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Two-Dimensional Phosphorene-Derived Protective Layers on a Lithium Metal Anode for Lithium-Oxygen Batteries Youngjin Kim, Dongho Koo, Seongmin Ha, Sung Chul Jung, Taeeun Yim, Hanseul Kim, Seung Kyo Oh, Dong-Min Kim, Aram Choi, Yongku Kang, Kyoung Han Ryu, Minchul Jang, Young-Kyu Han, Seung M. Oh, and Kyu Tae Lee ACS Nano, Just Accepted Manuscript • DOI: 10.1021/acsnano.8b00348 • Publication Date (Web): 01 May 2018 Downloaded from http://pubs.acs.org on May 1, 2018

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Two-Dimensional Phosphorene-Derived Protective Layers on a Lithium Metal Anode for Lithium-Oxygen Batteries Youngjin Kim,† Dongho Koo,† Seongmin Ha,† Sung Chul Jung,§ Taeeun Yim,∥Hanseul Kim,† Seung Kyo Oh,† Dong-Min Kim,† Aram Choi,† Yongku Kang,⊥ Kyoung Han Ryu,# Minchul Jang,¶ Young-Kyu Han,‡,* Seung M. Oh,† and Kyu Tae Lee†,* † School of Chemical and Biological Engineering, Institute of Chemical Processes, Seoul National University, 1, Gwanak-ro, Gwanak-gu, Seoul 08826 (Republic of Korea) ‡ Department of Energy and Materials Engineering, Dongguk University-Seoul, Seoul 04620 (Republic of Korea) § Department of Physics, Pukyong National University, 45, Yongso-ro, Nam-Gu, Busan 48513 (Republic of Korea) ∥Department of Chemistry, Incheon National University, 119 Academy-ro, Songdo-dong, Yeonsu-gu, Incheon 22012 (Republic of Korea) ⊥Advanced Materials Division, Korea Research Institute of Chemical Technology, Yuseong Daejeon 34114 (Republic of Korea) # Environment and Energy Research Team, Division of Automotive Research and Development, Hyundai Motor Company, 37 Cheoldobangmulgwan-ro, Uiwang, Gyeonggido 16082 (Republic of Korea) 1 ACS Paragon Plus Environment

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¶ Future Technology Research Center, CRD, LG Chem, Ltd., 188, Munji-ro, Yuseong-gu, Daejeon, 34122 (Republic of Korea) E-mail: [email protected], [email protected]

ABSTRACT: Li-O2 batteries are desirable for electric vehicles because of their high energy density. Li dendrite growth and severe electrolyte decomposition on Li metal are, however, challenging issues for the practical application of these batteries. In this connection, an electrochemically active two-dimensional phosphorene-derived lithium phosphide is introduced as a Li metal protective layer, where the nanosized protective layer on Li metal suppresses electrolyte decomposition and Li dendrite growth. This suppression is attributed to thermodynamic properties of the electrochemically active lithium phosphide protective layer. The electrolyte decomposition is suppressed on the protective layer because the redox potential of lithium phosphide layer is higher than that of electrolyte decomposition. Li plating is thermodynamically unfavorable on lithium phosphide layers, which hinders Li dendrite growth during cycling. As a result, the nanosized lithium phosphide protective layer improves the cycle performance of Li symmetric cells and Li-O2 batteries with various electrolytes including lithium bis(trifluoromethanesulfonyl)imide in N,N-dimethylacetamide. A variety of ex situ analyses and theoretical calculations support these behaviors of the phosphorene-derived lithium phosphide protective layer.

KEYWORDS: phosphorene, lithium phosphide, lithium metal, protective layers, lithium oxygen batteries

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A great deal of attention has been given to improving lithium-ion batteries (LIBs) to extend the driving range of electric vehicles.1-7 However, even with LIBs reaching their theoretical energy density limit, it is not possible for these batteries to power battery electric vehicles (BEVs), which can drive up to 500 km on a single charge. Therefore, new battery systems are being sought as an alternative to LIBs, and Li-O2 batteries are considered a promising candidate.8-13 The theoretical energy density (about 3,505 W h kg-1) of Li-O2 batteries is much higher than that (about 387 W h kg-1) of LIBs; this is because Li-O2 batteries store energy based on the conversion chemistry of oxygen instead of the intercalation chemistry, as used in current LIBs.14-17 Li-O2 batteries comprise an oxygen cathode, an electrolyte, and a Li metal anode. Although remarkable advancements have been achieved for Li-O2 batteries through recent pioneering works on oxygen cathodes and electrolytes,18-28 there has not been enough focus on the challenging issues for Li metal anodes. Affordable electrolytes for use in Li-O2 batteries are restricted because of the reactivity of superoxide radicals and Li metal. For example, while organic carbonates are very common and stable solvents for LIBs, they have been excluded from use in Li-O2 batteries. This is because organic carbonates chemically react with the superoxide radicals obtained from the oxygen cathode, leading to their irreversible decomposition to lithium alkyl carbonates such as Li2CO3.29-31 In this regard, amide-based organic solvents, such as N,Ndimethylacetamide (DMA), have been considered as promising electrolyte solvents because they are more chemically and electrochemically stable against nucleophilic attacks of superoxide than organic solvents such as carbonates, ethers, and sulfoxides. However, unfortunately, DMA is known to be decomposed upon contact with Li metal, which is accompanied by corrosion of the Li metal anode.32,

33

Therefore, Li metal protection is

necessary for operating Li-O2 batteries with various electrolytes.34-38 3 ACS Paragon Plus Environment

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Many efforts have been devoted to improving the Li metal stability using stable solid-electrolyte-interphase (SEI) layers formed by electrolyte decomposition, resulting in enhanced electrochemical performance.39-47 However, SEI layers are not stable enough to inhibit the deformation of the layers during cycling, resulting in an inhomogeneous local current and the formation of Li dendrites.48-51 In addition, although the SEI layers are not destroyed, the large current density enables Li metal plating on the SEI layers, leading to additional electrolyte decomposition caused by the exposure of new Li metal surfaces. For example, in previous work, the improved stability of LiNO3-derived SEI layers resulted in improved

cycle

retention

for

Li-O2

batteries

when

compared

with

lithium

bis(trifluoromethanesulfonyl)imide (LiTFSI) in DMA. Nevertheless, the LiNO3-derived SEI layers were not sufficiently stable, and they eventually deformed during cycling. This deformation caused repetitive exposure of new Li dendrites to the electrolyte, accompanied by the continuous consumption of LiNO3 during cycling; this resulted in depletion of the electrolyte and abrupt failure of the Li-O2 batteries.33, 52 This indicates that there are at least a few requirements for a Li metal protective layer in Li-O2 batteries. Electrolytes and Li plating should be thermodynamically stable and unfavorable on the protective layer surface, respectively. The mechanical strength of the protective layer should be sufficiently strong to suppress its deformation due to Li dendrite growth.53 In this paper, taking these requirements into consideration, a two-dimensional phosphorene-derived lithium phosphide on Li metal is introduced as a Li metal protective layer. The nanosized lithium phosphide protective layer exhibited the excellent electrochemical performance, with Li symmetric cells with various electrolytes showing stable cycle performance over 500 cycles, Li-O2 batteries showing no capacity fading over 50 cycles even with LiTFSI-dissolved in DMA, and suppressed Li dendrite growth on the lithium phosphide protective layer.


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RESULTS AND DISCUSSION Theoretical consideration for Li protection We propose a concept of a mechanically strong protective layer that thermodynamically inhibits electrolyte decomposition and Li dendrite formation. For this purpose, the first step is to find an electrochemically active protective layer that has a higher redox potential than the lowest unoccupied molecular orbital (LUMO) level of the electrolyte solvent (Figure 1). As electrolytes are electrochemically decomposed when the redox potential of the electrode surface is lower than the LUMO level of the solvent,54-58 electrolyte decomposition can be suppressed if the redox potential of the protective layer is higher than the LUMO level of the electrolyte solvent (less than about 0.8 V vs. Li/Li+). One appropriate candidate for this limited redox potential range is black phosphorus, which is electrochemically active with Li. Black phosphorus is spontaneously transformed into lithium phosphide (Li3P) on Li metal, acting as a protective layer. The redox potential of black phosphorus converting into Li3P is about 0.9 V versus Li/Li+, which is considered to be higher than the LUMO level of the electrolyte solvent.59, 60 However, as another form of black phosphorus, phosphorene is more practical for use as a protective layer than bulk black phosphorus as the two materials have similar redox potentials but different morphologies.61-63 Phosphorene is considered as a few layers of black phosphorus, and obtained through the exfoliation of two-dimensional layered black phosphorus.64-66 Phosphorene, with its thickness of tens of nanometers, can form thin protective layers on Li metal through conventional coating processes; in contrast, it is not


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Figure 1. Schematic diagram of the electrochemically active protective layer concept.


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possible for black phosphorus to be coated on Li in the form of a thin film because the thickness of the bulk particles can reach a few millimeters. The second step is to consider whether the plating of Li on the surface of Li3P is energetically unfavourable, in the sense that Li dendrite growth on a Li3P protective layer can be hindered even at large current densities. For the consideration of the thermodynamics of Li plating on Li3P, we first calculated the addition energy of a single Li atom on the Li3P surface using density functional theory (DFT) calculations. The (001) surface of crystalline Li3P was considered as the surface model because of its high stability and definite adsorption sites. In contrast to other surfaces, the (001) surface can have a top layer consisting of only Li atoms. Our ab initio molecular dynamics (AIMD) simulations of Li3P layer on Li metal show that only Li atoms at the surface are exposed to vacuum (Figure S1b), implying that the formation of Li top layer is energetically favored over that of P or mixed Li–P top layers. The calculated Li addition energy on the Li3P(001) surface was 0.70 eV, which is 50% lower than the average value (1.40 eV) of Li addition energies on Li(111), Li(110), and Li(001) surfaces (Figure 2a-c). This result indicates that the energy gain due to Li adsorption on Li3P is much smaller than that on Li metal. Thus, Li plating on Li3P is thermodynamically unfavourable compared with that on Li metal. We also calculated the removal energy of a single Li atom from the Li3P(001) surface to estimate the structural robustness of the surface. The calculated Li removal energy for the Li3P(001) surface was 3.06 eV, which is 79% greater than the average value (1.71 eV) of the Li removal energies from Li(111), Li(110), and Li(001) surfaces (Figure 2a). This result implies that the Li3P layer on Li metal is stably maintained during cycling because the removal of Li from the Li3P surface is energetically much more difficult than it is from the Li metal surface. Therefore, the Li addition and removal energy calculations suggest that the Li dendrite growth on the Li3P is thermodynamically hindered 7 ACS Paragon Plus Environment

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Figure 2. Thermodynamics of Li plating on lithium phosphide (Li3P). (a) Li addition and removal energies on Li metal and Li3P(001) surfaces. Structures of (b) Li(001) and (c) Li3P(001) surfaces. The Li addition and removal energies are defined as Eadd = –(Efinal – Esurf – ELi) and Erem = Efinal – Esurf + ELi, respectively. Efinal , Esurf, and ELi are the energy of Liadded or Li-removed surface, the energy of pristine surface, and the energy of an isolated Li atom, respectively. The Li addition and removal energies are the energy gain due to the adsorption of a Li atom on a pristine surface and the energy required to remove a Li atom from the pristine surface, respectively. (d) Density of states of bcc Li metal and amorphous Li3P. EF represents the Fermi level.


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without structural collapse during cycling. The thermodynamically unfavourable addition and removal of Li on the Li3P surface can be explained in terms of electrostatic interactions. Bader charge analyses of the atomic layers of the Li3P(001) surface show that the first-layer Li, second-layer Li, and second-layer P atoms (Figure 2c) have charge states of +0.85, +0.82, and –2.50 e per atom, respectively. The P atoms deprive the first-layer Li atoms of their electrons to share them with Li adsorbates, thereby weakening Li adsorption on Li3P(001). The removal of a single Li atom from Li3P(001) is also difficult because the Li atom is tightly bound through attractive electrostatic interaction between the positively charged Li and negatively charged P atoms on the Li3P surface. In addition to this thermodynamic aspect, the electrically insulating nature of Li3P may contribute to inhibiting Li dendrite growth by blocking electron transport, as evidenced by the calculated density of states of Li metal and Li3P (Figure 2d). At the Fermi level (EF), the electron density of Li3P is much lower than that of Li metal, indicating that the number of electrons contributing to the electric conductivity in Li3P is considerably smaller than that in Li metal. Finally, we need to consider whether the mechanical strength of Li3P is great enough to suppress the deformation of the Li3P protective layer due to Li dendrite growth. It is known that protective layers are easily deformed because of Li dendrite growth with a high shear modulus, indicating that a protective layer should be sufficiently strong to endure the pressure originating from Li dendrite growth. The shear modulus of Li3P obtained from DFT calculations is 34.47 Gpa,67 which is ten times higher than that of Li dendrites (3.4 GPa).53 This result indicates that even in terms of mechanical strength, Li3P is a promising candidate for Li protection.


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Synthesis of a phosphorene-derived protective layer A few layers of phosphorene sheets were obtained via facile exfoliation of black phosphorus in a carbonate-based solvent (ethylene carbonate: diethyl carbonate [EC:DEC] 1:1 volume ratio) using ultrasonication below 30 °C for 30 h. Figure 3a and b show high resolution transmission electron microscope (HR-TEM) image of exfoliated phosphorene and its corresponding fast Fourier transform (FFT) image, respectively. Phosphorene was several hundred nm in size. The FFT image reveals that the phosphorene layers retained the

crystallinity of black phosphorus during exfoliation.64, 65 Phosphorene-coated Li metal was obtained by spin-coating a solution of phosphorene in EC:DEC onto a Li metal foil in an Arfilled glove box, followed by drying. As shown in the X-ray photoelectron spectroscopy (XPS) spectra of P 2p for the phosphorene-coated Li metal at various etching times (Figure 3c), the phosphorene layers just above the Li metal surface chemically reacted to form Li3P.68 In contrast, the uppermost phosphorene layers remained unreacted because the reaction kinetics that forms Li3P is slow at room temperature. The spontaneous formation of Li3P is supported by ab initio molecular dynamics (AIMD) simulations, which revealed that P atoms on Li metal tend to mix with Li atoms to form Li3P with a large energy gain of 2.36 eV per P atom (Figure S1). We also investigated the stability of LixP layer on Li metal using DFT calculations. The LixP layer was found to be the most stable around the composition of x = 3 (Figure S2). This result suggests that the Li– P mixing on the surface of Li metal continues until the thermodynamically stable Li3P layer is formed. When an electrolyte solution of LiClO4 in EC:DEC (1:1 volume ratio) was dropped on the phosphorene-coated Li metal, a self-discharge reaction between the unreacted uppermost phosphorene layers and Li metal occurred to form Li3P, as evidenced by the XPS 10 ACS Paragon Plus Environment

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Figure 3. Characterization of the phosphorene-derived lithium phosphide (Li3P) protective layer. (a) High-resolution transmission electron microscope (HR-TEM) and (b) fast Fourier transform (FFT) images of exfoliated phosphorenes. X-ray photoelectron spectroscopy (XPS) spectra of P 2p for the phosphorene-coated Li metal at various etching times: (c) before and (d) after dropping an electrolyte solution. (e) Dynamic secondary ion mass spectrometry (SIMS) depth profiles of the phosphorene-coated Li metal. Time-of-flight SIMS (TOF-SIMS) mapping images of Li and P for (f) pristine and (g) cycled phosphorene-coated Li metal electrodes.


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spectra in Figure 3d. The thickness of the phosphorene-coating layers was roughly evaluated using time-of-flight secondary ion mass spectrometry (ToF-SIMS) under dynamic SIMS conditions (Figure 3e). The intensity of phosphorus abruptly decreased after sputtering with Cs+ ion for 50 min, indicating that the thickness of the Li3P layer was approximately less than 1 µm. In addition, the ToF-SIMS mapping image (50 µm × 50 µm) revealed that Li3P was uniformly coated on the Li metal foil (Figure 3f). The formation of thin, uniform Li3P layers is attributed to the uniform coating of phosphorene sheets on the Li metal surface via spin coating owing to their two-dimensional flat morphology. It should be noted that LiClO4 was used instead of LiPF6 as a salt in the electrolyte for the XPS and ToF-SIMS analyses to exclude phosphorus contamination caused by the electrolyte components. Moreover, to confirm that Li3P was uniformly coated on the Li metal surface, we compared the changes of the electrolytes and Li metals after the storage of each bare Li metal and phosphorene-coated Li metal in the electrolytes of 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in DMA for 10 days. Both sides of the Li metal foil were coated with phosphorene for this storage experiment. As shown in Figure S3, the electrolyte color changed to pale yellow after the storage of the bare Li metal; moreover, we observed that the Li metal also changed from a shiny silver color to black. These color changes occurred because DMA reacted chemically with the Li metal, thereby decomposing into amine-based compounds; this decomposition mechanism is discussed below. The color changes of electrolyte and Li metal can be used as an indicator to observe the exposure of Li metal to DMA. However, when the phosphorenecoated Li metal was stored in the electrolyte, no color changes were observed in the electrolyte or Li metal, which remained transparent and shiny silver even after 10 days. This means that DMA did not come into direct in contact with Li metal in the case of the phosphorene-coated Li. This supports the claim that the phosphorene was uniformly coated on the Li metal surface. 12 ACS Paragon Plus Environment

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Electrochemical performance The electrochemical performances of bare and phosphorene-coated Li metal electrodes were compared using a symmetric cell (Li/electrolyte/Li) with various electrolytes. First, the galvanostatic plating and stripping of Li were repeatedly carried out with 1 M LiTFSI in DMA as the electrolyte at current densities of 0.1 and 0.5 mA cm-2 (Figure 4a and b). The phosphorene-coated Li metal electrode exhibited highly stable and reversible Li plating and stripping over 500 and 100 cycles at 0.1 and 0.5 mA cm-2, respectively, despite using a DMA solvent. For the bare Li metal electrode, the polarization drastically increased during the initial 10 cycles, and failure was observed within tens of cycles. This failure is ascribed to severe decomposition of DMA, which is not stable against Li metal; this failure mechanism is discussed below. Moreover, to examine the stability of the phosphorene-derived protective layer in LiO2 cells, the galvanostatic plating and stripping of the phosphorene-coated Li metal electrode were evaluated using a symmetric cell with an oxygen-dissolved electrolyte of 1 M LiTFSI in DMA (Figure 4c). The phosphorene-coated Li electrode exhibited no failure and no substantial increase of polarization over 500 cycles even with the oxygen-dissolved electrolyte, which is similar to the case using the bare electrolyte without oxygen dissolution shown in Figure 4a. Moreover, we performed ex situ

31

P solid-state nuclear magnetic

resonance (NMR) analysis to observe the change in the Li3P phase before and after storing Li3P in the oxygen-dissolved electrolyte for 24 h. The Li3P phase was obtained through the electrochemical lithiation of phosphorus powder electrodes, where the electrodes were discharged with a Li counter electrode to 0.01 V versus Li/Li+. Following this, the lithiated phosphorus electrode was stored in the oxygen-dissolved electrolyte for 24 h. Figure S4 shows the 31P NMR spectra of the lithiated electrodes before and after storage in the oxygendissolved electrolyte. The singlet at –278 ppm and doublet at –6.5 and +5.5 ppm reveal the 13 ACS Paragon Plus Environment

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Figure 4. Electrochemical performance of Li metal symmetric cells. Electrolyte in the Li/electrolyte/Li symmetric cells: 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in N,N-dimethylacetamide (DMA) Charge–discharge voltage profiles for the plating/stripping of bare (black) and phosphorene-coated (red) Li metal electrodes. Current pulse: (a) ± 0.1 and (b) ± 0.5 mA cm-2 for each 0.5 h cycle. (c) Charge–discharge voltage profiles for the plating/stripping of phosphorene-coated Li metal electrodes using the electrolytes with oxygen dissolution (black) or without oxygen dissolution (red). (d) Rate performance of the phosphorene-coated Li symmetric cell. Current pulse: ± 0.1 through 3 mA cm-2 for each 0.5 h cycle. 14 ACS Paragon Plus Environment

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existence of Li3P and LiP, respectively.69 No change was observed in the 31P NMR spectra of the lithiated electrode before and after the storage. This indicates that the phosphorenederived Li3P protective layer is stable in the oxygen-dissolved electrolyte. Figure 4d demonstrates the rate performance of the phosphorene-coated Li using a symmetric cell with 1 M LiTFSI in DMA. The phosphorene-coated Li metal exhibited reversible plating and stripping with small polarization (about 150 mV) even at a high current density of 3 mA cm-2. This is not only because the Li3P layers are thin but also because the ionic conductivity of Li3P (10-4 S cm-1)70 is sufficiently high at room temperature, reaching as much as that of LISICON (10-4–10-6 S cm-1).71, 72 We evaluated the electrochemical performance of Li-O2 batteries using the bare and phosphorene-coated Li metal anodes (Figure 5). 1 M LiTFSI in DMA was used as the electrolyte, and the galvanostatic charging and discharging were performed at 250 mA g-1 (0.25 C). Despite using 1 M LiTFSI in DMA, the phosphorene-coated Li metal electrode achieved excellent cycle performance with no capacity fading over 50 cycles when cycled with a constant capacity of 1000 mA h g-1 (Figure 5a). However, the Li-O2 cell with the bare Li metal electrode failed within 10 cycles, which is consistent with the failure of the symmetric cell consisting of bare Li metal electrodes with 1 M LiTFSI in DMA (Figure 4a). This difference in performance is attributed to the improved chemical and electrochemical stabilities of the phosphorene-coated Li metal electrode, which suppressed electrolyte decomposition and Li dendrite growth. The corresponding voltage profiles of these electrodes for Li-O2 batteries are displayed in Figure 5b and c. The Li-O2 cell with the phosphorenecoated Li metal exhibited reversible voltage profiles over 50 cycles, although a large polarization was observed because an oxygen cathode including no catalysts was used. To clarify the reaction mechanism of the Li-O2 cell with the phosphorene-coated Li metal 15 ACS Paragon Plus Environment

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Figure 5. Electrochemical performance of Li-O2 batteries. (a) Cycle performance of Li-O2 cells, and the corresponding galvanostatic charge-discharge voltage profiles for (b) phosphorene-coated Li and (c) bare Li metal anodes using the 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in N,N-dimethylacetamide (DMA) as the electrolyte. (d) Cyclic voltammograms of the electrolyte (1 M LiTFSI in DMA) before and after the storage of bare Li metal for 24 h.


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electrode, we performed the ex situ X-ray diffraction (XRD) analysis of the fully discharged oxygen electrode. As shown in Figure S5, XRD peaks corresponding to a Li2O2 phase were clearly observed. This means that the Li-O2 cell with the phosphorene-coated Li metal electrode proceeded through the same reaction mechanism as the previously-reported Li-O2 cells, where Li2O2 was reversibly formed during charging and discharging.20, 31, 73 Therefore, the charging potential at about 4 V versus Li/Li+ in Figure 5b is attributed to the reversible decomposition of Li2O2. In contrast to the Li-O2 cell with the phosphorene-coated Li metal, the Li-O2 cell with the bare Li metal showed unusual charging profiles at voltages lower than 4 V versus Li/Li+ in the initial cycles. To demonstrate the origin of the low charging potentials, we compared the cyclic voltammograms of the electrolyte (1 M LiTFSI in DMA) before and after the storage of bare Li metal for 24 h (Figure 5d). A glassy carbon electrode (electrode area = 0.07 cm2), platinum flag, and silver wire were used as the working, counter, and reference electrodes, respectively. The redox potential of the quasi-reference electrode of a silver wire was calibrated using a Ferrocene redox couple. When the cyclic voltammetry was carried out with the pristine electrolyte before the storage of Li metal, no redox peaks were observed. However, the irreversible oxidation peak at about 2.8 V versus Li/Li+ was observed when we examined the electrolyte storing Li metal for 24 h. This means that the soluble decomposed products of DMA were irreversibly oxidized at about 2.8 V versus Li/Li+. Therefore, the unusual low charging potentials of the Li-O2 cell with the bare Li metal in the initial cycles are attributed to the irreversible oxidation of decomposed products obtained from DMA. Similar electrochemical behaviours were observed for Li symmetric cells with other electrolytes such as 1 M LiNO3 in DMA (Figure 6a) and 1.3 M LiPF6 in EC:DEC (Figure 6b and c). For 1 M LiNO3 in DMA, the symmetric cell with phosphorene-coated Li metal 17 ACS Paragon Plus Environment

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Figure 6. Electrochemical performance of Li metal symmetric cells with various electrolytes. Galvanostatic charge-discharge voltage profiles of Li metal symmetric cells for the plating/stripping of bare (black) and phosphorene-coated (red) Li metal electrodes. Electrolytes in Li/electrolyte/Li symmetric cells: (a) 1 M LiNO3 in N,N-dimethylacetamide (DMA) and (b), (c) 1.3 M LiPF6 in EC:DEC. Current pulse: (a) ± 0.1 mA cm-2 for each 0.5 h cycle, (b) ± 0.1 mA cm-2 for each 0.5 h cycle, and (c) ± 2 mA cm-2 for each 0.5 h cycle.


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electrodes exhibited stable, reversible Li plating and stripping over 500 cycles. However, the symmetric cell with bare Li metal electrodes showed an abrupt increase in polarization after about 110 cycles, although stable Li plating and stripping were observed before this increase. It was recently reported that LiNO3 decomposes on Li metal surfaces, forming stable SEI layers comprising lithium oxides, when LiNO3 in DMA is used as the electrolyte for a Li metal anode.32 The improved stability of LiNO3-derived SEI layers resulted in improved cycle retention compared with LiTFSI in DMA. Nevertheless, the LiNO3-derived SEI layers were not sufficiently stable, leading to the deformation of SEI layers during cycling. This deformation caused the continuous consumption of LiNO3 during cycling,32 resulting in the depletion of the electrolyte and abrupt failure of the symmetric cell. For 1.3 M LiPF6 in EC:DEC, the symmetric cells of both bare and phosphorenecoated Li metal electrodes showed no failure during cycling. However, while the polarization of the phosphorene-coated Li metal electrode was stable over 500 cycles, that of the bare Li metal electrode gradually increased after 230 cycles, eventually becoming much larger than that of the phosphorene-coated electrode. The stable polarization of the phosphorene-coated Li metal electrodes was also observed even at a high current density such as 2 mA cm-2.74 (Figure 6c) Further, we examined the electrochemical performance of the other Li metal batteries comprising LiCoO2, Li metal, and carbonate-based electrolytes. Figure S6 shows the cycle performance of LiCoO2 at different voltage ranges, such as 3 – 4.25 V and 3 – 4.4 V versus Li/Li+. While abrupt capacity fading was observed for the LiCoO2 cell with the bare Li metal after about 260 and 60 cycles, the LiCoO2 cell with the phosphorene-coated Li metal exhibited stable cycle performance over 300 and 100 cycles. The much better capacity retention of the LiCoO2 cell with the phosphorene-coated Li is attributed to suppressed


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dendrite formation during cycling, which is consistent with the cycle performance of the Li metal symmetric cells with a carbonate-based electrolyte in Figure 6b.

Protection and failure mechanisms To clarify that Li3P protective layers suppress Li dendrite growth, we performed ex situ SEM analysis to observe Li dendrite growth of the bare Li and the phosphorene-coated Li metal electrodes after Li plating at 0.1 mA cm-2 for 20 h. This experimental condition corresponds to the areal capacity of 2 mA h cm-2, which is similar to the high areal capacity value (about 2.25 mA h cm-2) of the practical Li metal batteries comprising LiCoO2 and Li metal, assuming that the reversible capacity and loading amount of LiCoO2 is 150 mA h g-1 and 15 mg cm-2, respectively. For this comparison, 1 M LiTFSI in EC:DEC (1:1 volume ratio) was used as an electrolyte instead of 1M LiTFSI in DMA, as it is not possible to observe Li dendrites for 1M LiTFSI in DMA owing to the severe electrolyte decomposition of DMA on bare Li metal. Figure 7 shows the SEM images of the bare Li and the phosphorene-coated Li metal electrodes after Li plating. While Li dendrites with a thickness of a few micrometers were clearly observed for the bare Li electrode, no dendrites were observed for the phosphorene-coated Li metal; instead, Li metal with agglomerated spherical morphologies of a size of tens of micrometers was observed after Li plating. This suggests that the growth of Li dendrites was suppressed on the Li3P protective layer. This is consistent with the theoretical calculations and the electrochemical performance of the Li-O2 cells and Li symmetric cells (Figures 4, 5, and 6). However, the surface roughness of the phosphorenecoated Li metal increased from the smooth surface after Li plating, indicating that Li metal plating was not completely uniform underneath the Li3P protective layer when the areal


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Figure 7. Effect of the phosphorene-derived Li3P protective layers on Li dendrite growth. Ex situ field emission scanning electron microscopy (FE-SEM) images of (a) bare and (b) phosphorene-coated Li metal electrodes with 1 M LiTFSI in EC:DEC (1:1 volume ratio) after Li plating at 0.1 mA cm-2 for 20 h (areal capacity: 2 mA h cm-2).


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capacity was as high as 2 mA h cm-1. This implies that Li dendrites may form underneath the Li3P protective layer as the amount of Li plating increases. We also examined the failure mechanism of Li metal in DMA-based electrolytes to understand how the phosphorene-derived Li3P layer protects Li metal against DMA-based electrolytes. When galvanostatic plating and stripping of the bare Li symmetric cell were performed with the electrolyte of 1 M LiTFSI in DMA as shown in Figure 4, a color change from transparent to pale yellow was observed in the electrolyte solution after 100 cycles (Figure 8a). This indicated that the electrolyte was decomposed during cycling as such a color change is characteristic of the chemical decomposition of amine-based materials.75 The same color change was also observed in the optical image of the bare Li metal electrode after 100 cycles (Figure 8b).However, for the phosphorene-coated Li metal electrode, no color change was observed in either the electrolyte or Li metal electrode after 100 cycles. Moreover, we performed ex situ SEM analysis to compare a change in the morphologies of the phosphorene-coated Li metal before and after cycling performed at ± 0.1 mA cm-2 for each 0.5 h cycle (areal capacity: 0.05 mA h cm-2) over 100 cycles. As shown in Figure 8c and d, no dendrites were observed after 100 cycles even with 1 M LiTFSI in DMA. As we compared the SEM images of the phosphorene-coated Li metal after Li plating at 2 mA h cm2

(Figure 8d) and 100 cycles at 0.05 mA h cm-2 (Figure 7b), the latter showed a smoother

surface despite the fact that the latter was obtained after 100 cycles. This is because much less Li was plated and stripped during cycling compared to the former. These results indicate that, in contrast to Li metal, the Li3P protective layer was stable even in 1M LiTFSI in DMA, thereby suppressing electrolyte decomposition and dendrite growth. This is consistent with the electrochemical performances of the bare and the phosphorene-coated Li metals shown in


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Figure 8. Decomposition of N,N-dimethylacetamide (DMA)-based electrolytes. Ex situ optical images of (a) electrolytes and (b) Li metal electrodes collected after 100 cycles. Ex situ field emission scanning electron microscopy (FE-SEM) images of the phosphorenecoated Li metal electrode with 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in DMA (c) before and (d) after cycling performed at ± 0.1 mA cm-2 for each 0.5 h cycle (areal capacity: 0.05 mA h cm-2) over 100 cycles. (e) Ex situ nuclear magnetic resonance (NMR) spectra of the electrolyte consisting of 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in DMA collected after 100 cycles for bare and phosphorene-coated Li metal symmetric cells. (f) Proposed decomposition pathways for DMA-based electrolytes.


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Figure 4. The suppressed electrolyte decomposition on Li3P is attributed to the fact that the redox potential of Li3P is sufficiently high in relation to the LUMO level of DMA. To elucidate the decomposition mechanism of the DMA-based electrolyte on the bare Li metal electrode, ex situ NMR and gas chromatography mass spectrometry (GC-MS) analyses were performed using the electrolyte collected after 100 cycles (Figure 8e). The ex situ NMR spectra revealed informative clues about the decomposition pathway of DMA. The singlet NMR peak at 3.19 ppm was assigned to N,N-dimethylamine. The ex situ GC-MS analysis provided further evidence for the decomposition of DMA into N,N-dimethylamine and N,N-dimethylacetoacetamide (Figure S7). Taking these side-products into account, we propose that the decomposition of DMA proceeds via a two-electron transfer reaction followed by further cascade-type chemical reactions (Figure 8f). The stable cycle performance and lack of color change for the phosphorene-coated Li metal electrode strongly indicate not only that the Li3P protective layer was stable against DMA but also that the protective layer was not deformed and Li metal was not plated on the protective layer surface during cycling. Otherwise, severe decomposition of DMA should be observed with Li metal being newly exposed to the electrolyte during cycling. This inference was supported by the ex situ ToF-SIMS analysis (Figure 3g) of the phosphorene-coated Li metal electrode. In the ex situ ToF-SIMS mapping image (50 µm × 50 µm), phosphorus was uniformly detected on the surface of the Li metal electrode, even after 100 cycles. When considering that the depth resolution of the ToF-SIMS instrument is less than 2 nm, the observation of a uniform distribution of P proves that i) the Li3P protective layer was not destroyed during cycling, ii) Li metal was not plated on the surface of the Li3P protective layer during cycling, and iii) electrolyte was rarely decomposed on the Li3P layer because SEI layers were negligibly thin enough to allow detection of P in Li3P. 24 ACS Paragon Plus Environment

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CONCLUSIONS We introduced a thermodynamic approach to Li metal protection for Li-O2 batteries, and demonstrated the excellent electrochemical performance of the phosphorene-coated Li metal electrode in Li symmetric cells which showed stable cycle performance over 500 cycles, and in Li-O2 batteries, which exhibited no capacity fading over 50 cycles despite using LiTFSI in DMA. Moreover, no growth of Li dendrites was observed on the Li3P protective layer. This improved performance is attributed to thermodynamic suppression of electrolyte decomposition and Li dendrite growth by the electrochemically active phosphorene-derived Li3P protective layer, as evidenced by a variety of ex situ analyses, theoretical calculations, and electrochemical performances. The redox potential of Li3P is about 0.9 V versus Li/Li+ higher than the LUMO level of the electrolyte solvents (less than ca. 0.8 V vs. Li/Li+), leading to the suppression of electrolyte decomposition. The thermodynamic preference for Li plating on Li3P protective layers based on Li addition and removal energy calculations suggests that the Li3P layer on Li metal hinders Li dendrite growth during cycling. Moreover, the high mechanical strength of the Li3P layer contributes to improving the electrochemical performance of the phosphorene-coated Li metal anode. Finally, we consider that this approach for Li metal protection can be extended to improve the electrochemical performance of Na metal batteries by developing Na3P protective layers.

METHODS Fabrication of phosphorene-coated Li metals: Black phosphorus (Smart Elements, 70 mg) was exfoliated in a solution of EC:DEC (1:1 volume ratio, 30 mL) in an Ar atmosphere by ultrasonication using an Elmasonic P 70H ultrasonic bath operating at 37 kHz below 30 °C for 30 h. A 100-µL sample of the as-prepared phosphorene suspension was spin coated onto a Li metal foil (diameter: 14 mm) in an Ar-filled glove box and then dried at room temperature. 25 ACS Paragon Plus Environment

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Electrochemical characterization: The electrochemical performance of symmetric cells consisting

of

phosphorene-coated

Li/electrolyte/phosphorene-coated

Li

and

bare

Li/electrolyte/bare Li was examined using a 2032 coin cell with various electrolytes, including 1.3 M LiPF6 in EC:DEC (3:7 volume ratio), 1 M LiTFSI in EC:DEC (1:1 volume ratio), 1 M LiTFSI in DMA, and 1 M LiNO3 in DMA. A phosphorene-coated Li metal symmetric cell was also examined using a Swagelok-type cell with oxygen-dissolved electrolyte of 1 M LiTFSI in DMA. Oxygen dissolution in electrolytes was performed by purging the cell with oxygen gas at 40 standard cubic centimeters per minute (sccm) for 15 min, and then the cells were rested for 1 h before electrochemical measurement. Galvanostatic charge and discharge experiments of the symmetric cells were performed for Li metal plating and stripping at 0.1, 0.5, and 2 mA cm-2 for each 0.5 h cycle. We evaluated the cycle performance of Li-O2 cells with a Swagelok-type cell consisting of a carbon cathode, a bare or phosphorene-coated Li metal anode (12 mm diameter), and an electrolyte of 1 M LiTFSI in DMA. To prepare the air cathode, a slurry comprising Ketjen Black (80 wt.%) and a poly(vinylidene fluoride) (PVdF) binder (20 wt.%) was coated on a carbon paper current collector. The electrodes were dried at 120 oC in a vacuum oven overnight, and the loading mass was approximately 0.2–0.4 mg cm-2. Galvanostatic cyclings were performed at 250 mA g-1 between 2.0 and 4.5 V (vs. Li/Li+) with a constant capacity of 1000 mA h g-1. To observe the color change of the electrolyte after cycling, beaker-type cells were used instead of coin cells. To examine ex situ

31

P solid state NMR, the Li3P phase was prepared using the

electrochemical lithiation of red phosphorus/carbon composites (7:3 weight ratio). The phosphorus electrodes were comprised of red P/C composites (70 wt. %), Super P (10 wt. %) and a poly (acrylic acid) (PAA) binder (20 wt. %). The LiCoO2 electrodes comprised LiCoO2 (90 wt. %), Super P (5 wt. %), and a PVDF binder (5 wt. %). The loading amounts of LiCoO2 were approximately 8 mg cm-2. Galvanostatic experiments of LiCoO2 were carried out 26 ACS Paragon Plus Environment

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between 3.0 – 4.25 V versus Li/Li+ and 3.0 – 4.4 V versus Li/Li+ at 1 C-rate (150 mA g-1) using 1.3 M LiPF6 in EC:ethylmethyl carbonate (EMC):DMC (3:4:3 volume ratio) and 1.3 M LiPF6 in EC: DEC (1:1 volume ratio), respectively.

Computational details: The DFT calculations were carried out using the Vienna ab initio simulation package (VASP). We employed the Perdew-Burke-Ernzerhof (PBE) functionals for electron‒electron interactions and the projector augmented wave (PAW) method for electron‒ion interactions. We expanded the electronic wave functions in a plane wave basis set of 271.6 eV. The valence electron configurations were set as 1s22s1 for Li and 3s23p3 for P. The P-covered Li metal surface was simulated using periodic slab geometry, where each slab consisted of 3 and 29 atomic layers for P and Li (18 P and 174 Li atoms), respectively, and the vacuum spacing was about 16 Å. For Brillouin zone integrations, a 4 × 6 × 1 k-point mesh was used. We used AIMD simulations to allow the P atoms to mix with Li atoms for the formation of Li3P. Similar calculation methods were described in our previous work on the lithiation of P.76

Material Characterization: TEM analysis was performed using a high-resolution transmission electron microscope (STEM, JEOL ARM-200F). XPS experiments were carried out in an UHV multipurpose surface analysis system (SIGMA PROBE, Thermo, UK) operating at base pressures below 10-9 mbar. The XPS system was connected with a glove box to prevent air exposure when transferring the Li metal samples. The photoelectron spectra were obtained following excitation by an Al K (1486.6 eV) anode operating at a constant power of 150 W (15 kV and 10 mA). ToF-SIMS analyses were performed on an a 27 ACS Paragon Plus Environment

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TOF-SIMS 5 system with a Bi+ primary ion beam source (ION-TOF GmbH, Germany). A pulsed 25-keV Bi+ beam was used as the analysis source. The target current was maintained at a pulsed current of 1 pA with a raster size of 50 µm × 50 µm. A Cameca IMS 4FE7 instrument using a 14.5-keV Cs+ primary ion beam was used to obtain dynamic SIMS depth profiles. Depth profiles were obtained with a raster size of 200 µm × 200 µm at a primary beam current of 40 nA and a sputtering rate of 18.5 nm min-1. NMR and GC-MS analyses were carried out using a Bruker ASCEND400 spectrometer and an Agilent 7890B GC-MS, respectively. Each 100 µL of collected electrolyte was dissolved in 1.5 mL of acetone-d6 (deuterated acetone containing 0.1% tetramethylsilane as an internal standard for NMR spectroscopy), and the resulting solutions were measured at ambient temperature (25 °C). For all ex situ experiments, the cycled cells were disassembled in an Ar-filled glove box, and then the electrodes were rinsed with DMC. In addition, the ex situ

31

P solid state magic-angle

spinning (MAS)-NMR spectra were recorded with a Bruker Avance II NMR spectrometer. The 31P chemical shifts were referenced to 85 % H3PO4. The 31P spectra were obtained at 163 MHz with a pulse angle of π/2, 4.0 µs pulses and 50.0 s pulse delays. The spinning speed was approximately 8 kHz. FE-SEM images were obtained using a FE-SEM (JSM-7800F Prime, JEOL Ltd) equipped with energy dispersive X-ray spectroscopy (EDS). The ex situ XRD data were collected on a Rigaku D/MAX2500 powder diffractometer using Cu-Kα radiation (λ = 1.5405 Å, 40 kV and 200 mA)

ASSOCIATED CONTENT Supporting Information. The Supporting Information is available free of charge on the


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ACS Publications website at DOI: Figures S1-S7, DFT calculations, photo images of Li metals and electrolytes, 31P solid state NMR spectra, ex situ XRD analysis, electrochemical data, and GC-MS analysis (PDF)

AUTHOR INFORMATION Corresponding Author * E-mail: [email protected] (K. T. Lee) * E-mail: [email protected] (Y. K. Han)

ACKNOWLEDGMENTS This research was supported by a National Research Foundation of Korea (NRF) Grant (No. NRF-2016R1A2B3015956,

2016R1A2B4013374,

2016R1C1B1009452

and

2016R1A6A3A01007701), and by the Ministry of Trade, Industry & Energy (MOTIE, Korea) under Industrial Technology Innovation Program. No. 10063288, ‘Development of highly stable inverse-electrolyte-systems for high-energy density (350 Wh/kg) Li-metal rechargeable batteries with long cycle life (500th)’. We thank the National Center for InterUniversity Research Facilities (NCIRF) at Seoul National University (SNU) for assistance with the 500-MHz solid-state NMR and FE-SEM experiments.

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