/or “superconcentrated” electrolytes.35-36 An interesting example is the 1:1 mixture of Li[TFSA] and glyme (triglyme or tetraglyme); Li+ and glyme form a stable (long-lived) complex cation, [Li(glyme)]+, owing to the chelate effect. Therefore, the 1:1 mixture can be regarded as a “solvate ionic liquid” consisting of only solvate [Li(glyme)]+ cations and [TFSA]− anions.34, 37-41 Not only do these highly concentrated electrolytes have interesting physicochemical and electrochemical properties,35,

42

but also solvents

complexed with Li+ might have increased chemical stability and electrochemical stability (shifted HOMO and LUMO of the solvent by the complex formation with Li+

38, 43

) relative to free solvents.

Recent use of highly concentrated electrolyte solutions in Li-S batteries,36,

44-46

has led to highly

efficient discharge/charge characteristics of Li-S cells, which can be attributed to reduced solubility of lithium polysulfide (Li2Sm, 2 ≤ m ≤ 8) reaction intermediates with increasing salt concentration.46 In the present study, we investigated ORR in highly concentrated Li[TFSA]/DMSO solutions. The dissolution of LiO2, the ORR intermediate, is enhanced in organic solvents such as DMSO having higher electron-pair donating ability than glymes,16 which can be attributed to greater solvation of LiO2. 16, 19-20

We first elucidated the solvation structure of Li[TFSA] in DMSO using Raman spectroscopy,

and the amount of free DMSO in the solution as a function of Li+ concentration was quantitatively evaluated. Second, the ORR was assessed using a rotating ring-disk electrode (RRDE), where the solubility of LiO2 was correlated with the concentration of free DMSO. Subsequently, discharge and ACS Paragon Plus Environment

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charge tests of Li-O2 cells were carried out using carbon nanotube (CNT) electrodes. The morphology of dominant ORR product (Li2O2) was affected by the solubility of LiO2.16, 23 The side reaction of DMSO with superoxide and Li2O2, which results in the generation of a byproduct LiOH, took place in the solution containing free DMSO, but was suppressed in the highly concentrated solutions (> 2.3 mol dm−3) owing to the relatively low solubility of LiO2.

2. EXPERIMENTAL SECTION Materials. Triglyme (water content < 50 ppm) and lithium bis(trifluoromethanesulfonyl)amide (Li[TFSA], battery grade, water content < 50 ppm) were kindly supplied by Nippon Nyukazai Co., Ltd. and Solvay Japan, respectively. Dimethyl sulfoxide (DMSO, super-dehydrated grade, water content < 20 ppm, Kanto Chemical) was used as received. The electrolytes were prepared by simple mixing of DMSO and Li[TFSA] in an Ar-filled glove box (VAC, [H2O] ~0.5 ppm). Measurements. The viscosities and densities of the electrolytes were measured using a Stabinger viscometer (SVM3000, Anton Paar). Ionic conductivity was measured using the complex impedance method in the frequency range of 500 kHz to 1 Hz with 100 mV amplitude by using a VMP3 (Bio-Logic) impedance analyzer. Two Pt-black electrodes (CG-511B, TOA Electronics) were dipped in the electrolyte solution, and the cell was thermally equilibrated at 30 °C for 1 h before conductivity measurement using a thermostat chamber (SU-261, Espec). Raman spectra of the electrolyte solutions were measured using a Raman spectrometer (RMP330, JASCO) with 532 nm laser excitation, which was calibrated using a polypropylene standard. The samples were sealed in a capillary tube and their temperature was adjusted to 30 °C using a Peltier microscope stage (TS62, INSTEC) with a temperature controller (mk1000, INSTEC). Raman spectral bands were analyzed for different concentrations using the JASCO Spectra Manager program. After baseline correction of the entire spectral range, the integrated intensity (area) in the wavenumber range of 630–730 cm−1 (including both free- and bound-DMSO) was normalized. The Raman spectra of the ACS Paragon Plus Environment

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solutions were deconvoluted into 5 peaks using a Gaussian-Lorentzian (pseudo-Voigt) function as shown in Figure S1. For the deconvolution of spectra, the peak position of the C-S symmetric and asymmetric stretching modes of free DMSO were fixed at 669 and 698 cm−1, respectively. The peak positions of the TFSA S-N stretch, and C-S symmetric and asymmetric modes of bound DMSO were not fixed. Pulsed-field gradient spin echo (PGSE) NMR measurements were carried out to evaluate the self-diffusion coefficients of the components of the electrolyte solutions. A JEOL ECX-400 NMR spectrometer with a 9.4 T narrow-bore superconducting magnet and a pulsed-field gradient probe was used for the measurements. The samples were inserted into a NMR microtube (BMS-005J, Shigemi) to a height of 3~5 mm to exclude convection. 1H, 7Li, and 19F NMR spectra were recorded for DMSO, Li+ ion, and [TFSA]− anion, respectively. The self-diffusion coefficients (D) were measured via the use of a modified Hahn spin echo-based PGSE sequence incorporating a pulsed-field gradient (PFG) in each τ period. The free diffusion echo signal attenuation (E) is related to the experimental parameters by the Stejskal equation with sinusoidal PFG: ln( E ) = ln( S S δ =0 ) =

