Trimethylsilyl azide (TMSN3) Enhanced Li-O2 Battery Electrolytes

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Trimethylsilyl azide (TMSN3) Enhanced Li-O2 Battery Electrolytes Bhaskar Akkisetty, Ryan Rooney, Shuqi Lai, Andrew A. Gewirth, and Paul V. Braun ACS Appl. Energy Mater., Just Accepted Manuscript • DOI: 10.1021/acsaem.9b00007 • Publication Date (Web): 18 Mar 2019 Downloaded from http://pubs.acs.org on March 19, 2019

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ACS Applied Energy Materials

Trimethylsilyl azide (TMSN3) Enhanced Li-O2 Battery Electrolytes Bhaskar Akkisetty, †‡ Ryan T. Rooney, Shuqi Lai, †‡ Andrew A. Gewirth, and Paul V. Braun*†‡ †Department

of Material Science and Engineering, University of Illinois at Urbana-Champaign, Urbana, IL 61801, USA ‡Materials

Research Laboratory, University of Illinois at Urbana-Champaign, Urbana, IL 61801,

USA Department

of Chemistry, University of Illinois at Urbana-Champaign, Urbana, IL 61801, USA

Beckman

Institute for Advanced Science and Technology, University of Illinois at UrbanaChampaign, Urbana, IL 61801, USA Key words: Electrolyte, Additive, TMSN3, Bis(TMS)peroxide, Lithium nitride, Diglyme ABSTRACT Li-O2 batteries (commonly called Li-air batteries) are plagued by reactions of reduced oxygen intermediates at the cathode, as well as instability of ether-based electrolytes on the Li metal anode. Here, we study the effect of adding a non-redox mediator, trimethylsilylazide (TMSN3), to a diglyme-based electrolyte. The TMSN3 enables formation of a stable bis(TMS)peroxide from Li2O2 during discharge. The bis(TMS)peroxide can subsequently undergo decomposition during charging with a reduction in overpotential of up to 1000 mV relative to the decomposition potential of Li2O2. TMSN3 also results in formation of a robust solid electrolyte interphase (SEI), consisting of a composite of lithium nitride, reduced nitrates, and siloxanes that appears to limit both Li dendrite formation and bis(TMS)peroxide shuttling. The synergistic effect of TMSN3 on both the cathode and anode enables a high areal capacity of 3.8 mAh cm-2 at a high rate of 637 µA cm-2 for up to 30 cycles. INTRODUCTION The ever-increasing demand for high energy density secondary batteries has led to considerable interest in aprotic Li-O2 batteries. While significant challenges remain, this system is appealing because it has the highest theoretical specific energy (3500 Wh kg-1) of all existing battery technologies.1-4 Energy storage in Li-O2 batteries mechanism is provided by formation 1 ACS Paragon Plus Environment

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(discharge) and decomposition (charge) of Li2O2, concurrent with Li deplating and plating, respectively, via the generation of superoxide.1-2 During discharge, in diglyme and related solvents containing a moderate donor number anion salt (e.g., CF3SO3- (DN, 16.9)), the intermediate species superoxide (O2-) first forms an ion pair with a solvated lithium ion which subsequently dissociates chemically, or electrochemically reacts with lithium to precipitate Li2O2 in the form of a thick film on what is typically a carbon cathode. During charge, Li2O2 is oxidized, perhaps going through an unstable Li2-xO2 intermediate, leading to evolution of O2 and plating of Li.5-7 Application of aprotic Li-O2 batteries is hampered by formidable challenges, including the low roundtrip efficiency (60%) due to large overpotentials, particularly during charging, and the very practical fact that these batteries require either O2 or highly purified air, which imposes considerable overhead. Suggested reasons for the overpotential include accumulation of decomposition products on the cathode surface, poor electronic and ionic conductivity of the Li2O2 particles, and lack of contact between discharged product (Li2O2 particles) and the current collector.2, 8 Additionally, undesirable side products can arise from chemical and electrochemical reactions of intermediate superoxide, and peroxide species with carbon, binder and electrolyte.8-11 Abraham et al., introduced Pearson’s hard-soft acid base concept to facilitate the oxygen reduction reaction (ORR) and oxygen evolution reaction (OER).12 The intermediate species during the ORR and OER, LiO2, is very unstable due to the hard acid (Li+) and soft base (O2-) interaction. Larger cations like K+, Cs+, tetrabutyl ammonium ion and ionic liquids are soft acids that can form a complex with a soft base (O2-) which can delay the side reactions and also enable oxidation of Li2O2 at lower overpotential.12-15 However, to date, with these additives the capacity and cycle number achieved remained limited.12-15 Another reason for the failure of Li-O2 cells to cycle is instability of the electrolyte at the Li metal anode. Conventional carbonate-based electrolytes can form a relatively stable solid electrolyte interphase (SEI), but are highly reactive with the Li-O2 cathode intermediates.16 Ether based electrolytes that are relatively stable towards the cathode do not form a stable SEI at the anode, leading to a limited cycle life. One promising approach has been to protect the Li metal by means of LiNO3 or bromide ionomer (e.g. lithium 2-bromoethanesulfonate) electrolyte additives.7, 17-18

