Enhanced Lithium Oxygen Battery Using a Glyme Electrolyte and

Research Article Cite This: ACS Appl. Mater. Interfaces 2018, 10, 16367−16375

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Enhanced Lithium Oxygen Battery Using a Glyme Electrolyte and Carbon Nanotubes Lorenzo Carbone,† Paolo Tomislav Moro,‡ Mallory Gobet,∥ Stephen Munoz,∥,# Matthew Devany,⊥ Steven G. Greenbaum,*,∥ and Jusef Hassoun*,‡,§ †

Chemistry Department, Sapienza University of Rome, Piazzale Aldo Moro, 5, 00185 Rome, Italy Department of Chemical and Pharmaceutical Sciences and §National Interuniversity Consortium of Materials Science and Technology (INSTM), University of Ferrara Research Unit, University of Ferrara, Via Fossato di Mortara, 17, 44121 Ferrara, Italy ∥ Department of Physics & Astronomy and ⊥Department of Chemistry and Biochemistry, Hunter College of the City University of New York, New York, New York 10065, United States # Ph.D. Program in Physics, City University of New York, New York, New York 10016, United States ‡

S Supporting Information *

ABSTRACT: The lithium oxygen battery has a theoretical energy density potentially meeting the challenging requirements of electric vehicles. However, safety concerns and short lifespan hinder its application in practical systems. In this work, we show a cell configuration, including a multiwalled carbon nanotube electrode and a low flammability glyme electrolyte, capable of hundreds of cycles without signs of decay. Nuclear magnetic resonance and electrochemical tests confirm the suitability of the electrolyte in a practical battery, whereas morphological and structural aspects revealed by electron microscopy and X-ray diffraction demonstrate the reversible formation and dissolution of lithium peroxide during the electrochemical process. The enhanced cycle life of the cell and the high safety of the electrolyte suggest the lithium oxygen battery herein reported as a viable system for the next generation of high-energy applications. KEYWORDS: lithium oxygen, MWCNTs, glyme electrolyte, reaction mechanism, long cycle life

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INTRODUCTION Recent advances in lithium ion batteries and vehicle manufacturing technology have increased the driving range of full electric cars (EVs) to 200 km per charge,1 whereas a new commercial EV from Tesla, Inc. can reach a range approaching 400 km.2 The latter number appears presently to be the limit ascribed to the state-of-the-art lithium ion battery, which is composed of a carbon-based anode, a lithium transition-metal oxide cathode, and an organic carbonate-based electrolyte.3 Further increase in the driving range may be achieved by switching to a completely different chemistry, such as that associated with lithium sulfur or lithium oxygen batteries.4,5 Lithium sulfur battery, with a theoretical energy density of about 3500 Wh kg−1, is a strong candidate to reach this challenging goal,6 whereas the lithium oxygen battery is indicated to be even more energetic.7 However, the Li/O2 cell is still far from application due to several problems affecting its safety and cycle life, and an electrochemical process still not fully clarified.8,9 The relevant safety issues are mainly associated with the use of lithium metal in the cell, and reactivity and flammability of the organic electrolytes commonly employed in the lithium ion battery.10 Furthermore, cycle life and reaction © 2018 American Chemical Society

mechanism are intimately connected and strongly associated with the electrode/electrolyte interface,11 electrode nature,12 and the operating condition of the cell, including the cell geometry, the use of new separators, and protections at the lithium metal side.13 The lithium oxygen electrochemical process involves the formation of lithium peroxide (Li2O2), reactive intermediates such as lithium superoxide (LiO2), and radicals,14,15 with a stability mainly determined by the electrolyte, the cathode support16,17 for the electrochemical reaction, and the current rate18 applied to the lithium oxygen cell.19 In particular, different types of aprotic organic solvents, such as ethers or glyme, sulfoxides, phosphates, nitriles, and ionic liquids, actually influence the formation of discharge products in Li/O2 batteries.20 Literature data focusing on X-ray diffraction (XRD), X-ray photoelectron spectroscopy, and gas chromatography−mass spectrometry measurements indicated that the glyme solvents lead to significant formation of Li2O2,20 whereas other papers have suggested the addition of ionic Received: January 13, 2018 Accepted: April 20, 2018 Published: April 20, 2018 16367