− γ 2 g 2 Dδ 2 (4∆ − δ )

π2

(1)

where S is the spin echo signal intensity, δ is the duration of the field gradient with magnitude g, γ is the gyromagnetic ratio and ∆ is the interval between the two gradient pulses. The detailed procedures for the method are described elsewhere.47-48 Hydrodynamic voltammetry was performed to investigate the ORR/OER in the electrolytes using an air-tight 4-electrode cell equipped with water jacket. The rotating glassy carbon (GC) ring−GC disk electrode was used as the working electrode. The inner and outer radii of the GC ring electrode were 2.5 and 3.5 mm, respectively, and the radius of the GC disk electrode was 2 mm. The reference electrode was Li metal soaked in 1 mol dm−3 Li[TFSA]/triglyme, and the counter electrode was Pt coil. The reference electrode was separated from the sample solution using a liquid junction (Vycor glass). Cyclic voltammograms (CVs) were recorded using ALS700E bipotentiostat system (ALS Co., Ltd, ACS Paragon Plus Environment

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Japan). Temperature was maintained by using a water circulator at 30 °C. The electrolytes were bubbled with dry O2 gas prior to use and CV was conducted with continuous gas flow (1atm). Freestanding, binder free carbon nanotube (CNT) electrodes were prepared by a chemical vapor deposition process consisting of the catalyzed growth of multiwalled carbon nanotube arrays from Fe nanoparticle catalyst supported on an Al2O3-coated Si wafer with C2H2 as a carbon precursor.49-51 After vacuum drying CNT electrodes at 100 °C for overnight, a Li-O2 cell was fabricated with the CNT electrode (~ 1 × 1 cm, ~ 1 mg cm−2), glass filter as a separator (Whatman GF/A), Li metal foil (15 mm diameter, Chemetall), and an electrolyte in an Ar-filled glove box. A stainless steel mesh was used as the current collector for CNT electrode. After assembling, the cells were pressurized by dry O2 gas to 25 psi to supply an adequate amount of O2. Galvanostatic discharge and charge tests were conducted using a VMP3 (Biologic) galvanostat at room temperature. X-ray diffraction (XRD) measurements on discharged CNTs were conducted using a Rigaku Smartlab. The CNT electrode was extracted from the cell in the glovebox immediately after the discharge measurement (discharge time = 160 h) and sealed in an air-tight XRD sample holder (Anton Paar). The discharged CNT electrodes were observed with a scanning electron microscope (Zeiss Merlin High-resolution SEM with In-Lens detector at an accelerating voltage of 3 kV). The SEM samples were washed with dry 1,2-dimethoxyethane to remove residual electrolyte and sealed in an aluminum laminate package with heat sealer inside the glovebox for transport, before being opened and quickly placed into the SEM chamber, minimizing air exposure.

3. RESULTS AND DISCUSSION Solvation Structure of Li[TFSA]. To investigate the solvation structure of Li+ ion in the solutions, concentration-dependent Raman spectra were recorded. The molar concentrations of Li[TFSA] (cLi[TFSA]) and the molar ratios of Li[TFSA]/DMSO (cLi[TFSA]/cDMSO, where cDMSO is the molar concentration of DMSO) in the solutions are shown in Table S1. As shown in Figure 1, pure DMSO exhibits 2 bands centered at 669 and 698 cm−1, which are assigned to the C-S symmetric and ACS Paragon Plus Environment

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asymmetric stretching modes of DMSO, respectively.52 Dissolving Li[TFSA] into DMSO introduced another band at ~740 cm−1, which can be assigned to the S-N symmetric stretching vibration coupled with the CF3 bending of [TFSA]−53. The integrated intensity of this [TFSA]− peak (at ~740 cm−1) was found to increase linearly with cLi[TFSA]/cDMSO (Figure S2). This trend indicates not only that the Raman scattering coefficients of the DMSO and [TFSA]− are independent of the molar ratio, but also that the Raman scattering coefficients for free- and bound-DMSO are comparable.

Figure 1. (a) Normalized Raman spectra of Li[TFSA]/DMSO mixtures in the range from 630 to 770 cm−1 measured at 30 °C. Expanded views for the peaks of (b) C-S symmetric stretching mode and (c) CS asymmetric stretching mode of DMSO.

In the electrolytes with concentrations ≤ 2.3 mol dm−3 (cLi[TFSA]/cDMSO ≤ 1/4), the Raman peaks of DMSO corresponding to the C-S stretching modes at 669 cm−1 and 698 cm−1 had shoulders at 676 and 709cm−1, respectively. These shoulders can be assigned to the DMSO bound to the Li+ ion (denoted as bound-DMSO), as DMSO molecules present in solvated [Li(DMSO)n]+ cations have higher vibrational frequencies for the C-S stretching modes.52 Decreasing the molar ratio of DMSO to Li[TFSA] resulted in increased intensities of the bound-DMSO at 676 and 709 cm−1 and the spectra had isosbestic points at 672, 688, and 704 cm−1 for concentrations ≤ 2.3 mol dm−3 (cLi[TFSA]/cDMSO ≤ 1/4).


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We can estimate the Li+ solvation number (n) of [Li(DMSO)n]+ from the relative contributions of free DMSO at 698 cm−1 and bound-DMSO at ~710 cm−1 to the integrated intensities of the Raman C-S asymmetric stretching mode, which change linearly with the molar ratio of Li[TFSA]/DMSO in the concentration range ≤ 2.3 mol dm−3 (cLi[TFSA]/cDMSO ≤ 1/4), as shown Figure 2. Here the band of the C-S asymmetric stretching mode of DMSO at around 700 cm−1 was used for evaluating the solvation number of Li+ ion in solution, because the difference of the peak positions between the free- and boundDMSO is larger than that of the symmetric stretching mode. Assuming that the Raman scattering coefficients for free- and bound-DMSO were identical and constant in the whole concentration range (vide supra), Ib / (If + Ib) and the concentrations of Li[TFSA] (cLi[TFSA]) in the solution are related as follows: c Li[TFSA] Ib cb = =n I f + I b c DMSO c DMSO

(2)

where cDMSO is the total molar concentration of DMSO, cb is the molar concentration of bound-DMSO, and n is the Li+ solvation number of [Li(DMSO)n]+ present in the solution. In the range of cLi[TFSA]/cDMSO ≤ 1/4, the solvation number n was found to be 3.7 from the slope (red line) shown in Figure 2, forming [Li(DMSO)4]+ complexes with Li+ as monodentate ligands, which is in agreement with previous reports. Other groups have already reported that the solvation number of Li+ in DMSO solution is ca. 4.52, 54-55 Similarly, we can estimate the molar fraction of free DMSO (black line in Figure 2) in the solution from the Raman spectra using If / (If + Ib) = cf / cDMSO, where cf is the molar concentration of free DMSO (cf = cDMSO – cb). As expected, the amount of free DMSO decreased with increasing Li salt concentration, where the molar ratio of Li[TFSA] and DMSO equivalent to the expected solvation number (n = 4) of Li+ for the 2.3 mol dm-3 solution (cLi[TFSA]/cDMSO = 1/4) gave rise to a small amount of free DMSO (8.7%). In the more concentrated solutions, greater than 2.3 mol dm−3 (cLi[TFSA]/cDMSO > 1/4), free DMSO was not detectable by Raman spectroscopy with experimental certainty. In the concentration higher than 2.3 mol dm−3, [TFSA]− anions start to participate in the ACS Paragon Plus Environment

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complex formation of Li+, forming contact ion-pairs (CIPs), where the anion directly interacts with the cation, and aggregates (AGGs) such as triple ions (e.g., [Li(DMSO)x]+−[TFSA]−−[Li(DMSO)x]+) (vide infra).

Figure 2. The fraction of the integrated intensities of Raman peaks corresponding to free DMSO, If / (If + Ib), and bound-DMSO, Ib / (If + Ib), in the Li[TFSA]/DMSO solutions. If is the integrated intensity of the free DMSO peak centered at 698 cm−1, and Ib is the integrated intensity of the bound DMSO peak at ~710 cm−1. Insets show sketches of the predominant solvation structures.

While the Raman peak positions of both C-S symmetric and asymmetric stretching modes of these highly concentrated electrolytes stayed nearly constant for electrolytes with cLi[TFSA]/cDMSO equal to or less than 1/4, a linear shift was found with increasing molar fraction of Li[TFSA]) in the solution with cLi[TFSA]/cDMSO greater than 1/4, as shown in Figure 3(a). In addition, the Raman spectra of these highly concentrated electrolytes, where the free DMSO is negligible (Figure 2), no longer pass through the isosbestic points (Figure 1), which is similar to spectral changes reported for the highly concentrated solutions of Li[TFSA]/N,N-dimethylformamide,56 suggesting a stronger Li+-DMSO interaction and a different solvation structure from electrolytes having cLi[TFSA]/cDMSO ≤ 1/4.


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Figure 3. Peak positions of the (a) C-S symmetric and asymmetric stretching modes of bound-DMSO, and (b) S-N stretching mode of [TFSA]− in the Li[TFSA]/DMSO solutions. Insets show sketches of predominant solvation structures.

The peak position of the S-N symmetric stretching vibration coupled with the CF3 bending of [TFSA]− anion, which is sensitive to complex formation of [TFSA]− with alkali metal cations, is plotted as a function of molar fraction of Li[TFSA] in the solution in Figure 3(b). The peak position stayed nearly constant at 740.5 cm−1 in the concentrations ≤ 2.3 mol dm−3 (cLi[TFSA]/cDMSO ≤ 1/4). Previous findings have shown that the peak appears at 744 cm−1 with

cations

in

liquid,

e.g.,

a

pure

ionic

53

for [TFSA]− anion interacting very weakly

liquid

of

1-ethyl-3-methylimidazolium

bis

(trifluoromethanesulfonyl)amide ([EMI][TFSA]), and it shifts to a higher frequency of 750 cm−1 for the [TFSA]− anion in complexes with Li+ ions.53 Therefore, [TFSA]− is uncoordinated in the solutions with cLi[TFSA]/cDMSO ≤ 1/4. The distance between anions and cations decreases and ion-pair formation occurs with increasing concentration, and as the amount of free-DMSO decreases, the [TFSA]− is expected to ACS Paragon Plus Environment

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become closer to the solvate [Li(DMSO4)]+ cations. However, [TFSA]− is not bound to Li+ in the solutions, even for cLi[TFSA]/cDMSO = 1/4, as suggested by Raman (Figure 3b). This observation can be attributed to the high electron-pair-donating ability of DMSO and its stronger affinity towards Li+ as compared to [TFSA]−, and the formation of a solvent-separated ion pair (SSIP) in the cLi[TFSA]/cDMSO = 1/4 solution, where [TFSA]− is not involved in the first solvation shell of Li+. In further support of this claim, the Gutmann donor number (DN) of DMSO is 30, which is much higher than that of other common solvents such as propylene carbonate (DN = 15) and 1,2-dimethoxyethane (DN=24), indicative of high electron-pair-donating ability. In contrast, the DN of [TFSA]− anion is reported to be as low as 7,57 therefore, the interaction between Li+ and DMSO is much stronger than that of Li+ and [TFSA]−, resulting in scarce participation of the [TFSA]− anion in coordinating to Li+, even at the critical electrolyte composition of cLi[TFSA]/cDMSO = 1/4. For electrolytes with cLi[TFSA]/cDMSO > 1/4, the [TFSA]− peak shifts to higher frequency (stronger S-N stretching vibration) with increasing Li[TFSA] concentration, as shown in Figure 3(b), suggesting increasing bound-[TFSA]− in the solution with higher cLi[TFSA]/cDMSO.53, 58 In these highly concentrated electrolytes, the coordination of Li+ ion cannot be comprised of DMSO molecules only, and the [TFSA] − anions participate in the coordination of Li+, resulting in the formation of CIPs and AGGs. Transport Properties of Electrolytes. Figure 4(a) shows the concentration dependencies of viscosity and ionic conductivity for Li[TFSA]/DMSO solutions at 30 °C. The viscosity increases monotonically with increasing Li[TFSA] concentration, which can be attributed to stronger coulombic interaction between solvated Li+ ion ([Li(DMSO)n]+) and [TFSA]− anion associated with reduced average distance between the cation and the anion. On the other hand, the ionic conductivity increases with increasing Li[TFSA] concentration in the range of 0−1 mol dm−3 and decreases at higher than 1 mol dm−3, rendering a maximum ionic conductivity at around 1 mol dm−3. This observation can be attributed to the fact that while the number of charge carriers (ionic species) increases with increasing concentration - to raise ion conductivity - increasing concentration leads to reduced diffusion


12

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coefficients of ionic species associated with increased viscosity of the solution and increased proportion of the solvated [Li(DMSO)4]+ cation having a larger hydrodynamic radius in the solution. The selfdiffusion coefficients of DMSO, Li+ ion, and [TFSA]− anion measured by PGSE-NMR decreased monotonically with increasing Li[TFSA] concentration, as shown in Figure 4(b). The reduction in the diffusion coefficient of each species with increasing salt concentration can be attributed primarily to the increase in viscosity of the solution. At concentrations lower than 2.3 mol dm−3, the self-diffusion coefficient followed the order DDMSO > DTFSA- > DLi+, which agrees well with previous findings for conventional electrolyte solutions.48, 59-60 Although the ionic radius of Li+ ion is the smallest among the species in the solution, the solvated Li+, [Li(DMSO)4]+, has the largest hydrodynamic radius, and thus Li+ has the lowest diffusion coefficient. DDMSO is greater than DLi+, suggesting that the free DMSO molecules diffuse faster than those involved in the solvation of Li+.

Figure 4. Concentration dependencies of (a) viscosity and ionic conductivity, (b) self-diffusion coefficients (D) of Li+, [TFSA]−, and DMSO and (c) DDMSO/DLi ratios for Li[TFSA]/DMSO solutions at 30 °C. The lines in the figure are just guides for the eye.

Although we could not distinguish bound-DMSO from free-DMSO using NMR, the ratio of diffusion coefficients of Li+ and DMSO (DDMSO/DLi) is a good metric for evaluating the stability of the solvate cations in the solution.39, 48, 60-61 The concentration of free-DMSO decreased with increasing Li[TFSA] concentration, and DDMSO/DLi decreased with increasing Li[TFSA] concentration, as shown in Figure 4 (c). A similar change in the ratio of self-diffusion coefficients was observed in our earlier work ACS Paragon Plus Environment

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for mixtures of Li[TFSA] with another monodentate solvent, tetrahydrofuran.48 If the lifetime of the solvate [Li(DMSO)4]+ cation was longer than the time scale of the PGSE-NMR measurement, 10−2–10−3 s, the diffusion coefficients of DMSO and Li+ would be identical in the 2.3 mol dm−3 (cLi[TFSA]/cDMSO = 1/4) solution. However, the DDMSO/DLi is greater than unity even at the concentrations at and higher than 2.3 mol dm−3, suggesting that the lifetime of the [Li(DMSO)x]+ is shorter than 10−3 s in the highly concentrated solutions. It is hypothesized that the ligand (solvent) exchange between solvate cations becomes slower with increasing Li salt concentration, and the DDMSO/DLi is reduced. During ligand exchange, free DMSO may be produced, however, the lifetime of free-DMSO in the highly concentrated solutions should be short due to high ligand exchange kinetics because we could not detect free-DMSO at higher than 2.3 mol dm−3 concentrations (cLi[TFSA]/cDMSO > 1/4) using Raman spectroscopy (Figure 1). ORR in Li[TFSA]/DMSO Solutions. To study the effect of the concentration of Li salt on ORR, RRDE measurements were carried out in O2-saturated Li[TFSA]/DMSO solutions as a function of Li+ concentration (Figure 5) in a 4-electrode cell, with Li metal soaked in 1 mol dm−3 Li[TFSA]/triglyme solution as the reference electrode. ORR current was observed for the disk electrode in each cell during the cathodic scan, and an anodic current corresponding to the oxygen evolution reaction (OER) was observed during the anodic scan. That the ORR current has a maximum, rather than a steady state, regardless of Li+ concentration, can be attributed to the fact that ORR kinetics under these dynamic conditions are fundamentally limited not by the constant mass transport, but rather electron transfer, such as through a growing layer of insulating Li2O2.

62-63

During ORR in the 0.1 mol dm-3 (Figure 5a)

and 1 mol dm-3 Li+ (Figure 5b) electrolytes, where excess DMSO is expected, large positive ring currents are observed, consistent with convection and oxidation of soluble LiO2 formed on the disk during O2 reduction (eq 3). 16-18 Further evidence for this hypothesis comes from cyclic voltammograms measured in 1 mol dm−3 Li[TFSA]/DMSO which show that the ORR is sensitive to rotation speed (Figure S3a). That disk and ring currents during ORR increase as rotation increases suggests that the


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ORR involves soluble intermediates, consistent with the claim that at higher rotation, disk passivation is slowed, and soluble LiO2 is transported to the ring at faster rates. On the other hand, negligible ring current was observed during the disk anodic scan, likely because unlike ORR, the OER mechanism for Li2O2 oxidation does not involve soluble superoxide-based intermediates.64

Figure 5. RRDE responses on GC disk / GC ring electrode measured at 1000 rpm at 100 m Vs−1 in O2 saturated (a) 0.1 mol dm−3, (b) 1 mol dm−3, (c) 1.7 mol dm-3, (d) 2.3 mol dm-3 and (e) 3 mol dm−3 Li[TFSA]/DMSO at 30 °C. Reference electrode: Li metal in 1 mol dm-3 Li[TFSA]/triglyme. The potential of the GC disk electrode was swept from the open circuit potential (OCP) of ca. 3 VLi to the negative direction and the sweep direction was reversed at 1.5 VLi. The potential of the GC ring electrode was set to be 3.3 VLi. (f) Concentration dependency of peak ring/disk current ratio (N), data labels are fraction of free DMSO determined by Raman analysis (see text)”

As shown in Figure 5, the anodic current of the ring electrode (during the ORR at disk electrode) decreases with increasing concentration of Li[TFSA] and is absent at 3 mol dm-3 Li+. We assess the solubility of LiO2 using the fraction of ORR charge composed of soluble LiO2, which we estimate using the peak ring/disk current ratio ratio during ORR, N = ip (ring) / ip (disk), where ip (ring) and ip (disk) are the peak currents of ring and disk electrodes, respectively.65-66 Figure 5 (f) shows that as the concentration of Li[TFSA] increases from 0.1 to 3 mol dm-3, N decreases from ~0.33 to ~0. Combined ACS Paragon Plus Environment