In another promising approach, Asadi et al. found that a thin, compact coating of Li2CO3/C

on the Li metal anode allowed a Li-O2 cell to cycle over 700 times in an air-like atmosphere, whereas the same cell containing a bare electrode failed after 11 cycles.19 Finally, Armand et al. 2 ACS Paragon Plus Environment

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investigated the effect of 2 wt% LiN3 in DME on Li-plating and stripping in a symmetrical cell configuration (Li//Li). They observed stable Li-plating and stripping over several cycles. In work somewhat related to the work here, they also studied the feasibility of using LiN3 as an additive on a Li-S full cell in a polymer electrolyte over 30 discharge-charge cycles at 70 °C.20 Here we report the use of trimethylsilyl azide (TMSN3) as an additive in a 1M LiTf in diglyme electrolyte in a Li-O2 cell. Upon the addition of TMSN3, we obtain up to a 1000 mV reduction in overpotential during charging and an enhanced cycling stability and an exceptional capacity of 3.8 mAh cm-2 at a high rate of 637 µA cm-2 over 10 cycles. We systematically studied the influence of both the trimethylsilyl (TMS+) and azide (N3-) ions on the electrochemical response. We hypothesize the TMS+ chemically converts solid lithium peroxide (Li2O2) to liquid bis(TMS)peroxide during discharge, which subsequently can undergo oxidation at reduced overpotential (~3.5 V) during charge. The N3- undergoes electrochemical oxidation (during initial charge) releasing N2 which reacts with the Li metal20 anode forming a stable, compact, highly conductive Li3N passivation layer, which both prevents the shuttling of bis(TMS)peroxide and promotes stable Li stripping and plating. EXPERIMENTAL SECTION Anhydrous diglyme (99.5%, Sigma Aldrich) was used as received, and tetraglyme (≥99%, Sigma Aldrich) was dried using 4Å molecular sieves for at least one week. Lithium trifluoromethanesulfonate (LiTf, 99.995%, Sigma Aldrich) was dried in a vacuum oven for 24 h at 140 C before being used. Nickel foam (MTI) was cleaned with isopropyl alcohol (IPA) and coated with a carbon slurry. The carbon slurry consisted of super P carbon (Alfa Aesar) and polytetrafluoroethylene (PTFE, Sigma Aldrich) prepared as indicated below. Whatman glass microfiber filters (Sigma Aldrich) were used as separators. Trimethylsilyl azide (TMSN3, 95%), trimethylsilyl chloride (TMSCl, 99%), and lithium azide solution (LiN3, 20 wt. % in H2O) were purchased from Sigma Aldrich. Pure LiN3 was prepared from the LiN3 solution as reported.20 Note, azides can explode, and were handled in an open container following instructions provided by Sigma Aldrich. Electrode preparation SuperP carbon and PTFE binder was mixed in a 90:10 ratio and ground in an agate mortar. The resulting black powder was placed in an 80:20 solution of isopropanol and DI water, and stirred 3 ACS Paragon Plus Environment


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for at least 12 h. The resulting viscous black slurry was coated on circular discs of Ni foam (10 mm diameter, ~0.6 mm thick) by ultra-sonication for 5 min and dried in vacuum for 6 h at 150 C. The carbon loading of the electrodes was 1- 2 mg cm-2. Li-O2 cell fabrication An in-house Swagelok type cell design based on a previous report21 was used to study the Li-O2 cell performance. Prior to assembling the cell, all the Swagelok cell parts and the 250 mL cylindrical glass container was dried in vacuum for 4-5 h at 75 C to remove moisture. The dried Swagelok cell components and O2 container were transferred inside the glove-box. The cell was assembled by stacking the components in the following order: Li metal disc anode, Whatman glass filter fiber separator, super P carbon-coated Ni foam cathodes. 1M LiTf in diglyme with and without added TMSN3, TMSCl and LiN3 was used as electrolytes. The volume of the electrolyte used was ~200 µL. The fabricated Swagelok cell was transferred to an O2 gas housing container (250 mL) inside the glove-box (at which point the container is filled with Ar). The O2 container with Swagelok cell was taken out of the glove-box and purged with ultra-high pure O2 gas for 30 min and maintained at 1 atm. Galvanostatic discharge/charge measurements were conducted using a Biologic VMP. All potentials are referenced against Li/Li+. Characterization After cycling, the Swagelok cell containing the TMSN3-containing electrolyte was transferred to the glove box antechamber and held for 15 min under vacuum to remove volatile components of the electrolyte. Fourier transform infrared (FT-IR) spectroscopy was measured on a Bruker Vertex70 FTIR using a globar source and a liquid nitrogen cooled mercury cadmium telluride detector. FT-IR includes KBr baseline correction. The pellets were prepared by grinding electrode material scraped off the substrate with KBr powder (dried in a vacuum oven) inside the glove box followed by pressing of a pellet. FT-IR spectra were collected under a N2 atmosphere. X-ray diffraction (XRD) was performed using a Siemens-Bruker D5000 with CuK radiation. Raman spectroscopy was performed using a confocal Raman microscope with a 532 nm excitation and a 10X magnification objective. The electrode was placed on a cover slip and sealed with a Mylar film (to prevent exposure to atmosphere) to collect the Raman spectrum. Field emission scanning electron microscopy (FESEM) was performed using Hitachi S4800 operating at 10 kV. Prior to 4 ACS Paragon Plus Environment

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XRD, Raman and FESEM measurements, cathode substrates were washed with acetonitrile to remove electrolyte. Nuclear magnetic resonance spectroscopy (H1-NMR) H1-NMR experiments were performed on a 500MHz Varian U500. To determine if bis(TMS)peroxide was present after cycling, the cell after discharge and charge cycling was transferred into a glovebox, and the separator and electrode was washed with CDCl3 to extract the electrolyte and any CDCl3-soluble products (such as bis(TMS)peroxide). The H1-NMR signal was calibrated with CDCl3. For the characterization of water-soluble parasitic products such as HCOOLi and CH3COOLi which might be present on the cathode, the cathode after dischargecharge cycling was first washed with acetonitrile to remove residual electrolyte, dried for 30 min and washed with D2O, and the NMR signal was calibrated with D2O. Differential electrochemical mass spectrometry Differential electrochemical mass spectrometry (DEMS) was performed to quantify O2 evolution of the TMSN3 containing Li-O2 cell during the charging process. The DEMS instrument was a RGA 200 from Stanford Research Systems with a quadrupole spectrometer and two-stage differential vacuum pumping system. The setup, scheme and procedures were adapted from a previous report.22-23 The electrochemical cell used to perform DEMS was adapted from previous reports and modified to connect directly to the DEMS setup.23 The Li-O2 cell was assembled in an Ar-filled glove box, then transferred out of the glove box. The cell consisted of a Li disk anode, a glass fiber separator, 10 to 15 drops of 1 M LiTF in diglyme electrolyte containing 6 v/v% TMSN3, and a carbon-coated Ni foam disk cathode. A spring was used to ensure electrical contact between the cathode and the cell.

Another

compartment was secured to the cell to create a larger headspace before purging with O2 at 1.27 atm for 10 minutes. The cell was isolated before discharging at 100 µA cm-2 for 5 h for a total capacity of 0.5 mAh. After discharge, the extra compartment was disconnected, and the main cell was connected directly to the DEMS setup. The headspace of the cell was purged rigorously with Ar to achieve an oxygen-free background. The cell was charged for 2 h at 250 µA cm-2, manually taking batch samples every 5 minutes (24 data points in total). The transfer line was purged with Ar to achieve the lowest possible O2 background before each measurement. During sampling, the 5 ACS Paragon Plus Environment


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cell rested at OCP. The headspace was purged with Ar before beginning the next 5-minute segment. The transfer efficiency of the cell (described previously) was determined to be 14% for the second transfer from the cell to the DEMS instrument. The measured O2 percentage from the DEMS data was modified by this transfer efficiency, as well as background O2 levels to determine the true O2 content in the cell headspace. The true O2 content and ideal gas law were used to determine the number of moles of O2 evolved during the 5-minute charging interval. The pressure of the cell was controlled at 1.27 atm for all procedures, the temperature of the room was 19 °C, and the cell headspace volume was measured to be 1.6 mL. The same calculation was performed for the determination of CO2 and N2 produced during charging. A leak test was performed to ensure an accurate baseline level of all species in the cell headspace. RESULTS AND DISCUSSION Electrochemical studies Cyclic voltammetry (CV) was performed in a three-electrode liquid cell to evaluate the stability of TMSN3 in the electrolyte from 2-4 V at a scan rate of 50 mV s-1. Shiny, low surface area glassy carbon was used as the working electrode, platinum (Pt) foil as the counter electrode, and lithium foil as the reference electrode. The reference electrode was enclosed in a vycor glass frit to avoid direct contact with the electrolyte. The electrolyte for CV contained 6 v/v% TMSN3 and 0.5 M LiTf in tetraglyme. A fresh solution was prepared for each CV measurement. The TMSN3 solution was prepared inside the glove-box and transferred out in a sealed three electrode cell and bubbled with O2 gas for 30 min to obtain the CV in the presence of O2. Figure 1 shows the CV plots of the TMSN3–containing electrolyte in the presence of Ar and O2. Little current is observed under Ar except for an irreversible anodic peak at 3.6 V corresponding to the oxidation of N3- to N2. Under O2, the cathodic sweep shows an onset reduction potential at 2.5 V ascribed to the ORR. The anodic peak at 3.6 V under O2 is larger than under Ar, which could be due to the cumulative oxidation of both N3- and Li2O2. This was further confirmed via DEMS as will be discussed in a later section. Based on the CV data, it appears TMSN3 is stable in the electrolyte solution over the investigated voltage range. As will be discussed, formation of N2 from the oxidation of N3- likely plays an important role in passivating Li metal, which is important for stable cycling.20

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Figure 1: CVs of three electrode Li-O2 cell using a 6 v/v% TMSN3–containing 0.5 M LiTf in TEGDME electrolyte after electrolyte saturation with O2 (black trace) or Ar (red trace). Glassy carbon was used as the working electrode, platinum foil as the counter electrode and Li metal as the reference electrode. A sweep rate of 50 mV s-1 was applied. Figure 2 shows discharge-charge profiles of Li-O2 cells containing various concentrations of TMSN3 and a TMSN3-free electrolyte (1M LiTf in diglyme). Cells were cycled at a constant current density of 127 µA cm-2, with a capacity limit of 0.64 mAh cm-2. The Li-O2 cell containing the TMSN3-free electrolyte showed a large 1.8 V overpotential and a round trip efficiency of less than 60% (Figure 2a). After adding 1, 3, and 6 v/v% TMSN3 to the electrolyte, the polarization between discharge and charge was reduced to 900 mV. In cells containing TMSN3 the first charge cycle oxidation onset potential starts at 3.5 V, which reduces to ~3V starting with the second charge cycle, which is close to the thermodynamic value of Li2O2. We hypothesize the reason for the change in voltage between the first and second cycle is that during the 1st charge the azide ion undergoes rapid oxidation predominantly at 3.5 V, before the oxidation of lithium peroxide begins. Once the azide in the electrolyte is consumed, the oxidation potential moves to ~3 V, corresponding to the oxidation of Li2O2 via the formation of bis(TMS)peroxide. The formation of bis(TMS)peroxide will be discussed in a following section. In a 1 v/v% TMSN3 cell the overpotential starts increasing after 10 cycles (Figure 2b). As the concentration of TMSN3 increased to 3 and 6 v/v% the number of cycles before the overpotential starts increasing increased



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to about 20 and 30 cycles, respectively (Figure 2c and d). While 30 cycles is not sufficient for practical applications, it is still significant given the state of art.

Figure 2: Electrochemical performance of Li-O2 cells at the indicated cycle numbers using (a) TMSN3-free electrolyte (1M LiTf in diglyme) (b-d) with the indicated volume % TMSN3 added to the electrolyte. All the cells were cycled at a constant current density of 127 µA cm-2. When we increase the concentration of TMSN3 to 12 v/v%, the overpotential of the first discharge is significantly larger than that with 1, 3 and 6 v/v%TMSN3 additive, and the cell failed after 2 cycles (Figure S1a, b). We suggest this is because the 12 v/v% TMSN3-containing electrolyte results in unstable Li plating and stripping (Figure S1c), possibly due to the formation of a thick and unstable passivation layer at the Li-metal-electrolyte interface. While we are not certain, this passivation layer may be formed from the reaction of TMSN3 with Li-metal, which would result in the formation of lithium azide (LiN3) (an insulator), and TMS+ side products in the presence of 8 ACS Paragon Plus Environment

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dissolved oxygen. The large overpotential could also be due to the poor Li-ion transport through electrolyte and carbon cathode as the surface may be fully coated with TMSN3. In contrast, the cell containing 6 v/v% TMSN3 in the electrolyte, showed a stable lithium plating and stripping behavior relative to the 12 v/v% (as well as the 3 and 1 v/v% TMSN3)-containing electrolyte. This could be the one of the reasons for the increase in cycling stability of the 6 v/v% TMSN3containing electrolyte relative to the other systems. XPS analysis of the Li metal anode after 20 cycles in the 6 v/v% TMSN3 electrolyte system indicates TMSN3 is only partially consumed through side reactions with Li metal during cycling, as XPS indicates only 1.6 atomic % Si (calculated from the survey spectra) on the electrode. XPS is highly surface sensitive, and if significant Si-containing products were present, this number would be much greater. H1-NMR on the electrolyte after 30 cycles initially containing 6 v/v% TMSN3 shows only a small amount of TMS is present in the electrolyte. The combination of these results suggests a significant fraction of the TMSN3 has simply evaporated. As the starting concentration of TMSN3 increases, it is of course available in the electrolyte for a greater number of cycles to in-situ generate bis(TMS)peroxide. Once TMSN3 has evaporated or is consumed, reactions between Li2O2, the carbon cathode and electrolyte occur, leading to an increase in overpotential.



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Figure 3: Electrochemical performance of Li-O2 cells using (a) 6 v/v% TMSCl and (b) 2 w/w% LiN3–containing electrolytes. (c, d) performance of Li-O2 cells using the 6 v/v% TMSN3containing electrolyte to high capacity limits using current densities of (c) 255 µA cm-2 (capacity limit of 1.3 mAh cm-2) and (d) 637 µA cm-2 (capacity limit of 3.8 mAh cm-2). To examine the influence of the TMS+ on the electrochemical oxidation of Li2O2, control experiments were performed by replacing TMSN3 with TMSCl. CV under argon confirmed that TMSCl is stable with no oxidation or reductions observed between 2 and 4 V (Figure S2). The charge-discharge profiles of 6 v/v% TMSCl in 1M LiTf in diglyme shows (Figure 3a) an oxidation (charging) onset potential of 3.2 V and a plateau centered at 3.6 V, as well as an increase in potential to ~ 4 V towards the end the charge. The capacity between 3.2 to 3.8 V is merely due to oxidation of Li2O2. The charge plateau between 3.8 and 4 V corresponds to the oxidation of both the Li2O2 and parasitic products (HCOOLi, CH3COOLi, and Li2CO3),11,

24

which was further 10

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supported by DEMS. The corrosive nature of chloride ions (Cl-) towards the carbon (on the substrate) and electrolyte25 causes the formation of parasitic products when TMSCl is present, even on the 1st charge. Interestingly, no increase in overpotential was observed for the measured 25 cycles, indicating perhaps the decomposition products are soluble (HCOOLi, CH3COOLi) rather than insoluble (Li2CO3). The electrochemical experiments demonstrate that, irrespective of the additive anion (N3- or Cl-), the charge plateau at ~ 3.6 V is persistent indicating that TMS+ effects the reduction in charge overpotential of the Li-O2 cell. The effect of N3- on the charge-discharge performance was investigated by adding LiN3 to a TMSN3-free electrolyte. Since the solubility of LiN3 in the electrolyte (1M LiTf in diglyme) is low, only a 2 wt. % solution was used. As expected, the flat charge curve observed at 3.5 V in the first charge cycle corresponds to the N3- oxidation, and the cell died early on the third discharge (Figure 3b). This could be due to the blockage of the electron/Li-ion path or clogging of the pore orifices by the electrochemically unreacted first and second discharge product (Li2O2) or side products. This observation further supports the notion that the TMS+ coming from the TMSN3 additive is largely responsible for the reduction in overpotential during charging. The Li-O2 cell containing 6 v/v% TMSN3 was also tested at high rates, large cycle number and at high capacity limits, comparable to that proposed for practical applications. At a current density of 255 µA cm-2 and a capacity limit of 1.27 mAh cm-2 an initial polarization of 0.9 V was observed, which increased to 1.2 V after 20 cycles (Figure 3c). When the current density was increased to 637 µA cm-2, the charge overpotential remained very low (640 mV) for the initial five cycles, even at a very high capacity limit of 3.82 mAh cm-2 (Figure 3d). By the end of the tenth cycle, the overpotential slightly increased. The Li-O2 cell containing 6 v/v% TMSN3 was also tested over 70 discharge-charge cycles at a current density of 637 µA cm-2, and capacity limit of 0.64 mAh cm-2. A stable capacity was observed up to 50 cycles and after that the potential during charging increased significantly (Figure S3). We suggest the remarkable capacity and improved efficiency and superior rate capability of the TMSN3-containing Li-O2 cell could be attributed to: (i) formation of a stable bis(TMS)peroxide that can undergo oxidation at a much reduced potential and faster rate, and (ii) the passivation of Li metal surface with Li3N, which can prevent side reactions with the electrolyte, allow stable Li-plating/stripping at reduced polarization, and prevent shuttling of the soluble bis(TMS)peroxide intermediate. 11 ACS Paragon Plus Environment


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Figure 4: DEMS showing gases evolving from the head space of a cell (a) during 1st charge (b) during 2nd charge. A constant charge current of 250 µA is used for both experiments, and data was collected for every 5 min. Prior to charge, the cells were discharged at a current of 100 µA to a capacity limit of 0.5 mAh. The 6 v/v% of TMSN3–containing electrolyte was used. DEMS was performed on the first and second charges to analyze the gases evolved from the cell and accumulated in the head space. DEMS analysis on the first charge of a cell with 6 v/v% TMSN3 showed N2 gas evolution during the initial 20 min of the total 120 min charge process. N2 gas evolution is due to the electrochemically induced decomposition of the azide ion at ~3.5 V as suggested by the CV results (Figure 2). The evolved N2 gas crosses over through the separator and reacts with the Li metal anode surface, creating a highly conductive Li3N layer (the influence of Li3N on the Li plating and stripping behavior will be discussed in a following section). DEMS shows that on the first charge the integrated number of moles of O2 evolved was 65% of the theoretical value of total number of moles of O2 consumed on first discharge (Figure 4a). This percentage increased to 80% on the second charge, while the amount of N2 gas detected on the 2nd charge decreased compared to the first charge (Figure 4b). This transformation was also clearly reflected in the charge profiles. The first charge onset potential at 3.5 V corresponds to the decomposition of N3-, whereas the reduced the second charge oxidation onset potential at ~3 V is probably from the decomposition of Li2O2 via the formation of bis(TMS)peroxide. Barile et al., reported a Li-O2 cell using a blank 1M LiTf in tetraglyme electrolyte showed an efficiency of 56%, and evolves CO2 gas from the decomposition of side products (Li2CO3, CH3COOLi, HCOOLi) above 3.6 V.22 These parasitic products are generated from the attack of highly basic Li2O2 and intermediate species (LiO2) on the carbon substrate and electrolyte. However, the Li-O2 cell 12 ACS Paragon Plus Environment

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containing TMSN3 does not show CO2 gas evolution throughout the first and second charging process.

Figure 5: H1-NMR spectra of 6 v/v% of TMSN3–containing electrolyte after the indicated discharge and charge cycles (a) broad, and (b) narrow range of chemical shifts. Analysis of bis(TMS)peroxide formation during cycling A simple iodometric chemical test was initially performed to identify the presence of bis(TMS)peroxide. The electrolyte solution from the cathode and separator was extracted with diglyme after the 1st discharge and 1st charge (two separate cells), and to that an excess amount of KI crystals were added. The cells were cycled at a current density of 127 µA cm-2 to a capacity limit of 0.64 mAh cm-2. The color of the solution after discharge changed from transparent to yellow due to the oxidation of iodide to iodine by the bis(TMS)peroxide (Figure S4). When the cell was discharged using a TMSN3-free electrolyte no yellow color was observed after KI addition 13 ACS Paragon Plus Environment


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due to the absence of peroxide species in the electrolyte. We also did not observe any change in color when the electrolyte solution containing 6 v/v% TMSN3 was tested with KI before cycling, providing evidence the color change is due to the presence of bis(TMS)peroxide. On the 1st charge the extracted solution did not show any color with KI (Figure S4), indicating formation and decomposition of liquid phase bis(TMS)peroxide occurs, respectively, on discharge and charge. We observed that dimethoxy ethane (DME), and diglymes (DEGDME) are stable towards KI whereas tetraglyme (TEGDME) oxidizes the iodide to iodine. Proton nuclear magnetic resonance spectroscopy (H1-NMR) was used to more quantitatively identify formation of bis(TMS)peroxide, and chemical changes to TMSN3 after cycling. We observed the TMSN3–containing electrolyte did not change in color and remained transparent even after 40 cycles, an indication no insoluble product is present in the electrolyte after cycling. After the first discharge, a significant change in the chemical shift value of TMSN3 was detected (Figure 5a and b). The new chemical shift () is 0.15, close to the literature value for bis(trimethylsilyl)peroxide (0.16).26 The chemical shift and overall spectra is also a nearly exact match with a bis(trimethylsilyl)peroxide-electrolyte solution we prepared by synthesizing bis(trimethylsilyl)peroxide using Li2O2 and TMSCl in 1M LiTf in diglyme. Note, in the absence of LiTf, using pure diglyme, no reaction was observed between Li2O2 and TMSCl. We suspect the LiTf acts as an acid catalyst for the formation of bis(trimethylsilyl)peroxide based on literature. Cookson, et al., first synthesized the bis(trimethylsilyl)peroxide using hydrogen peroxide (H2O2) and TMSCl in the presence of DABCO,36 and other reports use an acid catalyst such as lithium triflate.27-28 H1-NMR spectra of the as-synthesized bis(trimethylsilyl)peroxide closely matches previous reports.26,

28, 30

The H1-NMR spectra on 10th to 30th charge show only trace

bis(trimethylsilyl)peroxide in the electrolyte, along with a new peak at 0.08 ppm which could be from trimethylsilyl triflate or hexamethyldisiloxane. The peak close to 0 ppm may be solvated TMS+ in the electrolyte. During discharge, Li2O2 forms through the electrochemical reaction of Li+ with O2 on the carbon substrate. When TMSN3 was introduced in the electrolyte, it reacts with the Li2O2, forming bis(TMS)peroxide (equation 1). 2(CH3)3SiN3 + Li2O2

CF3SO3Li

(H3C)3Si

O O

Si(CH3)3 + 2LiN3

(eq 1)


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We suspect most of the bis(TMS)peroxide absorbs on the carbon until it undergoes decomposition upon charge. However, some bis(TMS)peroxide also goes into the electrolyte. This could be one reason that the DEMS shows only ~65 % of O2 evolution on 1st charge. The H1-NMR spectra after the first charge also shows the presence of unreacted bis(trimethylsilyl)peroxide in the electrolyte, perhaps coming from trace unreacted bis(trimethylsilyl)peroxide in the electrolyte. However we believe the quantity of unreacted bis(trimethylsilyl)peroxide (on charge) in the electrolyte is very small, as suggested by testing with KI; the extracted electrolyte after the 1st charge did not change color after exposure to KI (Figure S4). A similar mechanism was also observed by Zhou et al., where the solid discharge product Li2CO3 was converted into a liquid peroxodicarbonate (C2O62), in a Li-O2/CO2 cell using a super concentrated electrolyte solution, and the system then cycles via peroxodicarbonate (C2O62-) in the electrolyte.31 Another possibility is that TMS+ reacts with reduced O2 to form unstable TMSO2 (at room temperature) which undergoes further reduction or disproportionation to bis(TMS)peroxide. Formation of the intermediate trimethylsilyl superoxide (TMSO2) is very unlikely, because related radicals are very unstable even at 233 K.32-33 As shown in equation 1, the Li+ formed during discharge forms lithium azide (or is present as dissolved lithium azide in the electrolyte. X-ray photo electron spectroscopy (XPS) on the Li metal anode X-ray photo electron spectroscopy (XPS) was performed after 20 cycles to identify the influence of TMSN3 on the Li surface. Figure 6a shows the N 1s, Si 2p, F 1s, and C 1s spectra. The survey spectra, O 1s, Li 1s, and S 2p are shown in supporting information (Figure S5). The core level spectra of N 1s shows one major peak at 398.7 eV, corresponding to lithium nitride (Li3N).20, 34 The peak at 403.6 eV may be from reduced lithium nitrite (LiyNO2-x) and the higher binding energy peak at 407.3 eV could be from reduced lithium nitrate (LiyNO3-x).34 These compounds are perhaps formed by reaction of dissolved O2 with lithium azide on the surface of the Li metal. There is no FTIR and H1-NMR evidence of the presence of these reduced nitrate species on the cathode or in the electrolyte, suggesting they are insoluble in the electrolyte. It is well known that LiNO3 in the electrolyte results in the formation of an insoluble, passivating layer of reduced nitrate surface species on Li metal, which can minimize polysulfide shuttling in the Li-S battery system.34 The N 1s XPS spectra indicate TMSN3 promotes the formation of both Li3N and LiyNO3-x in our Li-O2 system. Together, we suspect these species passivate the Li metal surface preventing 15 ACS Paragon Plus Environment


