Aqueous Rechargeable Li and Na Ion Batteries - American Chemical

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Aqueous Rechargeable Li and Na Ion Batteries Haegyeom Kim,†,∥ Jihyun Hong,†,∥ Kyu-Young Park,†,∥ Hyungsub Kim,†,∥ Sung-Wook Kim,§ and Kisuk Kang*,†,‡ †

Department of Materials Science and Engineering, Research Institute of Advanced Materials (RIAM), Seoul National University, Gwanak-ro 1, Gwanak-gu, Seoul 151-742, Republic of Korea ‡ Center for Nanoparticle Research, Institute for Basic Science (IBS), Seoul National University, Gwanak-ro 1, Gwanak-gu, Seoul 151-742, Republic of Korea § Nuclear Fuel Cycle Development Group, Korea Atomic Energy Research Institute, 989-111 Daedeok-daero, Yuseong-gu, Daejeon 305-353, Republic of Korea 7.2. H2 and O 2 Evolution from Electrolyte Decomposition 7.3. Electrode Dissolution in an Aqueous Electrolyte 7.4. Proton Co-Insertion with Guest Ions 7.5. Outlook Author Information Corresponding Author Author Contributions Notes Biographies Acknowledgments References

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CONTENTS 1. Introduction 2. Overview 3. Aqueous Rechargeable Lithium Batteries 3.1. Introduction to Aqueous Rechargeable Lithium Batteries 3.2. Electrodes for Aqueous Rechargeable Lithium Batteries 3.2.1. Cathode Materials 3.2.2. Anode Materials 4. Aqueous Rechargeable Sodium Batteries 4.1. Introduction to Aqueous Rechargeable Sodium Batteries 4.2. Electrodes for Aqueous Rechargeable Sodium Batteries 4.2.1. Cathode Materials 4.2.2. Anode Materials 5. Host Materials for Alternative Guest Ions 5.1. Prussian Blue Analogues 5.2. Manganese Oxides 5.3. Titanium Oxides 5.4. Vanadium Pentoxides 6. Alternative Aqueous Rechargeable Battery Systems 6.1. Metal Anode in an Aqueous battery 6.2. Li/Na Aqueous Air Battery 6.3. Aqueous Cathode Battery 7. Challenges and Perspectives 7.1. Side Reactions with H2O and O2 in an Electrolyte

© 2014 American Chemical Society

1. INTRODUCTION The increased use of fossil fuel combustion to produce electricity increases carbon dioxide gas levels in the atmosphere, causing a greenhouse effect, which is a major cause of global warming.1,2 Thus, electricity from renewable and sustainable energy resources such as solar, wind, and tide has moved to the center of attention. However, renewable energy resources are only intermittently available and are dependent on the time, weather, season, and location, while the demands and consumption of electric energy are relatively constant. Consequently, large-scale stationary energy storage systems (ESSs) connected to renewable power plants have become key enablers of improving power reliability and quality as well as taking full advantage of high penetration of renewable energy sources.1,3,4 Essential criteria required for large-scale ESSs are (i) low costs of construction and maintenance, (ii) low risk of safety incidents for long-term utilization, (iii) high round-trip efficiency, and (iv) long cycle life. In addition, environmental benignness and nontoxicity should be considered. According to the Energy Storage Association, ESSs available for large-scale applications can be categorized into four types, mechanical, electrical, chemical, and electrochemical.3−5 Although energy storage is currently dominated by mechanical energy storage via pumped hydroelectricity, electrochemical ESSs, i.e., recharge-

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Received: April 28, 2014 Published: September 11, 2014 11788

dx.doi.org/10.1021/cr500232y | Chem. Rev. 2014, 114, 11788−11827

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Figure 1. Electrochemical stability range of water and redox potentials for electrode materials in LIBs and NIBs.

aqueous electrolyte. In this regard, the suitability of ARABs for large-scale applications has recently been revisited.6,17

able batteries, possess a number of desirable characteristics such as high round-trip efficiency, long cycle life, pollution-free operation, controllable power and energy to meet different grid functions, and low maintenance costs.1−4 In this respect, electrochemical ESSs, particularly rechargeable batteries, have been considered as promising candidates for power sources for large-scale applications. The most common rechargeable battery systems are lithium ion batteries (LIBs), which show the highest energy density, cycle stability, and energy efficiency6,7 among various rechargeable batteries.6,8 For portable applications, it is the most critical feature that the battery is capable of storing large amounts of energy in a given volume and mass in a short time. Thus, this feature has guided most investigations of LIBs to achieve high volumetric (W h L−1) and gravimetric (or specific) (W h kg−1) energy density combined with high volumetric (W L−1) and gravimetric (W kg−1) power density.8−11 As a result, LIBs have been successfully used as power sources in most of today’s portable electronics. Although LIBs have been optimized to meet the requirements of portable electronics, some intrinsic characteristics make the current LIBs less feasible for large-scale stationary ESSs, where the cost, safety, and long cycle life become relatively more important than energy densities.12,13 One of the major issues is safety. The high risk of safety incidents in LIBs is attributed to the flammable organic electrolyte and the thermal runaway caused by the reactivity of the electrode materials with electrolytes. In addition, the cost of LIBs is relatively high due to the special cell assembly technology, the requirement of a strictly dry environment during manufacturing processes, and the high price of transition metals, organic electrolytes, and lithium salts. Furthermore, the limited ion conductivities of the organic electrolytes require battery designs with thin electrodes for high power and energy efficiency. Aqueous rechargeable alkali-metal ion (Li+, Na+) batteries (ARABs) are promising alternatives for large-scale applications which resolve several challenges of conventional LIBs: (i) the safety issue of flammable organic electrolytes is fundamentally resolved, (ii) the rigorous manufacturing conditions are avoided, and the prices of the electrolyte solvent and salts are relatively low, and (iii) the ionic conductivity of the aqueous electrolyte is higher than those of organic electrolytes by 2 orders of magnitude,14−16 resulting in high round-trip efficiency and energy density even with bulky and scalable electrodes. In addition, they are based on the environmentally benign

2. OVERVIEW The overall electrochemistry of ARABs is identical to that of the conventional rechargeable battery system based on organic electrolytes, which transfers alkali-metal ions (Li+ or Na+) through the electrolytes and electrons through the external circuit between two electrodes (i.e., the anode and cathode). For example, the electrochemical reaction between the LiMn2O4 cathode and VO2(B) anode can be expressed as follows:16 LiMn2O4 ⇄ Li1 − xMn2O4 + x Li+ + x e−

(1)

VO2 (B) + x Li+ + x e− ⇄ LixVO2 (B)

(2)

However, the stable voltage window of aqueous electrolytes is narrower than that of the organic electrolytes used in the current batteries. The stable operating voltage window is approximately 1.23 V, beyond which the electrolysis of H2O occurs with H2 or O2 gas evolution. Thus, the energy density of the aqueous system is not generally as high as that of the organic system, which typically operates above 3.0 V for LIBs. Materials whose working potential is located between the H2 evolution potential and O2 evolution potential are suitable for the aqueous alkali-metal ion battery system. In addition, the electrode materials should be carefully selected with consideration of the pH, which is strongly associated with the H2 and O2 evolution potentials of the aqueous system.18,19 For decades, studies have explored a variety of electrode materials for LIBs and Na ion batteries (NIBs) with organic electrolytes, and a database on the working potentials (vs Li+/Li or Na+/Na) of these materials is available. Note that “potential” is not an absolute value, but a relative value compared to the reference. Hence, it is possible to convert the working potentials of such electrode materials in reference to the standard hydrogen electrode (SHE), which is more frequently used in aqueous systems. The standard reduction potentials for Li+/Li and Na+/ Na redox couples are approximately −3.04 and −2.71 V vs SHE, respectively. A simple potential diagram comparing the stable potential window (vs SHE) of water and the working potential (vs Li+/Li) of various insertion compounds on the same potential scale is shown in Figure 1. On the basis of this diagram, the feasibility of the electrode materials for the aqueous system can be determined with avoidance of electrolysis.20 11789


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Figure 2. Electrode materials for ARLBs. Electrode materials were categorized as oxide compounds, polyanionic compounds, and other compounds (Prussian blue analogues and organic electrodes).

systems will also be briefly introduced in the last section of this review.

The chemical stability of the electrode materials in H2O is another important aspect to consider in the aqueous electrolyte system. Li et al. discussed the thermodynamic relationship to evaluate the stability of the electrode materials in an aqueous medium on the basis of the following reaction:19 Li(intercalated) + H 2O ⇄ Li+ + OH− + 0.5H 2

3. AQUEOUS RECHARGEABLE LITHIUM BATTERIES 3.1. Introduction to Aqueous Rechargeable Lithium Batteries

(3)

The development of high-performance aqueous rechargeable lithium batteries (ARLBs) has received increasing attention since Li et al. developed ARLBs based on a LiMn2O4 cathode and a VO2(B) anode in 1994.16 Compared to conventional nonaqueous electrolyte LIBs, aqueous electrolyte systems are beneficial in terms of cost, safety, and power capability.15,22,23 The benefits of using cost-effective aqueous electrolytes make ARLBs promising candidates for large-scale ESSs. Moreover, a nonflammable electrolyte can be a critical enabler of safer rechargeable lithium battery systems. ARLBs also provide better power capabilities, since aqueous electrolytes have much higher ionic conductivity than nonaqueous electrolytes. In this section, we review the progress of ARLBs, with a focus on the development of active materials.

They derived two equations expressing the stability of the electrode materials as a function of the pH and Li + concentration (assuming [Li+] = [OH−]): V (x) = 3.885 − 0.118x (vs Li+/Li) (x = pH)

(4)

V (y) = 2.23 − 2kT ln(y) (vs Li+/Li) (y = Li+ concentration)

(5)

V(x) and V(y) are the potentials where hydrogen evolution occurs accompanying the spontaneous delithiation of the electrode material. When the redox potential is higher than V(x) and V(y), the electrode material is expected to be stable in the aqueous solution. For example, Li2Mn2O4 (∼2.97 V vs Li+/ Li) is stable in an aqueous LiOH solution at pH values above 8.19,21 Consequently, materials with adequate working potential and high stability in aqueous solution have been used for electrodes in aqueous alkali-metal ion batteries. Identical to the conventional LIBs and NIBs, ARABs adopted anode materials that have a lower working potential and cathode materials that have a higher working potential. The preferred LIB cathodes, such as LiCoO2, LiMn2O4, and LiFePO4 (3−4 V vs Li+/Li), are promising cathode candidates to meet these requirements of ARABs.15 On the other hand, the conventional carbon anode is not applicable to aqueous systems. This is due to its inherent instability in the aqueous medium at the Li-intercalated state because of the low redox potential. Instead, V-based oxide compounds (e.g., VO2(B), V2O5·nH2O, and H2V3O8) have been mainly considered as potential anode candidates.15 In this review, we introduce aqueous rechargeable batteries with an emphasis on faradic insertion-based electrode materials. Nonfaradic capacitive electrodes (i.e., activated carbon) will not be discussed in this review. The following sections will summarize the progress of ARABs based on intercalating cations such as Li+ ions and Na+ ions, as well as multivalent ions such as Mg2+, Ca2+, Zn2+, and Al3+ ions. Alternative

3.2. Electrodes for Aqueous Rechargeable Lithium Batteries

A variety of electrode materials for ARLBs have been introduced. Unlike the electrode materials used in organic electrolyte systems, the redox potentials of electrode materials in aqueous electrolytes should be within or near the electrolysis potentials of water. The red dotted lines in Figure 2 show the electrolysis potentials that generate O2 and H2 under neutral pH conditions.19 Materials with potentials outside this range cannot function properly since the electrode reaction will continuously involve the water-splitting reaction. We categorized representative types of electrode materials as oxide, polyanionic, and other compounds (as shown with their redox potentials in Figure 2). For both cathodes and anodes, oxides have been most extensively studied for use in ARLBs. Polyanionic compounds and other types, including Prussian blue analogues and organic compounds, have also been investigated. Here, we discuss the detailed electrochemical properties and issues of these electrode materials in ARLBs. 3.2.1. Cathode Materials. Cathode materials for ARLBs should be capable of performing repeated Li+ extraction and insertion. Thus, many of the lithium insertion compounds used in conventional LIBs have been considered. The Li + 11790


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Figure 3. (a) CV curves of LiMn2O4 in aqueous electrolytes with various pH values. Reprinted with permission from ref 39. Copyright 1996 Elsevier. (b) CV profiles of LiMn2O4 in different 0.1 M neutral electrolytes (solid line, KCl; dotted line, NH4Cl; dashed line, LiCl) at a scan rate of 0.05 V s−1. Reprinted from ref 47. Copyright 2003 American Chemical Society. Formal potentials of (c) the first redox couple (Ef1) and (d) the second redox couple (Ef2) at different electrolyte concentrations.

properties were measured using coin-type test cells. This first prototype cell exhibited remarkable reversibility, with an average voltage of 1.5 V and an energy density of 75 W h kg−1, which could compete with those of both Pb−acid and Ni−Cd batteries. The selection of electrolytes with an appropriate pH was important to enable efficient charging of cathode materials. For example, O2 evolution may occur when Li+ is extracted from the LiMn2O4 cathode at high pH, which deteriorates cycle performance. Further studies to understand the effects of pH on electrochemical reactions in the LiMn2O4 cathode were conducted by Pei et al.39 In their work, 5 M LiNO3 solutions with different pH levels (2.0−12.3) were used, and cyclic voltammetry (CV) was performed for tablet LiMn2O4 electrodes in a potential range from −0.5 to +1.4 V (vs the saturated calomel electrode (SCE)), as shown in Figure 3a. At low pH levels, a reduction peak of LiMn2O4 was clearly observed at about 0.5 V, while the oxidation peak slightly overlapped with that of water oxidation. However, at high pH levels, the reduction peak from LiMn2O4 disappeared and the oxidation peak of water significantly increased. This is attributable to the gradually decreasing O2 evolution potential in basic environments, which becomes even lower than the redox potential of LiMn2O4. On the contrary, H+ insertion reaction and Li+/H+ exchange may occur in a low-pH electrolyte,40 as demonstrated by Rao et al.46 These results suggest that appropriate electrolyte pH values are important to reliably operate LiMn2O4 cathodes in ARLBs. Considerable efforts have been devoted to enabling Li+ extraction and insertion into LiMn2O4 during cycling.24,47

deinsertion/insertion potentials of these compounds should be lower than the potential of O2 evolution to ensure the stability of aqueous electrolytes. However, to maximize energy density, it is also important to increase the Li+ deinsertion/insertion potential. With these considerations, various cathode materials have been proposed, including oxides (LiMn2O4, MnO2, LiCoO2, LiNi1/3Co1/3Mn1/3O2, and Na1.16V3O8),13,24−30 polyanionic compounds (LiFePO4, FePO4, LiMnPO4, Li(Fe, Mn)PO4, LiCoPO4, LiNiPO4, and LiCo1/2Ni1/2PO4),31−37 and Prussian blue analogues.38 During the early stages, cathode materials typically showed limited specific capacity and severe capacity degradation upon battery cycling. Studies have revealed that the limited specific capacity is caused by (i) H+ coinsertion into the structure, (ii) Li+/H+ exchange during battery cycling, (iii) water penetration into the structure, and (iv) dissolution of active materials in the aqueous electrolytes.16,19,39−42 Many studies have addressed these issues by modifying cathodes using dopants or additives and controlling the electrode/electrolyte interface using coatings or changing the electrolyte composition.39,43−45 In this section, we introduce the cathode materials that have been reported to date, categorizing them as oxides, polyanionic compounds, and Prussian blue analogues, with an emphasis on reaction mechanisms, electrochemical properties, and ways to improve electrochemical performance. 3.2.1.1. Oxides. 3.2.1.1.1. Spinel LiMn2O4. The first aqueous Li rechargeable battery, developed by Li et al., adopted spinelstructured LiMn2O4 as the cathode using 5 M LiNO3 and 0.001 M LiOH in water as the electrolyte.16 The electrochemical 11791


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Figure 4. (a) CV curves of LiAl0.1Mn1.9O4 obtained at different temperatures for 12 h. (b) CV curves of the electrode derived from calcination at 800 °C for different durations. (c) Cycle performances of LiAlxMn2−xO4 electrodes synthesized at 700 °C for 12 h. Reprinted with permission from ref 44. Copyright 2010 Elsevier. CV curves of (d) LiCr0.15Mn0.185O4 and (e) LiMn2O4 in LiNO3 electrolytes at a scan rate of 10 mV s−1. (f) Discharge capacity retention of Cr-doped and undoped LiMn2O4. Reprinted with permission from ref 49. Copyright 2007 Elsevier.

Jayalakshmi and co-workers demonstrated that Li+ extraction and insertion reactions occur preferentially in neutral electrolytes (LiCl in water) rather than H+ deinsertion/insertion.47 They prepared LiMn2O4 using a low-temperature solution combustion technique and conducted CV measurements using various concentrations of LiCl, KCl, and NH4Cl electrolytes at a voltage window between −1.0 and +1.5 V (vs Ag/AgCl). Figure 3b shows the CV curves for LiMn2O4 in 0.1 M KCl (solid line), LiCl (dashed line), and NH4Cl (dotted line) electrolytes. The first and second oxidation peaks are labeled I and II, and the two reduction peaks are labeled III and IV, respectively. Generally, when the designated cations are intercalated and deintercalated, the formal potentials for the redox couple should follow a linearly proportional relationship with the concentration of intercalants in the electrolyte. Note that the formal potential is given as Ef = (Ea + Ec)/2 (Ea is the anodic peak potential, and Ec is the cathodic peak potential). In all cases, the formal potentials of the first redox couple (I, IV) linearly increased with increases in the concentration of the electrolyte solution, which indicates that the first redox couple was based on the designated cation deinsertion/insertion reactions (Figure 3c). Similarly, the formal potentials of the second couple (II, III) in the LiCl electrolytes shifted to higher voltages with increasing electrolyte concentrations (Figure 3d). However, the formal potentials decreased in electrolytes containing K+ or NH4+, even though the concentration of the electrolyte solutions increased. This may have been because the competitive H+ insertion and extraction reactions occurred more favorably than the K+ and NH4+ reactions. On the basis of their experimental results, they suggested that the Li+ extraction and insertion reactions can take place in LiMn2O4 in neutral LiCl electrolytes, whereas H+ insertion is more favorable than K+ or NH4+ insertion. Tian and Yuan showed that the electrochemical performance of LiMn2O4 could be further improved by modifying the electrolyte composition.43 The electrochemical measurements were performed in a three-electrode system using electrolytes with different concentrations of Li salts; 1 M LiNO3, 5 M

LiNO3, and 9 M LiNO3 aqueous electrolytes with mild acidic conditions (pH 5.34−5.81) were used. They found that the rate capability of the electrode improved with higher concentrations of LiNO3 in the electrolytes because of the higher ionic conductivity. On the other hand, the reversibility of the electrode was best in 5 M LiNO3 and followed the order 5 M LiNO3 > 9 M LiNO3 > 1 M LiNO3 electrolytes. This indicates that the interplay between the electrode and the electrolyte affects the electrochemical performance of ARLBs. Doping metal ions in LiMn2O4 has been evaluated to improve the structural integrity of LiMn2O4 in aqueous electrolytes.44,48−51 Yuan et al. synthesized a series of Aldoped LiAlxMn2−xO4 (x = 0, 0.05, 0.1, 0.15) at various temperatures.44 While distinct redox couples were observed for all samples in the electrochemical cell using 5 M LiNO3, the polarization was generally lower for samples synthesized at lower temperatures (Figure 4a), probably due to the small particle size of LiAlxMn2−xO4. However, the two characteristic redox potentials of LiMn2O4 spinel were more discernible at 800 °C due to the improved crystallinity. It was also confirmed that a longer calcination time results in the two redox couples becoming more distinguishable (Figure 4b). The Al-doped spinel LiAlxMn2−xO4 generally improved cycle stability compared to undoped LiMn2O4 samples. On the basis of galvanostatic measurements at a current rate of 1000 mA g−1 (Figure 4c), the Al-doped samples could deliver a cycle life of more than 4000 cycles, which is not trivial to achieve even in conventional organic electrolyte systems. The improvement of cycle performance with Al doping was attributed to the depressed Jahn−Teller distortion and stabilization of the octahedral sites.44,50 Cvjeticanin et al. proposed Cr-doped LiCr0.15Mn1.85O4 as a potential cathode.49 CV analysis of LiCr0.15Mn1.85O4 was performed in saturated (∼9 M) LiNO3 electrolytes. Two distinct redox peaks were clearly resolved in LiCr0.15Mn1.85O4 in contrast to undoped LiMn2O4 (Figure 4d,e), indicative of a faster response. The LiCr0.15Mn1.85O4 electrode delivered more stable cyclability, which was also 11792


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Figure 5. (a) SEM image of porous LiMn2O4. Inset: TEM image. (b) Cycle performances of the solid and porous LiMn2O4 cathodes. (c) Rate capability of porous LiMn2O4 in ARLB systems. (d) Rate performances of nanostructured LiMn2O4 in ARLB and conventional LIB systems. Reprinted with permission from ref 55. Copyright 2011 The Royal Society of Chemistry. The performances of LiMn2O4 in LIBs have been reported previously (refs 56−58).

insertion in LiMn2O4 is possible, ensuring better cycle performance compared to that of an additive-free system. 3.2.1.1.2. Manganese Dioxide (MnO2). MnO2 possesses various polymorphs, namely, α-, β-, γ-, δ-, ε-, and λ-MnO2 structures.65−67 These polymorphs of MnO2 have been examined for supercapacitor applications in various aqueous solutions, including KOH, K2SO4, Na2SO4, and Li2SO4, which operate on the basis of electric double-layer capacitance.68−71 Among these polymorphs, λ-, γ-, and δ- MnO2 are capable of Li+ insertion into an aqueous solution and could be used as battery electrodes in ARLBs.26,27,72 Deutsher et al. investigated the λ-MnO2 phase for the first time using a saturated LiCl aqueous solution as the electrolyte.26 λ-MnO2 delivered a capacity of ∼160 mA h g−1 in the first discharge, and subsequent Li+ extractions and reinsertions were possible. However, cycle degradation was relatively fast, retaining only ∼50 mA h g−1 after 60 cycles. The electrochemical properties of γ-MnO2 were also examined by Yuan et al.27 While reversible Li+ deinsertion/insertion into γ-MnO2 was possible with LiOH electrolytes, mixed behavior of the insertion and the capacitive surface reaction was observed. Approximately 35 mA h g−1 of the full-cell capacity was delivered from an activated carbon anode and γ-MnO2 cathode, retaining 78% of the initial capacity after 1500 cycles. Similar behavior was observed in δMnO2, as demonstrated by Qu et al. A distinct redox couple was detected in Li+-containing aqueous electrolytes, indicative of lithium insertion into δ-MnO2; however, no cycle data were provided.72 Minakshi et al. attempted to improve the electrochemical properties of MnO2 in ARLBs by introducing several additives into the electrodes, such as TiS2, TiB2, Bi2O3, CaO, and CeO2.73−77 These additives were thought to stabilize MnO2 during battery cycling, depress phase deformation, and reduce side reactions in the electrochemical cells. 3.2.1.1.3. Layered LiCoO2. Layered LiCoO2, which is commonly used in conventional LIBs,78 has been investigated as a cathode in ARLBs. However, initial reports indicated that H+ insertion occurs preferentially over Li+ insertion with LiOH

attributed to depression of Jahn−Teller distortion by Cr doping (Figure 4f). Considerable effort has been devoted to the morphological optimization of LiMn2O4 by nanofabrication.52−55 Qu et al. developed porous LiMn2O4 using a monodispersed polystyrene (PS) colloid.55 The precursor solution containing ethanol, LiNO3, and Mn(NO3)2·4H2O was infiltrated into PS spheres, and the resultant powder was heated to remove the PS template. Figure 5a shows the microstructure of porous LiMn2O4 synthesized as described above. At a current rate of 100 mA g−1, the porous LiMn2O4 delivered 250 mA h g−1 and could retain 93% of the initial capacity after 10000 cycles at a current rate of 1000 mA g−1, making it superior to the microLiMn2O4 (Figure 5b). Even at a very high current rate of 10000 mA g−1, porous LiMn2O4 showed a capacity of ∼90 mA h g−1 (Figure 5c). Notably, the rate capability of porous LiMn2O4 in ARLBs is superior to that of the high-power nanostructured LiMn2O4 in LIBs with nonaqueous electrolytes (Figure 5d).56−58 The electrochemical performance with high power and long cycle stability of nanostructured LiMn2O4 was attributed to (i) a nanostructure that reduces the Li+ diffusion length and increases the electrode/electrolyte surface, (ii) nanograins that compensate the strain resulting from Jahn− Teller distortion, and (iii) inhibition of Mn dissolution in the electrolyte owing to high surface energy.59−61 Furthermore, (iv) the high ionic conductivity of aqueous electrolyte solutions was capable of delivering a high rate capability.25 An interesting approach to improve the cyclability of LiMn2O4 was recently introduced, in which vinylene carbonate (VC) was used as an electrolyte additive.51 The VC additive, which has been widely used in nonaqueous electrolyte systems to form a stable solid electrolyte interphase (SEI) layer on electrode materials,62−64 also forms a stable SEI layer on the surface of the electrode materials in aqueous electrolytes. While the nature of the SEI layer in the aqueous system remains unclear, it may prevent water molecules from penetrating into the electrode materials. As a result, stable Li+ deinsertion/ 11793


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Figure 6. (a) CV analyses and (b) formal potentials of LiCoO2 in different LiNO3 electrolytes. (c) Charge/discharge profiles and (d) cycle performance of LiCoO2 electrodes. Reprinted with permission from ref 20. Copyright 2009 Elsevier. (e) Capacity vs cycle number at different cutoff potentials. (f) Effect of higher cutoff potentials on the efficiency and discharge capacity of LiCoO2. Reprinted with permission from ref 81. Copyright 2011 Elsevier.

Figure 7. (a) Charge/discharge profiles at different current rates and (b) cycle properties of nano-LiCoO2.. Reprinted with permission from ref 82. Copyright 2010 Elsevier.

electrolytes.46,79 It was later reported that reversible Li+ extraction and insertion into layered LiCoO2 is possible using LiNO3 and Li2SO4 electrolytes.20,28,80,81 Ruffo et al. investigated the electrochemical behavior of LiCoO2 with different concentrations of electrolytes, including 0.1, 1.0, and 5.0 M

LiNO3 (Figure 6a).20 The redox reaction was observed at around ∼0.9 V (vs SHE) in the CV measurements of the LiCoO2 electrode. Importantly, the formal potentials increased linearly with the electrolyte concentration, as shown in Figure 6b, indicative of reversible Li+ ion insertion rather than H+ 11794


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Figure 8. CV analyses of NCM electrodes in 1 M Li2SO4 solution with various pH values. Reprinted with permission from ref 42. Copyright 2007 The Electrochemical Society.

capacities of 135 and 133 mA h g−1 at current rates of 5000 and 10000 mA g−1, respectively. This outstanding rate capability may be the result of a reduction in particle size combined with the intrinsically rapid ionic diffusion in aqueous electrolytes. The cycling performance was tested using a two-electrode cell with activated carbon counter electrodes at a current rate of 1000 mA g−1. Relatively small capacity fading was observed over 40 cycles. 3.2.1.1.4. Layered LiNi1/3Co1/3Mn1/3O2. Wang et al. introduced layered LiNi1/3Co1/3Mn1/3O2 (NCM), a promising alternative to LiCoO2 in conventional LIBs, as a possible cathode in ARLB systems. In their work, the effect of the pH on the electrochemical properties was investigated.41,42 Figure 8 shows CV analyses of NCM electrodes in aqueous electrolytes with Li2SO4 at different pH levels. The stability of NCM in aqueous electrolytes was affected by the pH of the electrolyte solutions. An oxidation peak was observed near 0.55 V vs SCE during the first charge process at all pH values, corresponding to the Li+ extraction process. However, subsequent cycling showed a shift in the oxidation peak to higher potentials, particularly at pH 7 and 9, indicative of side reactions. The stability of NCM in the electrochemical cell improved at pH 11, although the shift in the oxidation peak and decrease in peak current were still observed. The NCM electrode was further stabilized at pH 13, where no noticeable changes in the oxidation or reduction peaks were observed. Nevertheless, the O2 evolution potential at this pH (0.75 V vs SCE) began to overlap with the redox potential of NCM; thus, full utilization of Li+ extraction and insertion into the NCM electrode was not possible. The relative instability of NCM under low-pH conditions may be the result of H+ insertion into the NCM electrode. Shivashankaraiah et al. prepared polypyrrole (Ppy)coated NCM and measured its electrochemical properties using LiV3O8 as an anode in 5 M LiNO3 electrolyte.83 Ppy-coated NCM delivered 70 mA h g−1 in the initial discharge and retained a capacity of 55 mA h g−1 after 50 cycles. Liu et al. also

insertion. Galvanostatic cycling measurements were conducted to evaluate the LiCoO2 electrode in ARLB cells with 5 M LiNO3 electrolyte (Figure 6c,d). A flooded three-electrode cell was used, including LiCoO2, a Ag/AgCl reference electrode, and a large LixMn2O4 counter electrode. It was noted that the use of a large counter electrode with an initial composition of Li0.5Mn2O4 allowed the cell to cycle reversibly, whereas the use of other counter electrodes such as stainless steel, nickel mesh, or platinum foil changed the composition of the electrolyte and resulted in low reversibility. The LiCoO2 cathode cycled with relatively low polarization within the voltage window of 0.55− 1.15 V (vs SHE) and showed stable capacity retention over 90 cycles. In a voltage range of 0.55−1.2 V (vs SHE), the obtained capacity was approximately 115 mA h g−1, which corresponded to 0.42 Li+ in LiCoO2. In Figure 6e, Ruffo et al. further investigated the potential range where the LiCoO2 cathode could stably cycle.81 Galvanostatic experiments were conducted at a current rate of 27 mA g−1 with different voltage cutoffs. The electrochemical cell was precycled in a voltage range from 0.55 to 1.2 V (vs SHE), and the upper cutoff voltage increased by 0.05 V. Figure 6e shows that the charge and discharge capacities were affected by the upper cutoff potential. Within a voltage range of 0.55−1.4 V (vs SHE), about 0.5 Li+ could be reversibly extracted from the LiCoO2 cathode, which is close to the maximum practical capacity of LiCoO2. However, the Coulombic efficiency significantly decreased with increasing cutoff potentials, as shown in Figure 6f. The decrease in the reversibility of LiCoO2 above 1.3 V (vs SHE) was attributed to side reactions related to water decomposition and substrate oxidation. Tang et al. proposed a high-performance cathode for ARLBs based on nano-LiCoO2.82 The crystalline nanoparticles were synthesized using a sol−gel method, and the nano-LiCoO2 electrodes were tested using 0.5 M Li2SO4 electrolyte. Figure 7 shows that the nano-LiCoO2 delivered a high capacity of 143 mA h g−1 at a current rate of 1000 mA g−1 and retained 11795


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Figure 9. (a) Schematic of the preparation process of nanoporous NCM. (b) SEM and TEM (inset) images of porous NCM electrodes. (c) Charge/ discharge profiles of porous NCM electrodes. (d) Rate capability of porous NCM and bulk NCM electrodes. (e) Cycling behaviors of porous NCM electrodes. Reprinted with permission from ref 85. Copyright 2013 The Royal Society of Chemistry.

Figure 10. (a) Charge/discharge profiles of LiFePO4 in aqueous electrolytes for 10 cycles. TEM images of LiFePO4 after cycling in aqueous electrolytes (b) with O2 and (c) without O2. (d) Mössbauer spectroscopy of pristine LiFePO4 (upper) and LiFePO4 (below) after 100 cycles. (e) TEM image of carbon-coated LiFePO4. Inset: CV curves of carbon-coated LiFePO4 for 50 cycles. Reprinted with permission from ref 45. Copyright 2011 Elsevier.

Recently, Wang et al. reported a high-performance ARLB based on porous NCM synthesized from a vapor-grown carbon fiber template.85 Precursors of NCM were coated on the modified carbon fibers, and the subsequent calcination process removed the carbon fibers, resulting in an agglomerated form of porous NCM (Figure 9a,b). The nanoporous NCM electrode

reported NCM-based ARLBs paired with a LiV2.9Ni0.05Mn0.05O8 84

anode. While saturated LiNO3 electrolyte with a relatively low pH of 4.9 was used, NCM could deliver a capacity of 103.9 mA h g−1 in the first cycle and maintained 67.5 mA h g−1 after 50 cycles. 11796


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Figure 11. (a) Comparison of CV profiles of LiMnPO4 in different electrolytes. (b) CV analyses of LiMnPO4 and LiTi2(PO4)3 in 5 M LiNO3 electrolyte. (c) Charge/discharge profiles of the LiMnPO4/LiTi2(PO4)3 cell and (d) its cycle performance. Reprinted with permission from ref 89. Copyright 2012 Springer Science + Business Media.

delivered a capacity of 155 mA h g−1 at 1.5 C (1 C = 160 mA g−1) and retained a reversible capacity of 108 mA h g−1 at 45 C (Figure 9c). The cycle stability and power capability were significantly enhanced compared to those of conventional NCM electrodes, as shown in Figure 9d. More than 90% of the initial capacity could be retained after 50 cycles at a discharge rate of 3 C for both 80 and 180 C charging current rates (Figure 9e). 3.2.1.1.5. Na1.16V3O8. Vanadium-based cathode material (Na1.16V3O8, NVO) was recently proposed by Nair et al. for use in ARLB systems.30 The applicability of NVO was examined in a symmetric ARLB system in which both the cathode and anode consisted of NVO. The symmetric cell could be cycled in 4 M LiCl electrolyte, showing a linear charge/discharge profile with an average potential of about 0.8 V (vs SCE) in a voltage range from 0.0 to 1.0 V. The initial capacity was approximately 150 mA h g−1 at a current rate of 5 A g−1. Approximately 75% of the initial capacity was retained after 100 cycles. It was believed that the layered crystal structure of NVO facilitated lithium insertion/deinsertion into and from the layers, while the presence of sodium ions in the layer maintained the layered structure during the electrochemical lithiation/delithiation processes, resulting in excellent electrochemical performance in terms of rate capability and cycle stability. 3.2.1.2. Polyanionic Compounds. 3.2.1.2.1. Olivine LiFePO4. LiFePO4 with an olivine structure was introduced as a cathode for ARLBs by Manickam et al.31 The electrochemical properties of LiFePO4 were tested using a standard threeelectrode system with zinc foil, SCE, and saturated LiOH as the counter electrode, reference electrode, and electrolyte, respectively. Li+ deinsertion/insertion of the LiFePO4 was confirmed on the basis of electrochemical and ex situ analyses; however, the capacity utilizations of LiFePO4 were 41% (70 mA h g−1), 30% (50 mA h g−1), and 20% (40 mA h g−1) for the

first, second, and fifth cycles, respectively.86 The electrochemical oxidation of LiFePO4 was similar to that of nonaqueous media forming FePO4; however, the reduction of FePO4 was not fully reversible since it formed a mixture of LiFePO4 and Fe3O4.31 To elucidate the capacity degradation of LiFePO4, He et al. investigated its chemical and electrochemical stability in ARLB systems.45 As shown in Figure 10a, rapid capacity fading was observed after repeated battery cycling with 0.5 M Li2SO4 (pH 7) electrolyte. They found that O2 and OH− in the aqueous electrolytes critically deteriorated the cycle stability of LiFePO4. Parts b and c of Figure 10 show transmission electron microscopy (TEM) images of LiFePO4 after cycling in the presence and absence of O2 at pH 7. In the presence of O2, floccules were detected on the LiFePO4 surface, while no impurities were observed in the absence of O2. Mössbauer spectroscopy analysis also indicated that the Fe(III)-containing second phase was formed in the active materials after 100 cycles in natural electrolyte solution (Figure 10d). Thus, it was concluded that O2 and OH− result in side reactions and degrade the cycle life of LiFePO4 in ARLBs. They further demonstrated that a carbon coating effectively blocked the attack of dissolved O2 and OH− in the electrolytes and improved the cycle stability of LiFePO4 (Figure 10e). Liu et al. showed that a doped LiFePO4 such as LiMn0.05Ni 0.05Fe0.9PO4 exhibited an enhanced cycle stability and rate capability.87 In an ARLB system paired with a LiTi2(PO4)3 anode using saturated Li2SO4 solution electrolyte, the LiMn0.05Ni0.05Fe0.9PO4 cathode could deliver approximately 87 mA h g−1 of initial capacity with a potential plateau located at 0.92 V at 0.2 mA cm−2. After 50 cycles, about 55 mA h g−1 of capacity was retained. Additives such as CeO2 also improve the cycle stability of electrodes consisting of LiFePO4. Liu et al. reported that 2 wt % CeO2modified LiFePO4 showed better electrochemical performance than pristine materials, especially at elevated temperatures and larger scan rates. Electrochemical impedance spectroscopy tests 11797


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Figure 12. (a) Unit cell of the Prussian blue crystal structure containing an open framework of octagonal hexacyanometalate groups. Each of the eight subcells within the unit cell contains a large open “A site”. Charge and discharge profiles of (b) CuHCF and (c) NiHCF at a current rate of 50 mA g−1. Reprinted with permission from ref 38. Copyright 2011 The Electrochemical Society.

performance was observed over 50 cycles with negligible capacity fading after 5 cycles (Figure 11 c,d). Other studies have attempted to improve the electrochemical performance of LiMnPO 4 by using doping, as in the case of LiNi0.05Mn0.95PO4,90 and by adding additives such as TiS2.91 However, only marginal improvements were reported. 3.2.1.2.4. Other Polyanionic Compounds. Zhao et al. investigated the electrochemical behaviors of a series of LiFexMn1−xPO4 (x = 0.5, 0.4, 0.3, 0.2) compounds in an aqueous electrolyte.92 The carbon composite containing LiFexMn1−xPO4 binary olivine showed two potential plateaus at about 0.95 and 0.37 V (vs SCE), corresponding to the two redox couples of Mn2+/Mn3+ and Fe2+/Fe3+, respectively, in a neutral LiNO3 aqueous electrolyte. LiFexMn1−xPO4 (x = 0.5, 0.4, 0.3, 0.2) electrodes delivered discharge capacities of 110.22, 112.66, 111.08, and 90.50 mA h g−1 at a 0.1 C rate, respectively. Minakshi et al. reported the electrochemical properties of other olivine-structured LiMPO4 (M = Co, Ni) compounds.36,37,93 LiCoPO4 delivered a reversible capacity of 80 mA h g−1 for 25 cycles with discharge and charge potentials at 0.8 and 1.3 V (vs a Sn anode), respectively.36 Ex situ XRD analysis showed that the extraction of lithium proceeded via at least a two-phase reaction, with LiCoPO4 and CoPO4 phases. It was also demonstrated that lithium extraction/insertion was possible for a LiNiPO4 electrode in aqueous LiOH electrolyte.37 The first charge capacity was ∼90 mA h g−1, but the reversible capacity was less than 50%. CV and galvanostatic studies demonstrated that the deinsertion of LiNiPO4 led to an amorphous NiPO4 and a minor product of nickel(II) hydroxide (β-NiOOH). Nevertheless, it is noteworthy that the LiNiPO4 electrode was rarely active in conventional LIBs using organic electrolytes.94 The electrochemical properties of binary LiNi 0.5 Co 0.5 PO 4 olivine have also been studied. 93 LiNi0.5Co0.5PO4 in a potential range of −0.3 to +0.2 V (vs Hg/ HgO) showed an oxidation peak at −98 mV and a corresponding small reduction peak at −231 mV. A peak due to hydrogen evolution was also detected near −0.3 V. Ex situ XRD analysis of a charged sample identified a mixture of two

before and after CV measurements indicated that the CeO2 modification produced a good electrical contact between oxides.88 3.2.1.2.2. Trigonal FePO4. Trigonal FePO4, which is an analogue of olivine, was investigated as a potential cathode for ARLB systems in 2010.34 On the basis of ex situ X-ray diffraction (XRD) analyses, it was revealed that electroreduction of FePO4 formed the olivine LiFePO4. This indicates that the reaction mechanism of FePO4 in aqueous LiOH electrolyte is similar to that in nonaqueous electrolyte. The FePO4 electrode delivered a capacity of 65 mA h g−1 with an average potential of 0.5 V (vs a Zn anode) in the initial discharge; however, sufficient cycle data were not provided. 3.2.1.2.3. Olivine LiMnPO4. The electrochemical properties of olivine LiMnPO4 in ARLBs were investigated by Minakshi and co-workers.35 Using a LiOH electrolyte, electrochemical and ex situ analyses revealed that Li+ extraction and insertion into LiMnPO4 occurs via a reversible two-phase reaction between LiMnPO4 and MnPO4. Unlike LiFePO4, which undergoes a partial reduction, forming a mixture of LiFePO4 and Fe3O4, LiMnPO4 was stable during electrode oxidation and reduction. The LiMnPO4 electrode delivered an initial capacity of 70 mA h g−1, but showed relatively poor cyclability, retaining less than 20% of the initial capacity after 20 cycles. Later, Manjunatha et al. significantly enhanced the cycle stability by examining the effect of electrolytes.89 Figure 11a represents the CV curves of LiMnPO4 in different electrolyte solutions including LiOH, LiNO3, and Li2SO4. In saturated LiOH solutions, a broadened CV profile was caused by the high viscosity and low conductivity of the electrolytes. However, when 5 M LiNO3 was used, the potential gap between redox peaks decreased and the peak current increased. The electrochemical properties of LiMnPO4 were further characterized with LiTi2(PO4)3 as an anode in 5 M LiNO3 electrolyte. The redox potentials of both electrodes were within the stable potential range of the aqueous electrolyte, as shown in Figure 11b. LiMnPO4 delivered a capacity of 85 mA h g−1, corresponding to 0.5 Li+ in LiMnPO4. Moreover, stable cycle 11798


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Figure 13. CV profiles of the VO2(B) electrode in (a) (1) 0.1 M H3BO3 + 3 M KNO3 (pH 9.1) and (2) 0.1 M H3BO3 + 3.95 M LiNO3 (pH 9.1) (in the presence of Li+ in the electrolyte (line 2), the electrochemical activity of VO2(B) is observed at −0.60 and −0.73 V, which indicates that the Li+ deinsertion/insertion from or into VO2(B) occurs at these potentials) and (b) 0.1 M H3BO3 + 3.95 M LiNO3 at pH 6.2 (1), 8.2 (2), and 9.1 (3). (c) Variations of the peak potentials for lithium intercalaction and deinsertion. (d) Capacity retention for discharge (left) and charge (right) of VO2(B). (e) Percentage of VO2(B) dissolved in the electrolyte as a function of the electrolyte pH (0.1 M H3BO3 + 3.95 M LiNO3). (f) Capacity retention of VO2(B) electrodes in 0.1 M H3BO3 + 3.95 M LiNO3 electrolyte at pH 8.2. Reprinted with permission from ref 97. Copyright 1996 The Electrochemical Society.

and Prussian blue analogues. Among various cathode materials, LiMn2O4, LiCoO2, and NCM are the most promising candidates in terms of energy density due to their high operation voltages. However, limitations still remain for largescale ESSs, especially cyclability because of H+ insertion at low pH conditions and dissolution of transition metals during battery operations. The aforementioned problems could be retarded by carbon and organic coating, while further improvements are needed. The olivine-structured LiFePO4 is considered as another prospective cathode because of its high capacity combined with environmental-benign-ness and low costs. Nevertheless, the low voltage and poor cycle performance are not suitable for large-scale energy storage. Especially, residual O2 in the aqueous electrolyte severely deteriorates the cycle performance, decomposing LiFePO4. Other polyanions such as LiMnPO4 and LiCoPO4 delivered good cycle retentions, but the large hysteresis in the charge/discharge process is still problematic. Also, the electrochemical reaction mechanism of those polyanions are unclear yet, and thus, further study is necessary. Finally, while various cathode materials that do not contain Li in the crystal structure, including MnO2, FePO4, and Prussian blue analogues, have been widely studied, and they exhibit remarkable electrochemical performances, it must be noted that those cathode materials without Li are practically unfeasible unless the anode parts are prelithiated.95 3.2.2. Anode Materials. Since Dahn and co-workers constructed the first ARLB,16 various anode materials have been proposed. Materials that have been considered include oxides (VO 2(B),16,19,96−99 spinel Li2 Mn2 O4, layered γLiV3O8,28,100−105 H2V3O8,106 Na1+xV3O8,30,107 paramontroseite VO2,108,109 V2O5,110−112 and anatase TiO2113,114), polyanionic compounds (pyrophosphate TiP2O7 and Na superionic

phases, i.e., parent LiNi0.5Co0.5PO4 and a lithium-extracted Li1−xNi0.5Co0.5PO4 olivine-type new phase, suggestive of a delithiation process in aqueous electrolytes. 3.2.1.3. Prussian Blue Analogues. 3.2.1.3.1. Copper and Nickel Hexacyanoferrates. Recently, Wessells et al. proposed Prussian blue analogues as high-performance electrodes for ARABs.38 Prussian blue analogues have archetypal hexacyanometalate framework structures and can be described using the general formula AxPR(CN)6, in which nitrogen-coordinated transition-metal cations (P) and hexacyanometalate complexes (R(CN)6) form a face-centered cubic framework with large interstitial A sites (Figure 12a). The ionic occupancy in A sites varies between 0 and 2 according to valence changes in either or both of the P and R species. In their work, two electrode materials were prepared using a coprecipitation method: copper hexacyanoferrate (CuHCF) and nickel hexacyanoferrate (NiHCF). These materials were capable of intercalating a variety of ions, including Li+, Na+, K+, and NH4+, in aqueous electrolyte solutions (extraction/insertion of cations other than Li+ will be discussed in sections 4 and 5). The electrochemical activity of CuHCF and NiHCF with Li+ was examined at pH 2 by using three-electrode cells containing 1 M LiNO3 electrolyte. As illustrated in Figure 12b,c, both CuHCF and NiHCF were capable of delivering ∼58 mA h g−1 at 0.83 C (1 C = 60 mA g−1). At higher current rates such as 41.7 C, more than 60% of the capacity obtained at 0.83 C was retained. However, rapid capacity fading was observed for CuHCF and NiHCF electrodes. The poor cycle stability was attributed to dissolution of active materials, which may result from the larger Stokes radius of Li+ (2.4 Å) compared with the A sites (1.6 Å) in the Prussian blue structure.38 In this section, we have introduced representative cathode materials for ARLBs, including oxides, polyanionic compounds, 11799


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Figure 14. (a) Hydrothermally fabricated VO2(B) with morphologies of nanorods, nanoflakes, and nanoflowers (from left to right) and (b) CV curves of nanostructured VO2(B) corresponding to 30 cycles in 2.5 M LiNO3 + 0.01 M LiOH electrolyte. Reprinted with permission from ref 99. Copyright 2011 Elsevier.

conductor (NASICON)-type LiTi2(PO4)3)13,68,115−117 and organic polymeric compounds (polypyrrole and polyimides),118,119 redox reactions of which occur near the potential of hydrogen evolution (Figure 2 in section 3.2). On the basis of early studies, the majority of anode materials showed significant capacity fading during the electrochemical reactions19,39,96,97,100 and limited practical capacity compared to those in conventional LIBs. Investigations of anode materials have revealed the cause of the poor capacity retention: (i) dissolution of active compounds,30,97,110,120,121 (ii) irreversible structural transformation possibly derived from proton insertion,110,114 and (iii) spontaneous deinsertion reaction of the lithiated compounds accompanied by the decomposition of water, resulting in the formation of LiOH(aq) and hydrogen gas (LixA + xH2O → A + xLi+ + xOH− + (x/2)H2↑).19,96,115,116 Recent studies have shown that the electrochemical performance of anodes can be significantly improved by utilizing precisely controlled electrolytes in terms of pH, type and concentration of lithium salts, and residual O2105,122 and by coating the surface of active materials with protective layers.29,110,112 In this section, we briefly introduce anode materials that have been used in the ARLB system. The general electrochemical properties, issues, and challenges of the anode materials and recent progress that has addressed these challenges will be discussed. 3.2.2.1. Oxides. 3.2.2.1.1. Monoclinic VO2(B). The first ARLB by Li et al. used monoclinic structured VO2(B) as the anode material.16 VO2(B) possesses tunnels in the structure, through which lithium ions can rapidly intercalate and deintercalate reversibly.123 In addition, VO2(B) shows a moderate redox potential of −0.67 V compared to SCE (−0.43 V vs SHE, Figure 13a).16,97 Considering that the hydrogen evolution occurs at −0.43 V (vs SHE) at pH 7.29