DOI: 10.1021/acsami.7b19544 ACS Appl. Mater. Interfaces 2018, 10, 16367−16375

Research Article

ACS Applied Materials & Interfaces

ppm. The water content of the solvent was measured by Karl Fischer titration instrument (831 Karl Fisher Coulometer, Metrohm). The LiTFSI salt was dried under vacuum at 80 °C for 24 h. The electrolyte was formed in the glove box by dissolving LiTFSI in TEGDME with a ratio of 1 mole of salt to 1 kg of solvent (named TEGDME−LiTFSI 1 mol kg−1). Electrolyte Characterization. The thermal properties were evaluated by differential scanning calorimetry (DSC) using the Mattler Toledo DSC821 instrument. The TEGDME−LiTFSI electrolyte was sealed in an aluminum crucible under argon atmosphere before the DSC analysis. The measurement was performed by cooling down from −30 to −90 °C at the rate of −5 °C min−1. The lithium stripping deposition test was performed at a constant current rate of 0.1 mA cm−2 and a step limit of 1 h using a symmetrical lithium/lithium 2032 coin cell and a MACCOR 4000 series Battery Test System. The lithium/electrolyte interface resistance was measured using the latter cell configuration by electrochemical impedance spectroscopy (EIS) in the frequency range of 0.1 MHz to 0.1 Hz with a signal amplitude of 10 mV with a VersaSTAT MC Princeton Applied Research-AMETEK potentiostat. The impedance spectra were associated with the R(RQ)nQ equivalent circuit, where R is the resistance and Q is a constant phase element, and analyzed using the nonlinear least square fit by Boukamp Software;29,30 only the results with a χ2 lower than 10−4 were accepted. The voltammetry tests were carried out in Swedglock T-cell with lithium metal as the reference and counter electrode and a working electrode formed by depositing Super P (Timcal) on Cu or Al support for the cathodic and the anodic scan, respectively. Cyclic voltammetry (CV) for the cathodic region was performed in the potential range 0−2 V versus Li/Li+, whereas linear sweep voltammetry (LSV) for the anodic region was performed up to 5 V versus Li/Li+, both with a scan rate of 0.1 mV s−1. The measurements were carried out by a VersaSTAT MC Princeton Applied Research-AMETEK potentiostat. The self-diffusion coefficient of the 1H, 7Li, and 19F nuclei were measured by a Bruker 400 Avance III NMR spectrometer suppressing the convection effect by a pulsed-field gradient (PFG) doublestimulated echo sequence.31,32 The data were collected every 10 °C from 20 to 80 °C with a gradient strength g of 1−45 G cm−1, gradient pulse of 1−2 ms, and a diffusion delay Δ of 100−200 ms. The error on the self-diffusion coefficient is about 3−5%.31 The apparent lithium transport number (t+) was calculated by eq 133

liquid to the glyme for improving the electrolyte conductivity and enhancing the kinetics of both the O2 reduction and the lithium peroxide oxidation.21 The dominant side reactions at the oxygen electrode have been indicated to involve electrolyte decomposition with formation of Li2CO3 and lithium carboxylates, depending on the interplay between the electrode and the electrolyte.22 Furthermore, a fundamental study of the oxygen reduction reaction in nonaqueous electrolytes conducted in four solvents differing by the donor numbers (DNs), i.e., dimethyl sulfoxide (DMSO), acetonitrile (MeCN), dimethoxyethane (DME), and tetra ethylene glycol dimethyl ether (TEGDME), indicated that the solvent and the supporting electrolyte cations in the solution act in concert to influence the nature of reduction products and reversibility.23 The study concluded that high DN solvents provide increased stability for the complex [Li+(solvent)n···O2−] to the modulated Lewis acidity of the hard acid, leading to a distinct O2/O2− reversible couple, whereas low DN solvents lead to fast O2− decomposition or electrochemical reduction to O22−.23 The formation of this strongly nucleophilic specie decomposes the common organic carbonate-based electrolytes (characterized by the electrophilic carbonyl) during the discharge processes, causing fast capacity fading until cell failure,24 whereas endcapped glyme, based on ether bond with chemical formula CH3−(OCH2CH2)n−OCH3, appears more suitable with promising properties in terms of stability, ionic conductivity,25 and solvation ability for lithium salt dissociation.26 The glyme electrolyte stability may be further improved by including ionic liquids20 or suitable additives23 in the electrolyte for enhancing the solid electrolyte interface (SEI) characteristics. Ether- and glyme-based electrolytes,27 and high surface area carbons16 at the cathode side have demonstrated the most promising characteristics for efficient lithium oxygen batteries, with satisfactory cycle life, capacity higher than 500 mAh g−1 with respect to the carbon mass, and an operating voltage of about 2.8 V during the discharge process.28 In this work, we describe an enhanced system, formed by the combination of a safe electrolyte solution of lithium bistrifluoromethanesulfonimidate (CF3SO2NLiSO2CF3, LiTFSI) salt in tetra ethylene glycol dimethyl ether solvent (TEGDME) and multiwalled carbon nanotubes (MWCNTs), as the reaction host for the lithium oxygen reaction. The cell shows an outstanding cycle life, extended up to 250 cycles, due to the optimal characteristics of the electrolyte as revealed by NMR and electrochemistry. Reversible cell behavior involving formation and subsequent dissolution of Li2O2 is clearly detected by XRD, scanning electron microscopy (SEM), and transmission electron microscopy (TEM) on the MWCNT surface. Although further efforts are required for practical employment, the cell herein studied is proposed as a valid candidate for powering next-generation EVs with an extended driving range.

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t+ =

DLi DLi + DF

(1)

where DLi is the lithium self-diffusion coefficient and DF is the fluorine self-diffusion coefficient of the TFSI− anion. The apparent ionic conductivity of the electrolyte (δNMR) was calculated from the selfdiffusion coefficient by the Nernst−Einstein equation eq 234

δ NMR =

F 2[C] (DLi + DF) RT

(2)

where F is the Faraday constant (96 485 coulomb), [C] represents the solution concentration (mol cm−3), and R is the ideal-gas constant (8.314472 J K−1 mol−1). Furthermore, the ionic conductivity was measured by impedance spectroscopy (δEIS) in a symmetrical SS/SS 2032 coin cell with a Teflon O-ring separator in the temperature range of 20−80 °C, using a signal amplitude of 10 mV in the frequency range of 1 MHz to 0.1 Hz by employing a VersaSTAT MC Princeton Applied ResearchAMETEK potentiostat. Conductivity values δEIS and δNMR were combined in eq 3 to obtain the ionic association degree (α)35

EXPERIMENTAL SECTION

Electrolyte Preparation. The electrolyte was prepared by mixing tetra ethylene glycol dimethyl ether (CH3O(CH2CH2O)4CH3), hereafter named TEGDME, and lithium bistrifluoromethanesulfonimidate (CF3SO2NLiSO2CF3), herein named LiTFSI. Both chemicals were purchased from Sigma-Aldrich. Molecular sieves of 3 Å from Sigma-Aldrich were dried under vacuum at 250 °C for 48 h and used to reduce the water content of the TEGDME below 10 ppm. After cooling back to the room temperature, the dry molecular sieves were inserted in the TEGDME solvent, and hold for at least 3 days in an argon-filled glove box with water and oxygen content lower than 1

⎛ δ ⎞ α = ⎜1 − EIS ⎟ δ NMR ⎠ ⎝

(3)

Cathode Preparation and Cell Assembly. Multiwalled carbon nanotubes from Sigma-Aldrich, hereafter named MWCNTs, were annealed under nitrogen flow at 750 °C for 12 h. The active material (named N2@MWCNTs) was mixed separately with poly(vinylidene 16368


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Figure 1. Flammability test of the TEGDME−LiTFSI.

Figure 2. Characteristics of the TEGDME−LiTFSI electrolyte (1 mol kg−1). (A) Differential scanning calorimetry (DSC) performed between −30 and −90 °C under argon. (B) Lithium stripping/deposition voltage profile of a symmetrical lithium/lithium cell performed with a constant current of 0.1 mA cm−2 and a step limit of 1 h. (C) Time evolution of the lithium/electrolyte interface resistance, and in inset the corresponding Nyquist plots, investigated by impedance spectroscopy in symmetrical lithium with a signal amplitude of 10 mV in a 0.1 MHz to 0.1 Hz frequency range. (D) Electrochemical stability window of the electrolyte determined by linear sweep voltammetry (LSV) up to 5 V in the anodic region and by cyclic voltammetry (CV) between 0 and 2 V versus Li+/Li in the cathodic region (inset), using a Swagelok T-cell with lithium as the counter and reference electrode and Super P as the working electrode, at a scan rate of 0.1 mV s−1. difluoride) (Solvay) binder in a 80:20 mass ratio and dispersed in Nmethylpyrrolidone (Aldrich). The slurry was cast on a gas diffusion layer (TGP-H-030 carbon paper, Toray) by a Doctor blade. The active material ratio on the current collector was about 0.8 mg cm−2. The electrode foils were punched in disks, dried overnight at 100 °C, and finally introduced in an argon-filled glovebox with oxygen and moisture contents lower than 1 ppm. Lithium oxygen cells were then assembled by combining lithium metal, the N2@MWCNTs, and the TEGDME−LiTFSI in a top-meshed 2032 coin cell inserted into a glass tube filled subsequently by pure oxygen. Cathode and Cell Characterization. The powder structure was characterized by X-ray diffraction using a Bruker instrument equipped with a Cu Kα source and a graphite monochromator, in the 2θ/θ scanning mode. The morphology of the samples was evaluated by SEM (Zeiss EVO 40) with a thermionic electron gun equipped with LaB6 crystal. The lithium oxygen cell was studied by a galvanostatic test using a voltage cutoff performed at a constant current of 100 mA g−1 in the 1.5−4.5 V range. Galvanostatic cycling tests were also carried out at a current rate of 100 mA g−1 and limiting the capacity to 500, 1000, and 2000 mAh g−1. The current density and specific

capacity were referred to the N2@MWCNTs weight into the electrode, whereas the surface capacity was determined with respect to the electrode geometric area.

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RESULTS AND DISCUSSION

Several studies have focused on the characterization of glymebased electrolytes using various salts, concentrations, and experimental setups.18,36−38 In this work, we shed light on the performance of a TEGDME−LiTFSI electrolyte (1 mol kg−1)39−41 by combining calorimetry, electrochemistry, and NMR. The electrolyte concentration adopted in this work (i.e., 1 mol kg−1 salt to solvent ratio) was optimized by our group in a previous comparative study on the role of lithium salt in determining the performance of lithium oxygen batteries,39 whereas the combination of TEGDME and LiTFSI was suggested as an efficient electrolyte by our group as well as by other studies.40,41 16369


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Figure 3. Species mobility in the TEGDME−LiTFSI electrolyte (1 mol kg−1). (A) Self-diffusion coefficients of the 1H, 7Li, and 19F nuclei carried out by NMR. (B) Comparison of the conductivity values calculated from the self-diffusion coefficients and determined by electrochemical impedance spectroscopy (EIS). (C) Ionic association degree of the lithium salt determined by combining NMR and EIS data. (D) Lithium transport number determined using the NMR data.Temperature range 20−80 °C. All of the equations used for the determination of the parameters of the figure are reported in the Experimental Section.

A direct ignition test, performed to evaluate the flammability of the TEGDME−LiTFSI electrolyte, reveals the absence of fire evolution upon direct exposure to a butane flame prolonged over 20 s (Figure 1). This behavior suggests the low flammability of the electrolyte under the experimental setup adopted in our work, and its possible employment in a relatively safe lithium oxygen battery. However, further experiments with a more suitable setup, such as a direct test in a cell or more effective exposure to aggressive conditions, are certainly required to full-proof the safety level of the electrolyte. The DSC plot of Figure 2A indicates that the electrolyte freezes below −50 °C, thus allowing use at relatively low temperatures. The lithium stripping/deposition profile in a symmetrical Li cell in Figure 2B evidences a very low polarization (limited to about 20 mV during the whole test), demonstrating the stability of the electrolyte. Further proof of the electrolyte stability is shown in Figure 2C on the lithium/ electrolyte interface impedance trend and, in the inset, the corresponding Nyquist plots, as well as by the linear sweep voltammetry of the Li/electrolyte/SP cell in the anodic region (Figure 2D) and the CV of the cell in the cathodic region between 0 and 2 V (Figure 2D inset). Figure 2C reveals an impedance ranging from 150 to 300 Ω due to the solid electrolyte interphase (SEI) film formation.42 Figure 2D shows an electrochemical stability window extending from 0 to 4.7 V

