Recent Progress in Cathode-Free Zinc Electrolytic MnO₂ Batteries: Electrolytes and Electrodes

Abstract

Zinc–manganese dioxide (Zn–MnO₂) batteries, pivotal in primary energy storage, face challenges in rechargeability due to cathode dissolution and anode corrosion. This review summarizes cathode-free designs using pH-optimized electrolytes and modified electrodes/current collectors. For electrolytes, while acidic systems with additives (PVP, HAc) enhance ion transport, dual-electrolyte configurations (ion-selective membranes/hydrogels) reduce Zn corrosion. Near-neutral strategies utilize nanomicelles/complexing agents to regulate MnO₂ deposition. Moreover, mediators (I⁻, Br⁻, Cr³⁺) reactivate MnO₂ but require shuttle-effect control. For the electrodes/current collectors, electrode innovations including SEI/CEI layers and surfactant-driven phase tuning are introduced. Electrode-free designs and integrated “supercapattery” systems combining supercapacitors with Zn–MnO₂/I₂ chemistries are also discussed. This review highlights electrolyte–electrode synergy and hybrid device potential, paving the way for sustainable, high-performance Zn–MnO₂ systems.

1. Introduction

Zinc–manganese dioxide (MnO₂) batteries continue to dominate the primary battery market, making efforts to render them rechargeable highly promising [1]. However, rechargeable alkaline Zn–MnO₂ batteries have historically suffered from poor cycling performance, prompting recent research to focus on near-mild acidic or neutral electrolytes [1,2]. Although the first report of a Zn//γ-MnO₂ battery using a mild acidic ZnSO₄ electrolyte dates back to 1988 [3], widespread attention emerged only in 2012 when Zn//α-MnO₂ demonstrated excellent reversibility and cyclability, as illustrated in Figure 1a [4]. Initially, the electrochemical mechanism was attributed to Zn²⁺ ion (de)intercalation, facilitated by the large 2 × 2 tunnel structure of α-MnO₂ [4]. However, after a decade of research, alternative mechanisms have been proposed. A comprehensive review categorizes these mechanisms as follows [5]: exclusive Zn²⁺ (de)intercalation, H⁺ conversion, Zn²⁺ and H⁺ dual-ion (de)intercalation, mixed intercalation–conversion, deposition–dissolution coupled with (de)intercalation, and exclusive deposition–dissolution. Hybrid reaction mechanisms involving deposition–dissolution and (de)intercalation are also summarized in ref [6]. These conflicting explanations arise partly from the diverse crystal phases of manganese oxides [7] and partly from misattributing XRD peaks of parasitic layered hydroxide products (e.g., zinc basic sulfate, Zn₄(OH)₆(SO₄)·nH₂O [ZBS] in ZnSO₄ electrolytes) to electrochemically active phases such as zinc buserite, zinc birnessite, ZnMn₂O₄, MnOOH, Mn₂O₃, MnO, or Mn₃O₄ [5]. For instance, non-operando measurements of discharged MnO₂ cathodes may inadvertently lose intercalated water molecules, complicating the analysis [8,9].

To introduce the deposition–dissolution mechanism, which is the main mechanism for this review, key studies are examined here. In ref. [10], STEM imaging of discharged α-MnO₂ nanowires in 1 M ZnSO₄ (Figure 1b) revealed severe tunnel distortion near the surface, indicative of H⁺ intercalation, while the 2 × 2 tunnels showed no occupancy by heavy cations like Zn²⁺. EDS mapping further confirmed the absence of Zn in the discharged nanowires. Similarly, ref. [11] demonstrated no Zn incorporation in α-MnO₂ nanowires cycled in 2 M ZnSO₄ + 0.1 M MnSO₄ via EDS (Figure 1c). This suggests Zn²⁺ intercalation may require MnO₂ phases with larger structural channels, such as 3 × 3 todorokite, layered, or amorphous phases. Notably, both studies observed nanowire fragmentation (Figure 1d) and Zn-rich nanograin formation [10,11], implying Zn²⁺ (de)intercalation could occur in secondary manganese oxides formed during cycling. Additionally, significant Mn²⁺ concentration fluctuations during charge/discharge [11,12,13,14] support the widely accepted deposition–dissolution mechanism of manganese oxides (with or without Mn²⁺). For example, Zn//ZBS batteries utilizing ZnSO₄-based electrolytes with varying concentrations of MnSO₄ have been demonstrated (Figure 1e) [15,16]. These findings suggest that the dominant electrochemical process depends critically on electrolyte conditions. The key factors may be the solubility and remaining concentration of Mn²⁺ in the electrolytes during cycling. For flow batteries, when electrolyte volume exceeds the MnO₂ deposition requirement during charging, excess volume promotes MnO₂ dissolution and stabilizes pH, favoring an exclusive deposition–dissolution mechanism [17]. Furthermore, the formation of inactive materials (e.g., ZnMn₂O₄) in static batteries during cycling, which consume Mn²⁺, may also alter reaction pathways, as shown in Figure 1f [18].

Unlike conventional mildly acidic Zn–MnO₂ batteries, which rely on Mn⁴⁺/Mn³⁺ redox (theoretical capacity: 308 mAh/g), recent zinc electrolytic MnO₂ batteries utilize a two-electron Mn²⁺/Mn⁴⁺ redox process (theoretical capacity: 616 mAh/g) [19]. These systems also exhibit higher discharge voltages [20]. The remainder of this paper focuses on reviewing advances in cathode-free zinc electrolytic MnO₂ batteries, where MnO₂ is electrochemically deposited onto carbon current collectors from Mn²⁺-containing electrolytes [21]. The reactions are summarized in Section 2.2. Without special notes, ions other than Zn²⁺ and Mn²⁺ can be regarded as supporting materials.

Figure 1. The deposition–dissolution mechanism of MnO₂ in zinc-manganese batteries: (a) the first report on a Zn//α-MnO₂ battery with mild aqueous ZnSO₄ as electrolyte, where Zn²⁺ was proposed to insert into the 2 × 2 tunnel structure of α-MnO₂. Reproduced with permission from ref. [4]. Copyright 2012, John Wiley & Son; (b) STEM analysis of discharged α-MnO₂ nanowires. The (001) projections of MnO₂ demonstrate essentially “empty” tunnels without the presence of heavy cations. The high-magnification image (right) depicts non-uniform and anisotropic tunnel distortion caused by H⁺ intercalation. Reproduced with permission from ref. [10]. Copyright 2022, Springer Nature; (c) STEM-HAADF image of short nanorods and STEM-EDS mappings showing the elemental distributions of Mn, O, and Zn in the MnO₂ electrode during the first discharge cycle Reproduced with permission from ref. [11]. Copyright 2016, Springer Nature; (d) low-magnification TEM images of the α-MnO₂ electrode after 1, 10, and 50 cycles (from (left) to (right)), showing fragmentation during cycling. Reproduced with permission from ref. [10]. Copyright 2022, Springer Nature; (e) charge/discharge capacity and Coulombic efficiency of batteries based on cathode Zn₄SO₄(OH)₆·4H₂O (ZHS) with 2 M ZnSO₄ + x M MnSO₄ (variable x) electrolytes in the first cycle, where the ZHS serves as electrolyte salts. Reproduced with permission from ref. [15]. Copyright 2020, Elsevier; (f) Mn-based competitive capacity evolution protocol, where the initial MnO₂ is layered structure and the Mn²⁺ concentration decreased during cycling. Different mechanism stages are shown at the (upper) section along with the cycle profile partitioning and pH values at 1.85 V; the (middle) section illustrates micro-mechanism schematics across four regions; the (bottom) section presents the Mn²⁺ concentration and contribution ratio. Reproduced with permission from ref. [18]. Copyright 2022, Royal Society of Chemistry.

Figure 1. The deposition–dissolution mechanism of MnO₂ in zinc-manganese batteries: (a) the first report on a Zn//α-MnO₂ battery with mild aqueous ZnSO₄ as electrolyte, where Zn²⁺ was proposed to insert into the 2 × 2 tunnel structure of α-MnO₂. Reproduced with permission from ref. [4]. Copyright 2012, John Wiley & Son; (b) STEM analysis of discharged α-MnO₂ nanowires. The (001) projections of MnO₂ demonstrate essentially “empty” tunnels without the presence of heavy cations. The high-magnification image (right) depicts non-uniform and anisotropic tunnel distortion caused by H⁺ intercalation. Reproduced with permission from ref. [10]. Copyright 2022, Springer Nature; (c) STEM-HAADF image of short nanorods and STEM-EDS mappings showing the elemental distributions of Mn, O, and Zn in the MnO₂ electrode during the first discharge cycle Reproduced with permission from ref. [11]. Copyright 2016, Springer Nature; (d) low-magnification TEM images of the α-MnO₂ electrode after 1, 10, and 50 cycles (from (left) to (right)), showing fragmentation during cycling. Reproduced with permission from ref. [10]. Copyright 2022, Springer Nature; (e) charge/discharge capacity and Coulombic efficiency of batteries based on cathode Zn₄SO₄(OH)₆·4H₂O (ZHS) with 2 M ZnSO₄ + x M MnSO₄ (variable x) electrolytes in the first cycle, where the ZHS serves as electrolyte salts. Reproduced with permission from ref. [15]. Copyright 2020, Elsevier; (f) