| INTRODUCTION

The lithium-ion battery (LIB) has become an indispensable energy storage solution. However, it is still challenging to balance the cost, sustainability, and performance, where the cathode is often considered as the bottleneck. 1 To enable the long-term deployment of LIBs at scale, reducing and eliminating Ni and Co in cathodes is crucial due to the cost and supply chain issues associated with their limited and geographically concentrated supply. 2 Generally, applying earth-abundant materials, such as Fe or Mn, in said cathodes has the potential to overcome these materials supply challenges. Among the materials utilizing these elements, MnO 2 has one of the deepest histories of application as an electrode material, dating back to 1866. 3 In addition to its economic and environmental advantages, the high theoretical specific capacity (308 mAh/g, based on one-electron reaction), diverse polymorphs and chemical tunability make MnO 2 particularly desirable. Endowed by these merits, MnO 2 was commercialized as primary batteries utilizing electrolytic manganese dioxide (EMD) cathodes in the 1970s, which are now consumed at scales in the hundreds of millions of cells. 4 However, the commercial success of rechargeable MnO 2 cells remains limited due to issues associated with their cyclability.

Historically, MnO 2 , (i.e., pyrolusite MnO 2 ) was explored in secondary batteries to explore the benefits of nanostructures in the 2000s. Before that, this phase was recognized as electrochemically inactive in the bulk form, due to its sluggish lithium diffusion and rapid irreversible transformation to spinel LiMn 2 O 4 (LMO). 5 However, despite the impact of such nanostructures, the reported cycling stability of such materials remained untenable for commercial means. This behavior is thought to be largely a result of proton exposure, including the hydrolysis of PF 6

-and solvent (ethylene carbonate, EC) deprotonation at the cathode/electrolyte interface. 6 For example, Bruce et al. reported mesoporous β-MnO 2 of 10 nm wall thickness delivered a first cycle capacity of 292 mAh/g while only 50 charge/discharge cycles (53% capacity retention) were reported. 7 Though MnO 2 is known to suffer acutely from these concerns, such degradation is broadly observed in Mn-based cathodes. For example, it is a consensus that LMO suffers from unwanted Mn 2+ dissolution and performance decay in carbonate-based electrolytes. 8 Acid species generated in commercial carbonate electrolytes can accelerate the disproportionation reaction of Mn 3+ , forming soluble Mn 2+ . Different strategies have been implemented to stabilize cathode/electrolyte interface to reduce such dissolution, including surface coating, doping, electrolytes additives, and functionalized binders. 9 However, fundamentally suppressing Mn dissolution in sustainable Mn-based cathodes remains a grand challenge. 8 It is proposed that electrolyte redesign is a promising approach to minimize Mn dissolution. Ideally, the electrolytes should be acid-free during storage and electrochemical cycling. Thus, ether-electrolytes utilizing lithium imide salts with limited tendency of hydrolysis become candidates. The other enlightening fact is that the operation voltage of MnO 2 is below 4 V (vs. Li + /Li), which well matches the stability window of ether solvents. Our rationale for paring MnO 2 cathode and ether electrolytes as more sustainable cell chemistry is presented in Schematic 1.

Herein, we report that the application of an ordinary ether electrolyte (1 mol/L LiTFSI DOL/DME 50/50 in volume) in MnO 2 secondary cells vastly improves cycling performance via the suppression of Mn dissolution. We find that this stabilization is both a function of replacement of LiPF 6 in the electrolyte with LiTFSI, and the replacement of carbonate solvents with ether solvents. Application of this ether electrolyte was found to provide > 90% capacity retention over 200 cycles in a wide variety of MnO 2 polymorphs in both nano-scale and micron-scale morphologies, indicative of the versatility of this strategy. While X-ray adsorption near edge spectroscopy (XANES) spectra shows direct evidence of Mn 3+ in cycled MnO 2 , Mn dissolution is minimized in the ether system. This study not only offers a new strategy to enable the reversible cycling of Li-MnO 2 batteries, but also enlightens further examination of the role of electrolyte chemistry in the long-term stabilization of Mn-based cathodes.

| RESULTS AND DISCUSSIONS

Though many MnO 2 polymorphs have been reported, β-MnO 2 was initially chosen as a model material for its one-phase reaction mechanism (after initial cycles) at the nanoscale offering desirable structural integrity. 7a More importantly, the phase with tunnel homogeneity (1 × 1) is free of guest ions, including protons that are potentially detrimental to interfacial stability of cathodes. 10 In this work, a unique nanotubes-assembled hierarchical β-MnO 2 microsphere (denoted as h-β-MnO 2 ) was designed by a rapid thermal treatment (400°C, 4 h) from γ-MnO 2 as depicted in Schematic 1. Upon heating, structural water was continuously removed until above 300°C, shown by thermogravimetric analysis (TGA) in Figure 1A. The removal of crystalline water inside the tunnel structure was observed along with the shrinkage of the lattice parameter (Figure S1A). The subsequent γ-β phase transition (Figure S1B) occurred between 350°C and 400°C, is in good agreement with previous work. 11 As shown in Figure 1B, the h-β-MnO 2 of tetragonal rutile structure was identified as space group P4 2 /mnm. The material presents a Brunauer-Emmett-Teller (BET) specific surface area of approximately 27 m 2 /g with limited mesopores (Figure 1C).

To examine the structure properties, electron microscopic imaging including scanning electron microscope (SEM) and transmission electron microscopy (TEM) were performed. From the SEM images (Figure 1D), the h-β-MnO 2 presents particle sizes between 2 μm and 4 μm with a densely packed structure suggested by lowmagnification TEM images. The secondary structure is composed of nanotubes (Figure 1E). Moreover, small tube diameter (30-40 nm) has been demonstrated to accommodate large strain induced by ion insertion. 12 Unlike traditional chemical manganese dioxide (CMD), which tends to form dispersed nanotubes or nanowires, we believe that this hierarchical structure would produce a higher volumetric energy density for battery applications. 13 Meanwhile, from the high-resolution TEM (HRTEM) image (Figure 1F), ( 111) and (101) facets can be differentiated. The inset fast Fourier transform (FFT) pattern clearly shows (211) and (101) facets. We note that our h-β-MnO 2 is not single-crystalline, in contrast to nanosized β-MnO 2 directly obtained by hydrothermal reaction. 14 To demonstrate the universality of our electrolyte selection criteria, we also included other CMD, α-MnO 2 , γ-MnO 2 (Figures S1-S3), and thermally activated commercial EMD powder (Figure S4).

To examine the impact of electrolyte selection on the electrochemical behavior of MnO 2 , 1 mol/L LiPF 6 in EC: diethyl carbonate (DEC) (50:50 by volume, LP40) was chosen as a common commercial carbonate electrolyte to compare with the aforementioned ether system. Before examining their influence on secondary battery cycling, the primary battery performance was assessed (Figure S5), as electrolyte composition has been demonstrated to have large impact on the first discharge behavior. 15 In doing so, limited difference between carbonate and ether-based electrolytes were observed on the first discharge. At C/20, approximately 250 mAh/g was delivered with a discharge plateau at approximately 2.75 V. 7a When the mass loading was increased over 10 mg/cm 2 (~2.6 mAh/cm 2 ), negligible capacity drop (Figure S6) was detected. 16 Therefore, it is confirmed that the synthesized h-β-MnO 2 with low crystallinity displays sufficient discharge capacity and output voltage to act as a platform for future study into its rechargeability.

To avoid the parasitic oxidation of ether solvents (Figure S7), the operation window was initially set as 1.8 V-4.0 V. The cycling performance for both electrolytes is exhibited in Figure 2A,B. Though the initial specific capacities were comparable, the capacity loss per cycle between the ether and carbonate systems became distinct during cycling. After 100 cycles, h-β-MnO 2 cycling in ether maintained 144 mAh/g with a 78% capacity retention at C/20, compared to only 69 mAh/g (capacity retention of 38%) in carbonate. When the current density was increased to 1C, the electrode demonstrated 82% capacity retention after 300 cycles in ether, as opposed to 11% in carbonate. Voltage profiles at 1C in Figure 2C,D shows remarkable capacity fading and polarization in carbonate. To consolidate the universal influence of different electrolyte chemistry, other MnO 2 were synthesized and applied to long-term cycling (Figure S8), delineated in Figure 2E. Regardless of the phase, the ether electrolyte was found to preserve > 90% capacity retention after 50 cycles, whereas the carbonate system produced unstable, highly phase-dependent cycling. Impedance data in Figure S9 reveals that in carbonate, both R sei and R ct were altered considerably, likely due to the increased passivation layer and structure changes on their electrodes. 17 Figures 2F and S10A suggest that the degradation produced by the LP40 system is due to both the solvent and LiPF 6 , where a LiTFSI in EC/DEC analog produced slightly improved performance relative to LP40, yet less stable than the ether system. Meanwhile, we measured ex-situ XRD (Figure S10B) for cycled h-β-MnO 2 three electrolytes, where the negligible difference was noticed. These data provide clear evidence that the electrolyte chemistry governs the cyclability of MnO 2 secondary batteries.