versus Li/Li+, that is, a region in which only SEI film formation at about 0.8 V,43 reversible insertion of lithium into amorphous carbon at about 1 V,44,45 and lithium deposition/stripping around 0 V46 can be observed (see CV in the inset). Pulsed-field gradient (PFG) NMR data reported in Figure 3A show the self-diffusion coefficients of the 1H, 7Li, and 19F nuclei (tracking the TEGDME solvent, the lithium cation, and the TFSI anion, respectively) in the temperature range between 20 and 80 °C. The data indicate faster mobility of the solvent molecules compared to the ions. The NMR data allow the determination of the apparent electrolyte conductivity, which is compared in Figure 3B with the conductivity determined by impedance spectroscopy, to find the ionic association degree (Figure 3C). The lithium transference number (Figure 3D) between 20 and 80 °C can also be calculated (see the Experimental Section for the details).47,48 Increasing the temperature increases both conductivity and ionic association degree, from 3 × 10−3 to 10−2 S cm−1 (Figure 3B) and from 0.40 to 0.55 (Figure 3C), respectively. The apparent lithium transport number appears to be only slightly affected by the temperature increase, with a value of about 0.45 over the whole temperature range despite the increase in ion association at elevated temperature (Figure 3D). It should be mentioned that an apparent transport number close to 0.5 is partly a consequence of ion association.26,49 In summary, the 16370


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Figure 4. (A) XRD patterns of the bare multiwalled carbon nanotubes (MWCNTs) and of the same material after the annealing process under nitrogen atmosphere (N2@MWCNTs). (B, E) SEM images of the bare (E, bottom) and N2-annealed (B, top) MWCNTs. (C, D, F, G) TEM images at various magnifications of the bare (F, G, bottom) and the N2-annealed (C, D, top) MWCNTs.

Figure 5. (A) Galvanostatic voltage profile of the lithium oxygen battery using the TEGDME−LiTFSI (1 mol kg−1) electrolyte and the N2@ MWCNT electrode in a 2032 top-meshed coin cell. Bottom x-axis reports the capacity with respect to the N2@MWCNTs weight while top x-axis shows the capacity with respect to the geometric area of the electrode. Current rate of 100 mA g−1 with respect to the N2@MWCNTs weight. Voltage limit of 1.5−4.5 V. (B) XRD pattern of the cell at different states of the charge, marked by T1 (OCV), T2 (fully discharged), and T3 (fully charged). SEM images of the N2@MWCNT electrode at the OCV (C), at the discharged (D, E) state, and at the charged (F) state.

functional groups.53,54 The XRD patterns of Figure 4A reveal the graphitic nature of the MWCNTs and show only minor structural changes due to the thermal treatment. More relevant effects of annealing are observed for the MWCNT morphology, as revealed by SEM (Figure 4B,E) and TEM (C, D, F, and G) at various magnifications. The images, in particular the TEM, suggest that annealing is effective in removing impurities from the nanotubes (Figure 4F,C), increasing the electrode surface (see also low-magnification SEM in Figure S1 in Supporting Information (SI)), and enabling cavities and free spaces (Figure 4G,D), which may act as the host and the catalyst for the Li/O2 electrochemical process.22,55 The TEGDME−LiTFSI electrolyte and the nitrogenannealed MWCNT electrode (N2@MWCNT) were combined and studied in a lithium oxygen cell. The cell was fully discharged and charged, the voltage profile was collected (Figure 5A), and the reaction products were studied in terms of