(see Figure 2 in section 3.2), the thermodynamic stability of lithiated VO2(B) in electrolyte solutions with pH values higher than 7.29 suggests that VO2 can be a proper anode in the ARLB. The redox reaction of VO2(B) involves a lithium insertion/deinsertion reaction (VO2(B) + xLi+ + xe− ⇄ LixVO2(B)), while proton insertion and participation of H+ or OH− ions are not involved. This was demonstrated on the basis of the independence of the redox potential and pH of the electrolyte (Figure 13b,c). Despite the structural and thermodynamic stability of lithiated VO2(B), the ARLB constructed of VO2(B) generally shows poor capacity retention with continuous cycling.96−98 Detailed studies on VO2(B) conducted by Zhang et al. revealed that the cycle stability of VO2(B) is closely related to the pH of the electrolyte.97 Figure 13d shows that the discharge capacity and the Coulombic efficiency (i.e., capacity difference between the discharge and charge processes) of VO2(B) electrodes continue to increase with decreasing pH values from 11.3 to 6.2. This reflects the larger contribution of hydrogen evolution during discharge at lower pH values, which can reduce the stability of the battery. On the other hand, in a strong basic electrolyte (pH 11.3), the capacity decays rapidly due to the dissolution of VO2(B) into the electrolyte, which was revealed using atomic absorption spectrophotometry (AAS) analysis of the electrolytes after 10 cycles (Figure 13e). Zhang et al. showed that the relatively good capacity retention of the VO2(B) electrode could be achieved under mildly basic conditions (pH 8.2) as shown in Figure 13f. Nevertheless, several challenges remain for the utilization of VO2(B) in aqueous lithium battery systems. These include (i) the unwanted hydrogen evolution that reduces the Coulombic efficiency at relatively low pH,97 (ii) the instability of VO2(B) to oxidation over long-term air exposure,96 and (iii) the tedious synthesis procedure of metastable VO2(B).124,125 11800


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Figure 15. (a) CV curves of LiV3O8 in 1 M Li2SO4. (b) Cyclability of a LiV3O8//LiNi0.81Co0.19O2 cell for 1.3 and 1.9 V charge cutoffs. (c) X-ray diffraction patterns of LiV3O8 in the as-prepared state and after the first charge, first discharge, and 100th cycle (from top to bottom). Reprinted with permission from ref 100. Copyright 2000 Elsevier. (d) Color change of the electrolytes from colorless to yellow before (top) and after (bottom) the CV analysis of LiV3O8. Reprinted with permission from ref 120. Copyright 2010 Elsevier. (e) Cyclability of the LiV3O8//LiFePO4 ARLB in the electrolyte with/without the dissolved oxygen, with different concentrations of LiNO3 without the dissolved oxygen, and with 9 M LiNO3 without the dissolved oxygen (from left to right). Reprinted with permission from ref 105. Copyright 2013 Elsevier.

refs 107 and 108 indicated that the redox potential of P-VO2 was ∼ −500−200 mV vs SHE, which was a moderate voltage for anodes above the hydrogen evolution potential (the half-cell behavior of P-VO2 in ARLBs has not been investigated). P-VO2 showed a capacity of ∼65 mA h g−1 at a 10 mA g−1 current rate, and 60% of the maximum capacity could be retained at 200 mA g−1. In addition, it maintained 74% of the initial capacity at the 50th cycle.108 Isostructural VOOH showed electrochemical profiles similar to those of P-VO2, but slightly lower capacity by about 10 mA h g−1 was obtained with a retention of 37% at the 50th cycle.109 3.2.2.1.3. Layered γ-LiV3O8. Kohler et al. first introduced γLiV3O8, a possible cathode material in LIBs,127−129,106 as an anode material in an ARLB using 1 M Li2SO4 or 1 M LiCl electrolyte at pH 6.2.100 The CV profile of LiV3O8 in Figure 15a shows that the redox reaction occurs at an average potential of −100 mV vs SHE, which is higher than the potential of hydrogen evolution at pH 6.2 (−366 mV vs SHE), indicative of the thermodynamic stability of lithiated Li1+xV3O8. Despite their chemical stability, only 15−30% (45−90 mA h g−1) of the capacity achievable in organic electrolytes (230 mA h g−1) could be utilized, and only 25−40% of the initial capacity was retained after 100 cycles (Figure 15b). In 2007−2008, Wang et al. tested the LiV3O8 anode with various cathodes, such as spinel LiMn2O4, layered LiCoO2, and LiNi1/3Co1/3Mn1/3O2.28,101,102,118 They also reported a capacity of 70 mA h g−1 in an aqueous system with relatively low capacity retention below 50% after 100 cycles at neutral pH. Further studies examined the origin of poor capacity retention of γ-LiV3O8 anodes in the ARLB.100,120 Ex situ XRD analysis of the γ-LiV3O8 electrodes at different stages in the electrochemical cycling (Figure 15c) revealed the deterioration of the crystal structure. While the primary structure of γ-LiV3O8 was partly maintained during cycling, it

Many researchers have reported that the electrochemical property of VO2(B) could be enhanced through morphological optimization. Zhang et al. developed a synthesis method of single-crystalline VO2(B) nanoflowers via a hydrothermal route using V(IV)O(acac)2, a prefabricated V4+ precursor compound, and poly(vinylpyrrolidone) (PVP) polymer as a surfactant.98 The electrochemical activity of the nanoflower VO2(B) in an ARLB was examined in comparison with that of micro-VO2(B). More recently, Ni et al. comparatively investigated the relationship between the microstructure and electrode performances using VO2(B) nanorods, nanoflakes, and nanoflowers synthesized from a mixture of NH4VO3 and oxalic acid (Figure 14a).99 According to the authors, the VO2(B) nanoflowers with high surface area showed a higher capacity with better cycle stability compared to the other two morphologies in 2.5 M LiNO3 and 0.01 M LiOH aqueous electrolytes, as shown in Figure 14b. In addition, Wang et al. applied a carbon coating to the surface of VO2(B) to improve the cycle stability.126 3.2.2.1.2. Orthorhombic Paramontroseite VO2 (P-VO2) and Montroseite VOOH. VO2, which has an orthorhombic paramontroseite structure, was examined as a possible candidate for anode materials in ARLBs. It generally exhibits high electronic conductivity, which stems from the metallic electronic structure. Also, the crystal structure contains large tunnels (2.446 Å × 4.946 Å), which can facilitate the diffusion of lithium ions.108,109 In 2008, Wu et al. first developed a synthesis method for P-VO2 via a hydrothermal reaction,108 and Xu and co-workers synthesized P-VO2 from isostructural montroseite VOOH via a topochemical reaction.109 The authors constructed ARLBs with a P-VO2 anode and a LiMn2O4 cathode in 5 M LiNO3 and 0.001 M LiOH electrolytes. Synthetic P-VO2 electrodes could reversibly store and release lithium ions in ARLB systems. The estimation from the voltage−capacity curves of P-VO2//LiMn2O4 cells given in 11801


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Figure 16. (a) CV analysis of V2O5 (circles, boxed) and LiNi1/3Co1/3Mn1/3O2 (squares) in 5 M LiNO3 (pH 11). Reprinted with permission from ref 29. Copyright 2007 The Electrochemical Society. (b) Cyclability and Coulombic efficiency of a cell with bare V2O5 and LiMn2O4 in 5 M LiNO3. (c) X-ray diffraction patterns of the bare V2O5 anode before and after cycling. (d) X-ray diffraction pattern of the PPy-coated LixV2O5 anode after 40 cycles. Inset: SEM image of PPy-coated V2O5. Reprinted with permission from ref 110. Copyright 2007 Elsevier. (e) Cyclability (black) and Coulombic efficiency (red) of V2O5 xerogel vs excess amount of LiMn2O4 in saturated LiNO3. Inset: charge/discharge curves of the 1st, 2nd, and 100th cycles. Reprinted with permission from ref 111. Copyright 2009 Elsevier. (f) Cyclability (black) and Coulombic efficiency (blue) of the PPycoated V2O5−MWCNT hybrid vs LiMn2O4 in 0.5 M Li2SO4. Inset: TEM image of the PPy-coated hybrid. Reprinted with permission from ref 112. Copyright 2012 The Royal Society of Chemistry.

accompanied the formation of new compounds, such as LiV2O5 and V2O5 after 100 cycles.100 In addition, Caballero et al. showed that γ-LiV3O8 is unstable under typical operating conditions of ARLBs using 1 M LiNO3 electrolyte. The yellowcolored electrolyte after cycling, which was not observed in an open-circuit state, reflected the dissolution of the V-based compounds from lithiated γ-Li1+xV3O8 (Figure 15d).120 Recent studies have shown that the electrochemical performance of γ-LiV3O8 anodes could be significantly improved by the modification of the electrode morphology and electrolyte composition. Cheng et al. fabricated a macaroni-like nanostructured Li1+xV3O8 through a solution process using templates. The nanofabrication of Li1+xV3O8 could enable the utilization of 70% of the theoretical capacity, delivering 189 mA h g−1 at a 0.1 C rate.130 Zhao et al. revealed that the elimination of oxygen in the electrolyte and a high concentration of Li salts (9 M LiNO3) significantly increased the capacity retention of an ARLB constructed of γ-LiV3O8 and LiFePO4.105 As described in Figure 15e, the γ-LiV3O8//LiFePO4 cell exhibited approximately 90% capacity retention over 500 cycles without the dissolved oxygen in the electrolyte. These results indicate that the dissolution of γ-LiV3O8 during cycling can be effectively controlled through electrolyte optimization. On the basis of the fact that γ-LiV3O8 shows relatively good cycle stability over hundreds of cycles without the formation of new phases in LIBs with organic electrolytes129,131 and in ARLBs with optimized electrolytes,105 it can be assumed that the irreversible structural evolution of γ-LiV3O8 previously observed in ARLBs may be attributed to compositional changes of active materials induced by the dissolution of vanadiumrelated components during cycling. This suggests that γ-LiV3O8

can be used as a reliable anode material for ARLB systems, when the dissolution problem is settled. 3.2.2.1.4. H2V3O8 and Na1+xV 3O 8. Compounds with structures closely related to γ-LiV3O8 such as H2V3O8106 and Na1+xV3O830,107 were also examined as possible anode materials for ARLBs. H2V3O8, first introduced by Li et al. as an anode in an ARLB,106 possesses higher electronic conductivity due to the mixed valence of vanadium ions in the crystal than that of γLiV3O8. The structural H2O, which holds the VOx layers together with a hydrogen bond, enables structural stability during the lithium insertion and extraction. On the basis of these advantages, Li et al. achieved a large discharge capacity of 234 mA h g−1 with 72% capacity retention after the 50 cycles using hydrothermally synthesized H2V3O8 nanowires. Na1.16V3O8 and Na2V6O16·0.14H2O were recently proposed for use as anodes in ARLBs by Nair et al.30 and Zhou et al.,107 respectively. The replacement of lithium ions in LiV3O8 with sodium ions yielded the sodium vanadate which is a crystallographically “pillared” structure similar to γ-LiV3O8.30 Studies on Na1+xV3O8 in LIBs suggested that Na1+xV3O8 has a more resilient structure, with decreased vanadium dissolution, than layered γ-LiV3O8.132,133 Although no detailed structural analysis has been provided yet on the Na1+xV3O8 with lithium insertion and deinsertion in ARBLs, Na1.6V3O8 exhibited a high capacity of 150 mA h g−1 with relatively stable capacity retention of 75% at the 100th cycle in a symmetric cell with 4 M LiCl electrolyte.30 In addition, the Na2V6O16·0.14H2O nanowire showed a cycle performance of 80.1% at the 100th cycle and 77.0% at the 200th cycle. 3.2.2.1.5. V2O5. Vanadium pentoxide (V2O5) was recently proposed as an anode by Wang et al. in ARLB systems.29,110 V2O5 has been intensively investigated for its potential 11802


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Figure 17. CV and charge/discharge curves of (a) TiP2O7 and (b) LiTi2(PO4)3 in 5 M LiNO3. (c) Cyclability of the TiP2O7//LiMn2O4 cell (black) and LiTi2(PO4)3//LiMn2O4 cell (red) for 25 cycles. (d) X-ray diffraction patterns of the TiP2O7 anode before cycling and after 5 and 25 cycles (from bottom to top). Reprinted with permission from ref 117. Copyright 2007 Elsevier. (e) SEM images of the TiP2O7 anode before (left) and after (right) 10 cycles in 5 M LiNO3. Reprinted with permission from ref 121. Copyright 2014 Elsevier.

the cycles. The formation of various phases was attributed to the simultaneous insertion of lithium ions and protons into TiO2. Later studies of the nanostructured TiO2 anode showed electrochemical behaviors different from those of the reported micro-TiO2 anodes. Nanostructured TiO2 from electrodeposition by Wu et al.136 and nanotube arrays from the anodization by Liu et al.137 exhibited redox reactions at an average potential of −900 mV (vs SCE) in 1 M LiOH electrolyte and mixed 1.5 M LiOH and 4 M KOH electrolyte. The observed redox potential is about 400−500 mV lower compared to the earlier reports on the TiO2 anode.114 In addition, the nanostructured TiO2 anodes showed higher reversibility to Li storage in aqueous electrolytes without proton insertion reactions, resulting in a stable capacity retention during cycling (94.6% at the 100th cycle).136,137 Since it remains unclear how the nanostructured TiO2 avoided the proton insertion reactions and irreversible phase transformations, more detailed investigations are required for TiO2 anodes in ARLB systems. 3.2.2.1.7. Lithium Manganese Oxides. Li et al. and Pei et al. introduced spinel LiMn2O4 as an anode material in ARLBs.39,96 It was shown that the discharge profile of LiMn2O4 possesses a voltage plateau at −100 mV vs SHE with a limited capacity of 0.62 Li+, which results in the formation of Li1.62Mn2O4.3 However, the voltage plateau of a LiMn2O4 anode disappears in subsequent cycles, and the capacity decreases rapidly, indicative of the poor reversibility. On the other hand, Wang et al. constructed an ARLB using defect spinels such as Li2Mn4O9 and Li4Mn5O12 anodes with LiMn2O4 cathodes.125 The Li2Mn4O9//LiMn2O4 cell and Li4Mn5O12//LiMn2O4 cell delivered 100 mA h g−1 of capacity (based on the weight of the LiMn2O4 cathode; the capacity based on defect spinel anodes could not be determined due to the lack of anode loading information in ref 124) at an average potential of 1.0− 1.3 V for five cycles. Since early works in the 1990s,39,96,125 there have been few successful further works on the utilization of lithium manganese oxide anodes in ARLBs. This is probably due to the following intrinsic problems of trivalent manganese ions in reduced

application as a cathode in LIBs due to its high theoretical capacity (294 mA h g−1 for insertion of 2 Li+ ions into V2O5) and high electronic conductivity based on the mixed valence of vanadium during reduction.134,135 In ARLBs, V2O5 exhibits an adequate redox potential as an anode at around −300 mV vs SHE (Figure 16a); however, in early works, it exhibited poor capacity retention (15% after 40 cycles) in the ARLB constructed of a LiMn2O4 cathode and 5 M LiNO3 electrolyte, as shown in Figure 16b.29,110,112 The rapid capacity loss was dominantly caused by dissolution of vanadium ions, accompanied by crystal structure changes and amorphorization during repeated cycling (Figure 16c).110 To address this issue, there have been efforts to inhibit the dissolution of vanadium ions by surface coatings with various materials. Wang et al. demonstrated that V2O5 coated with polymeric compounds such as polypyrrole (PPy) and polyaniline (PAn) could enhance the cyclability (PPy, 86% for 40 cycles; PAn, 80% for 40 cycles) without causing amorphorization after 40 electrochemical cycles (Figure 16d).29,110 Stojkovic et al. showed that the fabrication of V2O5 xerogel effectively improved capacity retention (89% after 100 cycles, Figure 16e).111 Recently, Tang et al. fabricated a hybrid of V2O5 nanowires and multiwalled carbon nanotubes (MWCNTs), which was coated with PPy.112 The hybrid composite could retain the initial capacity with negligible loss over 500 cycles in 0.5 M Li2SO4 electrolyte (Figure 16f). 3.2.2.1.6. Anatase TiO2. Anatase TiO2, a stable anode material for LIBs, was first applied to ARLBs by Minakshi et al.114 The CV profile of TiO2 in 5 M LiOH electrolyte showed two cathodic peaks at −400 and −500 mV (vs SCE) and an anodic peak at −200 mV, indicative of the copresence of reversible and irreversible reactions during cycling. The authors also observed phase transformation of TiO2 during the reduction process, resulting in the formation of various titanium oxides, such as LixTiO2, Ti2O3, Ti2O, and TiO, from ex situ XRD, X-ray photoelectron spectroscopy (XPS), and Fourier transform infrared (FTIR) analyses. Although LixTiO2 (formed through lithium insertion into TiO2) disappeared reversibly after the charge process, the other phases remained throughout 11803


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Figure 18. (a) TEM images of LiTi2(PO4)3 uniformly coated with a carbon layer (∼25 nm thick) via chemical vapor deposition. (b) CV curves of pristine LiTi2(PO4)3 (top) and carbon-coated LiTi2(PO4)3 (bottom). (c) Cyclability and Coulombic efficiency of the C-coated LiTi2(PO4)3// LiMn2O4 cell in 1 M Li2SO4. Reprinted with permission from ref 13. Copyright 2007 Wiley-VCH Verbg GmbH & Co. KGaA. (d) Cyclability and Coulombic efficiency of LiTi2(PO4)3//partially delithiated Li1−xFePO4 in excess as a lithium reservoir in 2 M Li2SO4 at constant pH 8. Reprinted with permission from ref 115. Copyright 2011 The Electrochemical Society.

absence of the oxidative water decomposition by using a Li1−xFePO4 cathode, the LiTi2(PO4)3//Li1−xFePO4 cell in 2 M Li2SO4 at constant pH 8 resulted in a Coulombic efficiency of greater than 98% (Figure 18d) and retained 90% capacity after 100 cycles of repeated charge and discharge. 3.2.2.3. Organic Compounds. Electroactive organic-based molecular and polymeric compounds, which can store lithium ions and electrons in their conjugated chemical bonds, have attracted recent interest as sustainable electrode materials in LIBs with organic electrolytes.139−145 Lately, several attempts to utilize the organic electrode materials in aqueous lithium battery systems have been reported in an effort to construct “greener” aqueous batteries. The organic materials have been mainly applied to anodes in aqueous lithium batteries due to their redox voltage typically around −0.8 to +0.3 V vs SHE (2.2−3.3 V vs Li+/Li) and the absence of lithium ions in pristine states. The first demonstration of organic anodes in ARLBs was done for PPy paired with a spinel LiMn2O4 cathode in saturated Li2SO4 electrolytes (pH 7).118 Lithium ions could be reversibly stored and released in and out of the PPy anode through a “doping and dedoping” mechanism at an average potential of −0.27 V vs SCE. The same author constructed an ARLB with a PPy anode and layered LiCoO2 cathode.146 While the capacity of PPy reached only 40% of the practical capacity attained in organic electrolytes, the capacity was stably maintained over 120 cycles. Recently, polyimide with conjugated carbonyl groups based on the 1,4,5,8-naphthalenetetracarboxylic dianhydride (NTCDA) moiety was introduced as an anode material in ARLBs by Qin et al.119 The energy storage mechanism in polyimide is known as “enolization”, which represents the storage of lithium ions during reduction with the charge redistribution within conjugated carbonyl groups ((CO)2n) in aromatic molecules (Figure 19a). The polyimide anodes

states: (i) large distortion of the host frameworks induced by the Jahn−Teller active Mn3+ ion, which is a common cause of capacity fading in LIBs as well,39 (ii) dissolution of Mn2+ ions into electrolytes via disproportion of Mn3+ ions,138 and (iii) the relatively low theoretical capacity (95 (1000) 88 (250) 88 (1000) 99.8 (1000) 75 (20)

Specific capacities are calculated with the weight of this electrode. bCapacity based on total electrodes including both cathode and anode.

coatings on the electrode or stabilization of the electrode surface using additives may prevent side reactions, as has been shown previously.43−45 However, only a few of the electrode materials that have been introduced here are chemically stable without surface treatments. Therefore, it is important to identify electrode materials with appropriate redox potentials to attain chemical stability during battery operation.