The instability of cycled MnO 2 in carbonate is conjectured to be a result from Mn dissolution and accelerated disproportionation reaction of Mn 3+ in an acidic environment. To provide direct evidence of Mn 3+ in cycled MnO 2 , we employed transmission X-ray microscopy (TXM) with XANES to observe the changes of chemical states of Mn during cycling. Figure 3A suggests that the valence status of Mn is on average is close to 3+ at the discharged state. Figure 3B reflects the spatial distribution of Mn 3+ at an area scale of approximately 10 μm 2 . Only a trace amount of Mn 4+ is detected, consistent with the high capacity during the first discharge. Abundant Mn 3+ offers a foundation for studying the effects of different electrolytes. Because pristine MnO 2 is free of Mn 3+ and the cycled electrodes contain a limited amount of active material, we used LiMn 2 O 4 powder containing 50% Mn 3+ in a soak test under elevated temperature. This experiment directly reflects the stability of Mn 3+ in carbonate and ether (Figure 3C). We observed no change of powder and electrolytes in ether. In contrast, Mn dissolution happens in carbonate. The chemical state of dissolved Mn is reported to be 2+. 18 During electrochemical cycling, dissolved Mn 2+ can migrate through electrolytes and deposit on Li metal. Therefore, we detached cycled Li-MnO 2 cells and characterized the Li metal anode. As expected, in carbonate, Mn dissolution and deposition process existed more significantly (Figures 3D and S11A,B). 8,19 We also observed a temporal accumulation of Mn on the Li metal. In contrast, no Mn signal was detected with the usage of the ether electrolyte (Figures 3E and S11C,D).

Such difference was also substantiated by other MnO 2 samples such as electrolytic manganese dioxide (EMD), as displayed in Figure S12A-D. It is observed that 1 mol/L LiTFSI EC/DEC also resulted in Mn dissolution (Figure S12E,F), which can explain the notable capacity fading. To summarize, our hypothesis that the ether system without proton generation can resolve Mn dissolution fundamentally in LIBs is supported. More discussion is exhibited in the supplementary discussion and Table S1. At the same time, optical images (insets of Figure 3D,E) manifest that separator in carbonate appeared in a brownish color, a sign of electrolytes decomposition triggered by Mn 2+ . 20 We also found that reducing the charging cut-off voltage from 4.0 V to 3.7 V resulted in much-improved cycling stability possibly due to the limited volume expansion and irreversible phase transformation (Figure S13). Therefore, we focused on h-β-MnO 2 being charged to 3.7 V. In terms of cyclability (Figure S14), after 200 cycles, 130 mAh/g and 120 mA/g could be delivered at 0.2 C and 0.5 C, corresponding to capacity retention of 83% and 86%, respectively. β-MnO 2 of 12 h-annealing also demonstrated excellent reversibility (Figure S15) with almost no capacity fading after 300 cycles. All the above data convince that MnO 2 cathode can be highly rechargeable.

To provide a more holistic perspective of the benefits of the ether electrolyte in rechargeable Li-MnO 2 batteries, we apply commercial EMD as the cathode to evaluate cycling stability and the rate performance. At 0.2 C, the EMD presented 130 mAh/g after 200 cycles (96% retention) when applied with the ether electrolyte (Figure 4A). The voltage profile in Figure 4B process (Figure 4C). Ex-situ XRD reveals minimal irreversible lattice expansion along with amorphization during long cycling occurs (Figures 4D and S16). This can be attributed to the highly disordered spinel-related phase. 21 At 1C, 2C, and 3C, 125 mAh/g, 119 mAh/g, and 110 mAh/g could be output, respectively. Even cycled at 60°C at C/2 (Figure 4F), the EMD cathode still shows a capacity retention of >90% after 50 cycles in the ether electrolyte, suggesting high electrochemical stability.

Therefore, we propose a viable strategy to eliminate Mn dissolution in LIBs based on the application of ether electrolytes, which utilize neither PF 6 -nor EC that has been historically applied in Li-MnO 2 batteries (Table S2).

Figures 5 and S18 show our improved cyclability of MnO 2 as well as the promise of the ether electrolyte in Li-MnO 2 cells. In addition, in Figure S19, the cost per kWh for various cathodes were shown. Given the desirable technological maturity in synthesis and modification of MnO 2 , there are abundant opportunities for further increasing the cycling life and energy density. For example, Li 0.33 MnO 2 synthesized from EMD by LiNO 3 pre-lithiation indicates no dissolution in ether and excellent cyclability (Figure S20).

| CONCLUSION

In summary, we demonstrated that common ether electrolytes can repurpose traditional primary Li-MnO 2 batteries to be rechargeable batteries while preserving the low-cost feature. Regardless of the phase or morphology, we found that while Mn dissolution from MnO 2 cathodes is a universal phenomenon with carbonate electrolytes, ether-based electrolytes naturally without proton generation can prevent Mn-based cathodes from dissolution and rapid cell failure. As a result, MnO 2 cathode is exposed to be highly rechargeable as an intercalative cathode. For example, the 1 mol/L LiTFSI DOL/DME (50/50 in volume) ether electrolyte, paired with commercial EMD, was able to deliver a high capacity retention (96%) at 0.2 C for over 200 cycles. In the future, costeffective MnO 2 as well as further engineered MnO 2 can be more authentically characterized and understood by applying ether electrolytes. Such a simple yet effective strategy toward electrolyte design can pave the way for developing low-cost and more sustainable rechargeable batteries compared with today's LIBs systems. S2 presents more information of the prior studies.