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decomposition of electrolyte and inhibiting Li-dendrite growth. There are two peaks at 100.9 and 102.6 eV in the Si 2p spectra assigned to Si-C (from silanes) and -Si-O. These species are associated with silicones or lithium trimethylsilanolate (TMSOLi) obtained from the reaction of

Figure 6: (a) XPS spectra of Li metal anode after 20 cycles using 6 v/v% TMSN3–containing electrolyte. Electrochemical tests of symmetric cells with (b) additive-free electrolyte (c) with 6 v/v% TMSN3 electrolyte, and (d) 2 w/w% LiN3–containing electrolyte. The symmetric cells were cycled at a current density of 200 µA cm-2 to a capacity limit of 0.4 mAh cm-2. moisture contamination with the TMS ion on the surface of lithium metal.35 The F 1s spectra show two peaks at 685.02 eV and 688.96 eV corresponding to the LiF and -CF3 from lithium triflate, respectively.20, 36-37 The low intensity of LiF indicates that the side reaction of lithium triflate with Li metal is suppressed. The C 1s spectra show 2 peaks at 285 and 286.9 eV indicative of aliphatic 16 ACS Paragon Plus Environment

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hydrocarbons and ethers/alkoxy, both from the electrolyte solvent (diglyme).36-37 The peaks at 289.2 and 293.2 eV are assigned to the alkyl carboxylates, which could arise from the decomposition products of diglyme and from residual lithium triflate, respectively.36-37 Overall, XPS results indicate that the SEI layer attained from the TMSN3–containing electrolyte is rich in Li3N, suggesting the TMSN3 additive is the basis of a highly conducting and robust SEI layer.

Figure 7: SEM micrographs after (a) 40 h of cycling using the additive-free electrolyte (b) 740 h of cycling using the 6 v/v% TMSN3–containing electrolyte. Lithium metal plating and stripping behavior The effect of TMSN3 on the Li-plating/stripping behavior was examined and compared with a TMSN3-free electrolyte using a symmetric cell. Here we show only the additive-free and 6 v/v% TMSN3 data, as the 6 v/v% TMSN3 system was found to exhibit better results than the 1 and 3 v/v% TMSN3 systems. The symmetric Li cells were cycled at a high current density of 0.2 mA cm-2 with a capacity limit of 0.4 mAh cm-2. The voltage vs. time plots for the TMSN3-free electrolyte shows a very high nucleation barrier on the first charge (Li-plating) (Figure 6b), whereas the first charge with TMSN3 was very flat indicating a negligible nucleation barrier (Figure 6d). Moreover, the overpotential between charge and discharge in a TMSN3-containing cell was significantly reduced to 30 mV from 100 mV in an TMSN3-free electrolyte cell, perhaps implying that TMSN3 results in formation of a Li-ion conductive layer (Li3N) at the interface and promotes smooth Li-plating. The Li-plating/stripping performance in the TMSN3-free electrolyte is unstable after a few cycles, while the cell containing TMSN3 showed a very stable Liplating/stripping over several cycles. A SEM micrograph of the Li metal surface after 42 h of cycling using the additive-free electrolyte shows considerable cracking and exhibits a very porous surface, indicative of an unstable SEI (Figure 7a), in agreement with that the voltage fluctuates 17 ACS Paragon Plus Environment


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and the overpotential increases with cycling. A SEM micrograph of the Li metal cycled using the 6 v/v% TMSN3–containing electrolyte shows a very smooth and dense surface, and no crack formations even after 740 h (Figure 7b). Additional SEM micrographs of Li-metal after cycling are shown in the supporting information (Figure S6). Comparing the Li stripping/plating behavior of 6 v/v% TMSN3–containing electrolytes with that of the 2 w/w% LiN3 electrolyte shows TMSN3 to be more effective than LiN3 in imparting a stable cycling performance. However, while the Listripping/plating performance using a 2 w/w% LiN3 electrolyte was not stable over many cycles, the difference may simply be that the solubility of LiN3 in the electrolyte is at most 2 w/w% while TMSN3 is highly soluble in electrolyte.

Figure 8: SEM micrographs of the cathode after discharge to capacity limits of (a) 0.6 mAh, (b) 1.5 mAh, (c) 3 mAh, and (d) after charge to 3 mAh. A constant current density of 127 µA cm-2 was applied for all cells. Scale bar is 1 µm.


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Cathode analysis Products generated on the cathode during discharge and charge of a Li-O2 cell using the 6 v/v% TMSN3 in 1M LiTf in diglyme electrolyte were analyzed by means of FESEM, powder XRD, H1NMR, Raman and FT-IR spectroscopy. Figure 8 shows micrographs of the cathode (SuperP carbon coated Ni foam) after the 1st discharge and 1st charge cycle at a current density of 127 µA cm-2. Initially, the discharge capacity was limited to 0.5 mAh. At this discharge, no changes in morphology of the cathode were observed (Figure 8a), perhaps due to the conversion of solid Li2O2 formed to liquid bis(TMS)peroxide. Both Raman and XRD also did not detect any Li2O2 after discharge when the capacity was limited to 0.5 mAh, possibly because the small amount of Li2O2 on the cathode is below the detection limits of Raman and XRD. When the capacity limit was increased to 1.5 mAh, there is a clear indication of toroid and plate like Li2O2 particles (Figure 8b). Upon increasing the capacity limit to 3 mAh, an increase in toroid particles were observed (Figure 8c), which disappeared when the current sweep is reversed on the 1st charge (capacity limit of 3 mAh) (Figure 8d), providing indirect evidence for formation and decomposition of Li2O2 during discharge and charge, respectively. The appearance of Li2O2 particles at the higher capacity limit indicates conversion of solid Li2O2 to bis(TMS)peroxide probably terminates before full discharge. This is reasonable, as we observed that the chemical synthesis of bis(TMS)peroxide requires an ~3:1 molar ratio of TMSN3 and Li2O2. The capacity of 1.5 mAh is equivalent to the ~ 1:1 molar ratio of TMSN3 and Li2O2, calculated based on 150 µL of an electrolyte which contains 6 v/v% TMSN3. Even assuming all the TMSN3 in electrolyte is freely available, 6 v/v% TMSN3 in 150 µL of electrolyte is not sufficient to complete the reaction, so precipitation of Li2O2 is seen at 1.5 mAh capacity limit. For the measurement of Raman and XRD, a SuperP carbon coated GDL paper was used as cathode substrate, and the cells were cycled at a current density of 63.7 µA cm-2 and a capacity limit of 1.5 mAh. XRD and Raman clearly confirm the formation and decomposition of Li2O2 on discharge and charge, respectively (Figure S7). Since parasitic products such as alkyl carboxylates are amorphous and Raman signals from these products are very week, FT-IR and H1-NMR were used to detect any parasitic products upon discharge and charge as described in the methods. Cells were cycled at a current density of 127 µA cm-2 to a capacity limit of 0.64 mAh cm-2. The peaks in the FT-IR spectra between 1000 and 1400 cm-1 are assigned to lithium triflate from the electrolyte. FT-IR and H1-NMR confirm that 19 ACS Paragon Plus Environment


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there is no evidence of decomposition products after the first discharge and charge (Figure 9 and Figure S6) which we propose is due to the in-situ transformation of the highly reactive discharge product to the stable bis(TMS)peroxide, which may suppress formation of parasitic side-products. However, after 25 cycles the FT-IR spectra shows a clear evidence of formation of lithium carbonate (Li2CO3), and lithium formate (HCOOLi), and the intensity is increasing in cycles 25 to 40. The formation of lithium formate is also observed in the H1-NMR spectra after 10 cycles (Figure S8). The partial consumption of TMSN3 in side reactions with Li metal and moisture contamination, and its volatile nature at room temperature, results in there not being sufficient TMSN3 to convert all the Li2O2 product into bis(TMS)peroxide over an extended number of cycles. As a result, side products from the Li2O2 or intermediate superoxide’s reaction with carbon and electrolyte will still form and cause an increase in overpotential with prolonged cycling.

Figure 9: FT-IR spectra of the cathode after the indicated discharge or charge cycle using the 6 v/v% TMSN3–containing electrolyte at a current density of 127 µA cm-2 and a capacity cut-off of 0.64 mAh cm-2. 20 ACS Paragon Plus Environment

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CONCLUSION TMSN3 was examined for the first time as an additive in a 1M LiTf in diglyme Li-O2 battery electrolyte. Discharge-charge analysis demonstrates that TMS+ and N3- both positively influence electrochemical performance. The TMS+ enables oxidation of Li2O2 at lower potential (~3.5 V) through the formation a of a liquid phase and stable bis(TMS)peroxide, whereas the N3undergoes electrochemical decomposition at ~3.5 V, generating N2, that can cross through the separator and react with the Li metal anode to form a robust, and highly conductive Li3N passivation layer. The Li3N passivation layer can effectively prevent shuttling of bis(TMS)peroxide, and results in stable Li-metal plating and stripping. The formation of bis(TMS)peroxide was verified with H1-NMR spectroscopy. The consequence of the synergistic effect of TMSN3 on both the cathode and anode is an exceptional capacity of 3.8 mAh cm-2 at a high rate of 637 µA cm-2 over several cycles. These results reveal the potential benefits of liquid phase peroxides in Li-O2 electrolytes including a reduced overpotential during charge and passivation of the Li metal anode. This class of additive can widen the choice of electrolyte additives for other Li metal battery systems, that otherwise may be unstable against a Li metal anode. ACKNOWLEDGMENTS This work was supported by the National Science Foundation Engineering Research Center for Power Optimization of Electro Thermal Systems (POETS) with cooperative agreement EEC1449548. BA thanks the United States-India education foundation and Institute of International Education for the Fulbright-Nehru post-doctoral research fellowship. The authors thank Ashish Kukarni for FT-IR measurements, and Minjeong Shin for NMR measurements. ASSOCIATED CONTENT Supporting Information Cyclic voltammetry, XPS, XRD, SEM and H1-NMR presented in supporting information. This material is available free of charge via the Internet at http://pubs.acs.org. AUTHOR INFORMATION Corresponding author *E-mail:

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Graphical abstract


Trimethylsilyl azide (TMSN3) Enhanced Li-O2 Battery Electrolytes

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