TEGDME−LiTFSI electrolyte shows an ionic conductivity higher than 3 × 10−3 S cm−1 in the whole range of temperature, stable overvoltage upon cycling of 0.02 V, low lithium/ electrolyte interface resistance, and wide electrochemical stability window extended up to 4.5 V, which are characteristics comparable to those of the carbonate electrolytes commonly used in lithium ion batteries.50,51 In addition, bonuses including the relatively low volatility, hence the low flammability,52 and the modest reactivity with the nucleophile products compared to other electrolytes, such as those based on DMSO, MeCN, and DME,23−26 make this electrolyte particularly suitable for lithium oxygen batteries. Therefore, the solution is subsequently used in a Li/O2 cell employing multiwalled carbon nanotubes (MWCNTs) at the cathode side. Prior to use as the electrode support for the Li/O2 cell, the MWCNTs were treated under nitrogen atmosphere at 750 °C (see the Experimental Section) to remove possible undesired 16371


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ACS Applied Materials & Interfaces structure by ex situ XRD (Figure 5B) and morphology by ex situ SEM (Figure 5C−F). The voltage profile of Figure 5A shows that the N2@MWCNT electrode reaches a capacity of about 16 Ah g−1 with respect to the N2@MWCNTs weight (a value of about 9 mA cm−2 as referred to the electrode geometric area; see Figure S2 in SI). It is worth mentioning that the cell was cycled using the voltage cutoff of 1.5−4.5 V, which is a cycling mode aimed to detect the full charge discharge products instead of optimizing cell stability, as confirmed by the capacity decay over the course of the test (Figure S3 in the SI). Lithium oxygen cell stability may be achievable only by limiting the discharge capacity during cycling to avoid a decrease of electrode conductivity by an excessive Li2O2 accumulation at the electrode surface, as reported by previous works56 and fully confirmed hereafter. The XRD patterns of the N2@MWCNT electrode collected from the lithium oxygen cells at different states of charge reported in Figure 5B at the OCV (marked by T1 in Figure 5A), fully discharged state (marked by T2 in Figure 5A), and fully charged state (marked by T3 in Figure 5A) clearly show the formation and complete dissolution of lithium peroxide during the galvanostatic test, without any residual product. This is in line with the high coulombic efficiency, i.e., approaching 100%, demonstrated by the galvanostatic test of Figure 5A. Furthermore, the corresponding SEM images (Figure 5C−F) show the formation of particles of a size of about 50−100 μm (Figure 5D,E) and suggest an electrodeposition mechanism of Li2O2 at the N2@MWCNT surface rather than the solution mechanism involving the formation of the reaction products at the electrode/electrolyte interface, which is instead observed in cells using ionic liquid electrolytes.57−66 According to the surface mechanism, the nucleation and growth of the Li2O2 crystallites evolve over the sites of the carbon support surface by a direct electrodeposition, which is strongly affected by the current rate applied to the lithium oxygen cell, leading generally to a micrometric nonuniform deposit.57−62 Instead, solution mechanism occurs when the OOR proceeds within the solution at the electrode/ electrolyte interface, rather than at the electrode surface, which may cause the precipitation of Li2O2 crystallites when the concentration of this specie exceeds the solubility limit of the electrolyte solvent, and consequent formation of small particles with a homogeneous distribution over the electrode.63−66 Therefore, the micrometric particle size of the lithium peroxide and its distribution at the electrode surface observed in Figure 5D,E suggest for our cell a surface mechanism. The SEM shown in Figure 5F evidences a complete dissolution of the products upon full charge, thus further suggesting the reversibility of the electrochemical process. The lithium oxygen battery was then cycled using a capacity limited to 500 mAh g−1 with respect to the MWCNTs weight to demonstrate superior cycle life. Figure 6 shows that the cell under this condition may achieve 250 cycles of continuous charge and discharge without any signs of polarization increase (Figure 6A) or capacity fading (Figure 6B). The cell operates at about 2.7 V and, therefore, holds a theoretical energy density of about 1350 Wh kg−1 considering the carbon mass. The energy content of the cell may be increased by raising the capacity limit up to 1000 or even 2000 mAh g−1, as suggested by the voltage profiles reported in Figures S3 and S4, respectively, in the SI section.66 Taking into account the cell voltage and the delivered capacity, under the various regimes, we can estimate the energy density according to a normalization factor (N) of 1/3, typically considered in lithium ion batteries,67 or a more restrictive factor

Figure 6. (A) Galvanostatic voltage profile and (B) corresponding cycling trend of the lithium oxygen battery using the TEGDME− LiTFSI (1 mol kg−1) electrolyte and the N2@MWCNT electrode in a 2032 top-meshed coin cell. Current rate of 100 mA g−1. Capacity limit 500 mAh g−1. Current density and specific capacity are referred to the N2@MWCNTs weight.

of 1/5, which may be more appropriate for the lithium oxygen battery, to include anode and electrolyte masses and inactive element weight such as current collectors and cell case (see Table S1 in SI). Table S1 reveals a minimum value of the energy density of 270 Wh kg−1 for the lithium oxygen cell limited to 500 mAh g−1 considering a normalization factor of 1/ 5, which approaches the energy density of the lithium ion battery based on LiCoO2 cathode and graphite anode reported for comparison in the same table. Remarkably, the maximum value of the energy density for the Li/O2 battery with a capacity limit of 2000 mAh g−1 and considering the less restrictive normalization factor of 1/3 exceeds 1700 Wh kg−1. These performances well suggest the suitability of the lithium oxygen cell studied in this work as a high-performance energy storage system. However, we may point out that a too high capacity limit may lead to a decay of the cell performances due to the electrode insulation, as indeed evidenced by Figure S3 in SI. In summary, we reported new insight on the characteristics of the TEGDME−LiTFSI electrolyte by combining PFG-NMR and electrochemical analyses to shed light on very important ion mobility properties of the charge carriers within the glyme media, such as self-diffusion coefficients, practical and apparent conductivities, ionic association degree of the lithium salt, and lithium transport number, which are crucial parameters allowing the electrolyte operation in efficient batteries. 16372


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ACS Applied Materials & Interfaces Diffraction patterns and electron microscopy evidenced in detail the structure and morphology of multiwalled carbon nanotubes before and after the annealing process under nitrogen atmosphere, which actually allowed the operation of the electrode in lithium oxygen cells with enhanced characteristics. Hence, we have originally shown a lithium oxygen cell using a carbon nanotube electrode with a cycle life as long as 250 cycles.

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AUTHOR INFORMATION

Corresponding Authors

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*E-mail: [email protected] (S.G.G.). *E-mail: [email protected] (J.H.).

CONCLUSIONS The cycling performance and stability of the lithium oxygen cell may be ascribed to the nature of the electrolyte due to the formation of strongly nucleophilic species (in particular O22−) during the redox process, which affect cells using common organic carbonate-based electrolytes characterized by the electrophilic carbonyl,24 whereas increased cycling stability is obtained in end-capped glyme (i.e., CH3−(OCH2CH2)n− OCH3) based on the less electrophile ether bond and more stable molecules.25 Literature papers reported lithium oxygen cells using an electrolyte formed by dissolving LiPF6 in propylene carbonate with stability limited to less than 10 cycles,24 whereas the stability was extended over 50 cycles using LiClO4 in TEGDME68 and over 150 cycles using an electrolyte based on TFSI and TEGDME,28,69 as summarized by a recent review on lithium oxygen cells.5 In this work, the lithium oxygen cell using the nitrogen-annealed MWCNT electrode and TEGDME−LiTFSI electrolyte revealed remarkable cycle life, extended up to 250 cycles, and excellent characteristics suitable for advanced high-energy storage. This outstanding performance was attributed to the excellent ion mobility and stability of the electrolyte, as revealed by NMR and EIS, as well as to an extremely reversible electrochemical process promoted by the optimal morphology of the MWCNTs evidenced by SEM and TEM. Formation and dissolution of Li2O2 during reduction and oxidation, respectively, were clearly observed by X-ray diffraction and electron microscopy. Micrometric crystalline peroxide is stably supported by the nanotube aggregates at the electrode surface, with no signs of electrode/electrolyte degradation or cell performance decay. The results obtained in this study indicate this cell as a stable, long-life, and relatively safe battery for application in modern EVs. The practical energy density of the cell is estimated to range from 270 to 1700 Wh kg−1, depending on the capacity limit adopted for cycling; however, the study reveals that excessive formation of reaction products may lead to electrode insulation, cell degradation, and capacity decay.

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100 mA g−1 and a capacity limit of 1000 mAh g−1 (Figure S4a) and 2000 mAh g−1 (Figure S4b); estimated energy density for the lithium oxygen battery studied in this work (Table S1) (PDF)

ORCID

Mallory Gobet: 0000-0001-9735-0741 Jusef Hassoun: 0000-0002-8218-5680 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the grant “Fondo di Ateneo per la Ricerca Locale FAR 2017” and the collaboration project “Accordo di Collaborazione Quadro 2015” between University of Ferrara (Department of Chemical and Pharmaceutical Sciences) and Sapienza University of Rome (Department of Chemistry). This research was also partially supported by the RISE Program at Hunter College, grant #: GM060665. The authors thank Daniela Palmeri (Electronic Microscopy Centre, Department of Chemical and Pharmaceutical Sciences, University of Ferrara, Italy) for performing electron microscopy images. The NMR measurements at Hunter College were supported by a grant from the U.S. Office of Naval Research, and the NMR Facility is supported in part by a National Institutes of Health RCMI infrastructure grant (MD007599).

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REFERENCES

(1) Tarascon, J. M.; Armand, M. Issues and Challenges Facing Rechargeable Lithium Batteries. Nature 2001, 414, 359−367. (2) Narins, T. P. The Battery Business: Lithium Availability and the Growth of the Global Electric Car Industry. Extr. Ind. Soc. 2017, 4, 321−328. (3) Lai, Y.-Q.; Xu, M.; Zhang, Z.-A.; Gao, C.-H.; Wang, P.; Yu, Z.-Y. Optimized Structure Stability and Electrochemical Performance of LiNi0.8Co0.15Al0.05O2 by Sputtering Nanoscale ZnO Film. J. Power Sources 2016, 309, 20−26. (4) Horie, H.; Abe, T.; Kinoshita, T.; Shimoida, Y. A Study on an Advanced Lithium-Ion Battery System for EVs. World Electr. Veh. J. 2008, 2, 25−31. (5) Carbone, L.; Greenbaum, S. G.; Hassoun, J. Lithium Sulfur and Lithium Oxygen Batteries: New Frontiers of Sustainable Energy Storage. Sustainable Energy Fuels 2017, 1, 228−247. (6) Ma, L.; Hendrickson, K. E.; Wei, S.; Archer, L. A. Nanomaterials: Science and Applications in the Lithium-sulfur Battery. Nano Today 2015, 10, 315−338. (7) Li, Y.; Wang, X.; Dong, S.; Chen, X.; Cui, G. Recent Advances in Non-Aqueous Electrolyte for Rechargeable Li-O2 Batteries. Adv. Energy Mater. 2016, No. 1600751. (8) Christensen, J.; Albertus, P.; Sanchez-Carrera, R. S.; Lohmann, T.; Kozinsky, B.; Liedtke, R.; Ahmed, J.; Kojic, A. A Critical Review of Li/Air Batteries. J. Electrochem. Soc. 2012, 159, R1−R30. (9) Grande, L.; Paillard, E.; Hassoun, J.; Park, J. B.; Lee, Y. J.; Sun, Y. K.; Passerini, S.; Scrosati, B. The Lithium/air Battery: Still an Emerging System or a Practical Reality? Adv. Mater. 2015, 784−800. (10) Wang, Q.; Ping, P.; Zhao, X.; Chu, G.; Sun, J.; Chen, C. Thermal Runaway Caused Fire and Explosion of Lithium Ion Battery. J. Power Sources 2012, 208, 210−224. (11) Adams, B. D.; Black, R.; Williams, Z.; Fernandes, R.; Cuisinier, M.; Berg, E. J.; Novak, P.; Murphy, G. K.; Nazar, L. F. Towards a

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsami.7b19544. Low-magnification SEM image of carbon nanotubes after the annealing process under nitrogen atmosphere (Figure S1); additional characteristics of the lithium oxygen cell using TEGDME−LiTFSI electrolyte and N2@MWCNT electrode; voltage profile of the cell cycled at a current of 100 mA g−1 and cutoff of 1.5−4.5 V, referred to the N2@MWCNTs weight and to the electrode geometric area (Figure S2); voltage profiles (Figure S3a) and cycling trend (Figure S3b) of the cell cycled at a current of 100 mA g−1 and cutoff of 1.5−4.5 V; voltage profiles of the cell cycled at a current rate of 16373


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Enhanced Lithium Oxygen Battery Using a Glyme Electrolyte and

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