operation in aqueous battery systems will be briefly introduced, and approaches to resolving these problems will be discussed. 7.1. Side Reactions with H2O and O2 in an Electrolyte

The stability of intercalated Li in the host electrode is an important issue in an aqueous electrolyte system. Side reactions between intercalated Li and H2O are detrimental to the cycle stability of aqueous cells. The thermodynamic relationship is well described in eqs 3−5 provided in section 2. Undesirable reactions can occur in electrode materials when they have a redox voltage lower than V(x), as shown in eqs 3 and 4.19 The majority of positive electrodes are stable against this reaction because their redox potentials are generally higher than V(x) (2.626 V vs Li+/Li in neutral solution).122 However, the chemical stability of the anode is more vulnerable because anode redox potentials are typically located near V(x). For example, the redox potential of the most common anode, LiTi2(PO4)3/Li1+xTi2(PO4)3, is 2.45 V vs Li+/Li.117 Thus, the lithiated phase is theoretically unstable in pH 7 aqueous electrolyte. While Li1+xTi2(PO4)3 may be stable in aqueous solutions with pH values above 10, unwanted side reactions can occur at the cathode in alkaline media, such as dissolution of the electrode (e.g., LiFePO4).45 The chemical stability of electrode materials further decreases when dissolved O2 is present in the electrolyte. Many reports described ARAB operation in air, which contains 21% oxygen. This significantly increases V(x) through the following reaction. Details on the derivation of this equation can be found in previous reports.122

7.2. H2 and O2 Evolution from Electrolyte Decomposition

The evolution of H2/O2 from aqueous electrolytes is an important factor that should be considered when designing for a long-term cycle life. Contrary to organic electrolytes with wide electrochemical stability windows (>4.0 V), the stability range of aqueous electrolytes is relatively narrow (∼1.23 V). Beyond this window, aqueous electrolytes are prone to decompose, generating H2 and O2 gases. Kinetic limitations of this electrolysis may expand the stability limit of some aqueous electrolytes to ∼2.0 V. However, side reactions derived from the electrolysis may inevitably occur during long-term cycling. Decomposition of organic electrolytes also occurs in LIBs, which can partly contribute to stable cycling through the formation of a protective film, a so-called SEI layer, on the electrode surface. However, the decomposition of aqueous electrolytes, and the consequent generation of O2 and H2 gases, does not form any protective films on the electrode surface. As a result, it is important to control the operating potential of both cathodes and anodes in aqueous rechargeable batteries. The catalytic effect of the active electrode material itself on water electrolysis is another factor that should be considered. The electrolysis overpotential of water generally depends on the relative activity of water-splitting catalysts. To minimize side reactions resulting in H2/O2 evolution, a noncatalytic electrode must be used in ARAB systems. For example, some electrode materials such as delithiated spinel LiMn2O4 (λMnO2) and Li2MnP2O7 are reportedly water oxidation catalysts.247,248 This indicates that the delithiation of spinel LiMn2O4 may promote O2 evolution, which in turn deteriorates ARLB cycle performance.

Li(intercalated) + 0.25O2 + 0.5H 2O ⇄ Li+ + OH−

V (x) = 4.268 − 0.059(pH) (V vs Li+/Li)

Notably, in the presence of O2, none of the materials introduced in this review can be used as anodes for ARLBs regardless of the electrolyte pH (e.g., V(x) = 3.855 V vs Li+/Li in a pH 7 electrolyte, and V(x) = 3.442 V vs Li+/Li at pH 14). In addition, these side reactions may occur in ARSBs at similar voltages. Taken together, precise control of the cutoff voltage at a certain pH and elimination of residual O2 in the electrolyte are crucial for stable operation of aqueous batteries. Carbon

7.3. Electrode Dissolution in an Aqueous Electrolyte

Electrode dissolution in water-based electrolytes can occur in ARABs and limit long-term cyclability. The dissolution of cathodes at a specific pH has often been noted in previous 11821


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reports. For example, LiFePO4 decomposes to Fe3O4 under high pH conditions (∼pH 10),31 and Prussian blue analogues are not generally stable at pH 7.38 In addition, dissolution may be accelerated with nanostructured electrodes due to their large surface areas. To resolve the problem of electrode dissolution during battery cycling, surface protection via the application of coatings or the addition of surface stabilizers should be considered.43−45 7.4. Proton Co-Insertion with Guest Ions

Co-insertion of protons with guest ions into the host electrode has been observed in ARABs. This is commonly observed in layered structured materials such as LiCoO 2 and MnO2.81,169,173 A recent first principles study demonstrated that substitutional protons blocked Li diffusion pathways, reducing the cyclability of layered-type electrodes in aqueous solutions.249 Proton co-insertion can be reduced by increasing the pH. For example, LiCoO2 electrodes contain significant numbers of protons in the lattice after cycling at low pH, but lithium ion deinsertion/insertion is stable at pH values above 9 without proton co-insertion.122 It should be noted, as described above, that battery operation in alkaline electrolytes can result in side reactions such as dissolution of the electrode. Since using neutral aqueous electrolytes (pH 7) is advantageous in ARABs, the development of electrode materials with minimal proton deinsertion/insertion is essential.

Haegyeom Kim received a B.S. degree (2009) from the Department of Materials Science and Engineering of Hanyang University and an M.S. degree (2011) from the Department of Energy, Environment, Water and Sustainability of the Korea Advanced Institute of Science and Technology. He is currently studying the development of graphenebased electrode materials for energy storage devices such as Li rechargeable batteries, Na rechargeable batteries, and supercapacitors as a Ph.D. candidate in the Department of Materials Science and Engineering of Seoul National University.

7.5. Outlook

The utilization of aqueous electrolytes in rechargeable Li and Na batteries helps in meeting the essential requirements of large-scale ESSs, especially in terms of low cost and environmental sustainability. “Rocking-chair”-type ARLBs and ARSBs possess promising characteristics such as high round trip efficiency, high energy density, and controllable power and energy that can be tailored to meet different requirements due to the variety of chemistries available. While current ARLBs and ARSBs require further improvements in cycle stability, some reports describe improved capacity retention by optimizing battery operation conditions. Moreover, recent studies have proposed that host materials with an “open framework” or organic materials are capable of reversibly taking up and releasing guest ions in their flexible structure via rearrangement of chemical bonds, thereby significantly improving the stability of ARLB or ARSB systems. Several alternative ARAB system designs are also available, as discussed above, to address the current challenges. These possibilities include ARABs incorporating (i) multivalent ions, (ii) oxygen gas electrodes, and (iii) hybrid electrolytes. On the basis of recent advances and further developments expected in the near future, ARABs are predicted to play an important role in grid-scale stationary ESSs connected to renewable power plants, which will promote a “greener” future.

Jihyun Hong received a B.S. degree (2009) and an M.S. degree (2011) in materials science and engineering from the Korea Advanced Institute of Science and Technology. His focus is on high-capacity layered Li-excess transition-metal oxides and biological organic compounds involving cellular energy transduction as cathodes for sustainable Li ion batteries using combined experiments and DFT calculations as a Ph.D. candidate in the Department of Materials Science and Engineering at Seoul National University.

AUTHOR INFORMATION Corresponding Author

*Phone: +82-2-880-7088. Fax: +82-2-885-9671. E-mail: [email protected]. Author Contributions ∥ Ha.K., J.H., K.-Y.P., and Hy.K. contributed equally to this work.

Notes

The authors declare no competing financial interest. 11822


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Kyu-Young Park received a B.S. degree (2010) from the Department of Materials Science and Engineering and an M.S. degree (2012) from the Department of Energy, Environment, Water and Sustainability of the Korea Advanced Institute of Science and Technology. He is currently a Ph.D. candidate in the Department of Materials Science and Engineering of Seoul National University. He is studying the polyanion electrode for Li ion batteries using combined structural analysis and ab initio calculations.

Kisuk Kang is a professor of materials science and engineering at Seoul National University (SNU), where he received his B.S. He did his Ph.D. and postdoctoal studies at the Massachusetts Institute of Technology. Before he joined SNU, he was a professor at the Korea Advanced Institute of Science and Technology until 2010. Since 2013, he has been a tenured professor at SNU. His research laboratory at SNU focuses on developing new materials for Li ion batteries and post Li battery chemistriessuch as Na, Mg, and metal−air batteries using combined experiments and ab initio calculations. Over the past 5 years, he has published papers in more than 100 international journals and issued more than 10 patents in this field. His accomplishments have been recognized by the International Society of Electrochemistry Young Investigator Award (2011), Inaugural Energy and Environmental Science Lectureship Award (Royal Society of Chemistry, 2012), PBFC Award (Korean Electrochemical Society, 2013), and prestigious Korean Young Scientist Award by the President of Korea (2013).

Hyungsub Kim received a B.S. degree in materials science and engineering from Hanyang University and an M.S. degree from the Korea Advanced Institute of Science and Technology. He is currently a Ph.D. candidate in the Department of Materials Science and Engineering of Seoul National University. He is working on the

ACKNOWLEDGMENTS This work was supported by (i) IBS-006-D1, (ii) Human Resources Development program (20124010203320) of the Korea Institute of Energy Technology Evaluation and Planning (KETEP) grant funded by the Korea government Ministry of Trade, Industry and Energy, (iii) the National Research Foundation of Korea Grant funded by the Korean Government (MEST) (NRF-2009-0094219), (iv) the Energy Efficiency & Resources of the Korea Institute of Energy Technology Evaluation and Planning (KETEP) grant funded by the Korea government Ministry of Trade, Industry & Energy (MOTIE) (No.20132020000270), (v) This work was supported by a grant from the Nuclear Research & Development Program of Korea Science and Engineering Foundation (KOSEF) funded by the Ministry of Education, Science and Technology, Republic of Korea. (vi) This work was also supported from the World Premier Materials grant funded by the Korean Ministry of Knowledge Economy.

design of an open framework electrode for Na rechargeable batteries.

Sung-Wook Kim received a Ph.D. degree (2010) from the Department of Materials Science and Engineering of the Korea Advanced Institute

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