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with our previous Raman results, the inverse correlation between peak current ratio and Li+ concentration suggests that the fraction of excess DMSO determines the solubility of LiO2. The collection efficiency of the RRDE evaluated from peak current density using soluble potassium ferricyanide/ferrocyanide redox couple (4 mmol dm−3 K4[Fe(CN)6] + 1 mol dm−3 KCl aqueous solution) is 0.53. Although the highest peak current ratio is lower than the collection efficiency of the RRDE, this is easily attributable to the fact that even at the highest LiO2 solubilities, LiO2 generated at the disk electrode can be partially converted to insoluble Li2O2 chemically and/or electrochemically before transport to the ring can occur. Indeed, Bruce and co-workers16 and Kwabi et al.8 estimate that chemical and electrochemical rates of LiO2 conversion can be as high as ~0.01 and ~1 s-1 respectively. 3 mol dm−3 Li[TFSA]/DMSO solution has a relatively high viscosity (Figure 4), and this may also suppress the transport of LiO2 from the disk to ring. To elucidate the effect of viscosity, RRDE voltammograms were measured in 3 mol dm−3 Li[TFSA]/DMSO at higher rotation speeds (Figure S3b), however, an increase of the ring current was not observed even at 3000 rpm. This suggests that the LiO2 hardly dissolves in 3 mol dm−3 Li[TFSA]/DMSO i.e. the higher viscosity is not the major factor. Note that RRDE in 1 mol dm−3 solution with higher rotation speed (Figure S3a) shows increasing both of disk and ring currents (vide supra). Given that Li+ is a hard acid, it can be reasonably assumed that at least the Li+ of LiO2 should be well-solvated for dissolution of LiO2 to occur. Raman spectroscopy measurements (vide supra) revealed that free DMSO decreased with increasing concentration of Li[TFSA] and that free DMSO is negligibly present in solutions of higher concentration than 2.3 mol dm−3 (cLi[TFSA]/cDMSO > 1/4). Thus, we hypothesize that the solubility of LiO2 decreased with increasing concentration of Li[TFSA] because amount of the free DMSO, which can contribute to the solvation of LiO2, decreased. In the 2.3 mol dm−3 solution (cLi[TFSA]/cDMSO = 1/4), however, there was a positive ring current despite the negligible amount of free DMSO in the solution determined by Raman measurements. A possible reason is that at 2.3 mol dm-3, while the average ratio of DMSO to Li+ is 4:1, there are local variations in DMSO:Li+ ratio that result in greater than expected LiO2 solubility. This


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could occur if one dissolution mechanism is that the LiO2 strips DMSO molecules from the neighboring solvate [Li(DMSO)4]+ cations, and the resulting [Li(DMSO)x]O2 complex dissolves and CIPs of [Li(DMSO)y]+−[TFSA]− are formed simultaneously. As we move into more and more concentrated solutions, the ability of the solution to maintain coordinative saturation of the lithium through transfer of DMSO becomes less and less, and hence the solubility of LiO2 decreases. Indeed, the Raman spectra of Li[TFSA]/DMSO solutions with higher concentration than 2.3 mol dm−3 (cLi[TFSA]/cDMSO > 1/4) suggested the formation of CIPs (Figure 3), and the PGSE-NMR data indicated that the solvate structure of [Li(DMSO)4]+ changes dynamically in the solution where cLi[TFSA]/cDMSO = 1/4 (Figure 4). Above 2.3 mol dm-3, though, RRDE data suggest that the solubility of LiO2 is almost negligible i.e. in the extremely concentrated 3 mol dm−3 solution. Indeed, the cathodic current of ORR at the disk electrode in 3 mol dm−3 solution is almost independent of the rotation speed (Figure S3b), indicating that the ORR product adheres to the electrode surface and is further reduced electrochemically and/or disproportionates chemically to produce Li2O2 on the surface of the electrode. Indeed, Li2O2 formation in a Li-O2 cell with 3 mol dm−3 Li[TFSA]/DMSO was confirmed by XRD measurement (vide infra). In further support for the claim that higher Li+ concentrations decreases LiO2 solubility, we note that the coulombic efficiency (defined as area (electric charge) of anodic peak divided by that of cathodic peak, assuming pure Li2O2 formation and degradation) of OER/ORR at the GC disk electrode becomes higher with increasing concentration of Li[TFSA] (black lines Figure 5a, b, e), being 22, 32, and 58 % in the 0.1, 1, and 3 mol dm−3 solutions, respectively. It should be noted that although there is a possibility of formation of minor byproducts such as Li2CO3 during ORR and then subsequent decomposition above 4 V during charge,50, 67 it is reasonable to assume that the main contributor to the observed trend in coulombic efficiencies is the difference of LiO2 solubility. As the solubility of LiO2 decreases, its diffusivity to the bulk electrolyte is suppressed, and disproportionation/further electrochemical conversion to Li2O2 is more likely to occur at the electrode surface, resulting in higher coulombic efficiency. For the LiO2-insoluble electrolyte (3 mol dm−3 solution), for which the ring


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current was negligible, we can reasonably assume that all superoxide formed at the electrode surface was converted to Li2O2 with no loss to solution. In this case, the coulombic efficiency for the ORR/OER at the GC disk electrode was the highest observed, but still less than 60%. This suggests that some irreversible processes occurred during the ORR and/or OER reactions in the highly concentrated solution.67 One possibility is that a side reaction of LiO2 with the electrolyte and/or the GC electrode took place during the ORR, and that electrically resistive species such as Li2CO3 were produced on the electrode. 9, 20, 22, 50 Li-O2 Battery. Figure 6 (a) shows the initial discharge and charge curves of Li-O2 cells with 1 and 3 mol dm−3 Li[TFSA]/DMSO. The battery tests were carried out galvanostatically at a current density of 25 mAg−1-CNT with a cut-off condition of maximum capacity of 4000 mAhg−1 based on the mass of CNT. The discharging and charging overpotentials of the cell with 3 mol dm−3 solution were slightly higher than those of the cell with 1 mol dm−3 solution. This is attributed to the higher solution resistance and slower ORR/OER kinetics in the 3 mol dm−3 solution. Both of the cells showed a relatively flat voltage plateau at around 2.7 V owing to the ORR during the discharging. During charging, ca. 2500 mA h g−1 of the discharge capacity was recovered at lower than 4 V, before the voltages of the cells rose to higher than 4 V for the remaining ca. 1500 mA h g−1. The charging capacity observed at lower than 4 V could be mainly attributed to the OER of Li2O2.68-70 The voltage rising to higher than 4 V indicated that substances other than Li2O2 were involved in the charging reaction. The charging capacities of the cells at lower than 4 V were similar, and this suggested that the amounts of the discharge product Li2O2 accumulated in the CNT electrodes were almost identical, regardless of differences in the solubility of LiO2. The solubility of LiO2 in 1 mol dm−3 Li[TFSA]/DMSO solution is relatively high, and diffusion of LiO2 from the CNT surface to the electrolyte would be expected to take place. Presumably, the disproportionation reaction of LiO2 mainly occurred in the electrolyte contained in the cathode layer, i.e., diffusion layer consisting of CNTs, and the discharge product was mainly accumulated within the cathode layer. ACS Paragon Plus Environment

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Figure 6. (a) Galvanostatic charge-discharge curves of [Li metal|Li[TFSA]/DMSO|CNT/O2] cells measured at a current density of 25 mA g−1-CNT. (b) XRD patterns of CNT electrodes discharged at 25 mA g−1 for 4000 mAh g−1-CNT in Li[TFSA]/DMSO electrolytes. The CNT electrode was removed from the cell immediately after the discharge (discharge time = 160 h). Al peaks originate from the XRD measurement cell.

Figure 7 shows the SEM images of CNT electrodes after discharge to a capacity of 4000 mA h g−1 in Li[TFSA]/DMSO solutions. Disk-like particles of 100−200 nm are clearly observed in the CNT electrode discharged in 1 mol dm−3 Li[TFSA]/DMSO (Figure 7 (a)), which is reminiscent of Li2O2 from previous reports.9,

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XRD patterns of the CNT electrodes after the discharge in 1 mol dm−3

solution are shown in Figure 6 (b), with diffraction peaks assignable to Li2O2 appearing at 2θ = 32.8, 34.8 and 58.5°, in addition to the peaks of Li2O2, the peaks of LiOH, which is known as a byproduct of ORR in DMSO-based solutions,29 are also observed at 2θ =35.7 and 51.4 ° for the CNT electrode ACS Paragon Plus Environment

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discharged in 1 mol dm−3 solution. The side reaction during the ORR will be discussed later. The disklike particles of Li2O2 were also generated during the discharge process in the highly concentrated solution of 3 mol dm−3 as shown in Figure 7 (b). However, the particle size is less than 100 nm and is smaller than that generated in 1 mol dm−3 solution (Figure 7 (a)). This is probably because the nucleation rate of Li2O2 increases with decreasing solubility of LiO2 and/or the formation of Li2O2 proceeds via a surface pathway inhighly concentrated solution.16, 21 Several reports have shown that the surface pathway is prone to form a film of Li2O2 rather than large, toroidal Li2O2.16, 23 Under this scheme, the slightly higher overpotentials for the ORR and OER in the Li-O2 cell with 3 mol dm−3 Li[TFSA]/DMSO solution (Figure 6 (a)) can be explained as a consequence of higher Li2O2 nucleation rates, which decreases the number of free sites for ORR/OER on the electrode.16, 18

Figure 7. SEM images of CNT electrodes discharged at 25 mA g−1 for 4000 mAh g−1-CNT in (a, c) 1 mol dm−3 and (b) 3 mol dm−3 Li[TFSA]/DMSO. The CNT electrode was removed from the cell immediately after the discharge (discharge time = 160 h). As shown in Figure 6 (b), the byproduct LiOH was formed in 1 mol dm−3 solution, however, LiOH was not detected for the electrode discharged in 3 mol dm−3 solution. In fact, Kwabi et al. have reported the reaction of DMSO with Li2O2 and superoxide species.29 The superoxide species accelerates the decomposition of DMSO and conversion of Li2O2 to LiOH, which has a flake-like morphology. Indeed, in the case of electrolyte solution containing excess DMSO (1 mol dm−3), in addition to the disk-like particles (Figure 7(a)), flake-like particles were observed for the electrode discharged in 1 mol dm−3 Li[TFSA]/DMSO (Figure 7 (c)). We believe that O2•– radical species dissolved in the solution attack free DMSO and the formation of LiOH proceeds. The flake-like particles were not found for the electrode discharged in 3 mol dm−3 solution. This is because the solubility of LiO2 is very low in 3 mol ACS Paragon Plus Environment

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dm−3 solution, and would be expected to quickly disproportionate/be further reduced on the electrode surface, as shown by RRDE data (Figure 5). Moreover, any trace amounts of dissolved LiO2 are unlikely to attack DMSO as negligible free DMSO exists at that concentration. Since the decomposition of DMSO and the conversion of Li2O2 to LiOH are rather slow in the absence of superoxide species,29 side reactivity between DMSO and Li2O2 is suppressed in the highly concentrated solution. The byproducts of ORR except for LiOH were not detected by XRD, however, the low coulombic efficiency of OER/ORR on the GC disk electrode in the highly concentrated solution indicated other irreversible side reactions proceeded on the carbon electrode (Figure 5 (e)). Although we could not detect Li2CO3, which is a possible byproduct due to the side reaction between LiO2 and carbon electrode,

9, 20, 22, 50

it might be generated on the CNT electrode. Previously, it was reported that

the Li2CO3 formed on a carbon electrode during ORR causes the voltage to rise to higher than 4 V during charging of a Li-O2 cell. 50 Therefore, the voltage plateau at higher than 4 V observed for the cell with 3 mol dm−3 solution (Figure 6 (a)) might originate from the decomposition of Li2CO3 on the CNT.

4. CONCLUSIONS The solvation structure of Li[TFSA] dissolved in DMSO was analyzed using Raman spectroscopy.

Li+ is solvated by DMSO and forms solvate [Li(DMSO)4]+ in the solutions at

concentrations lower than 2.3 mol dm−3, and free DMSO decreases with increasing concentration of Li[TFSA]. At concentrations higher than 2.3 mol dm−3, where the molar ratio of Li[TFSA]/DMSO is higher than 1/4, all of the DMSO molecules are involved in the solvation of Li+, and free DMSO is hardly present in the solution. In highly concentrated solutions, [TFSA]− also participates in the coordination of Li+. PGSE-NMR data suggested that the solvation structure of Li+ in the highly concentrated solution is dynamic, and that the solvate [Li(DMSO)x]+ ions exchange DMSO molecules at a high rate (within the timescale of 10−3 s). ACS Paragon Plus Environment

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Varying concentrations of free DMSO in Li[TFSA]/DMSO have profound effects on the ORR. RRDE data demonstrated the solubility of LiO2, which is an intermediate of ORR, decreases with increasing Li[TFSA] concentration (with decreasing free DMSO), and LiO2 is almost insoluble in the extremely concentrated solution of 3 mol dm−3 Li[TFSA]/DMSO because free DMSO, which can solubilize LiO2, is hardly present in the solution. Decreasing solubility of LiO2 increases the rate of deposition of Li2O2 on the electrode surface, leading to an increase in the coulombic efficiency (assuming pure Li2O2 formation and decomposition) of OER/ORR at the rotating disk electrode. Li-O2 cell tests with CNT-based positive electrodes showed that the particle size of Li2O2 decreased with increasing Li[TFSA] concentration, suggesting that Li2O2 nucleation and deposition processes are affected by the solubility of LiO2. Lastly, the solubility of LiO2 also had significant effects on the side reactions in a Li-O2 cell. With electrolytes containing a large amount of free DMSO (e.g., 1 mol dm−3 Li[TFSA]/DMSO), the solubility of LiO2 is relatively high, and the side reactions between DMSO and superoxide proceeded easily, resulting in the formation of LiOH in addition to Li2O2 upon discharge. On the other hand, in the highly concentrated solution (e.g., 3 mol dm−3 Li[TFSA]/DMSO), LiO2 is insoluble, and the side reactions which generate LiOH were suppressed. In summary, this study demonstrated that the concentration of Li salt in the electrolyte solution can have a strong effect on the solubility of the LiO2 intermediate, which in turn strongly affects parasitic reactivity, as well as Li2O2 nucleation and growth processes during the ORR.

ASSOCIATED CONTENT Supporting Information. Molar concentration, density, viscosity, ionic conductivity, and self-diffusion coefficients of Li[TFSA]/DMSO solutions (Table S1), Deconvolution of a Raman spectrum (Figure S1), Integral


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intensity of Raman peak of [TFSA]− anion (Ianion) (Figure S2), and Hydrodynamic voltammograms (Figure S3). This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION Corresponding Authors * [email protected] (K. D.) * [email protected] (Y. S.-H) Notes The authors declare no competing financial interest.

ACKNOWLEDGMENTS This study was supported in part by the Research & Development Initiative for Scientific Innovation of New Generation Batteries (RISING) program from the New Energy and Industrial Technology Development Organization (NEDO) of Japan and the JSPS KAKENHI (Nos. 14J00165 for RT, 15H03874, 16H06368 for KD) from the Japan Society for the Promotion of Science (JSPS). The work conducted at MIT was partly supported by Toyota Motor Europe. D.K. and Y.S.H. thank the SkoltehMIT Centre for Electrochemical Energy Storage. The authors thank Dr Koffi P. C. Yao and Dr HaoHsun Chang for their helpful discussion and technical assistance with the experiments.


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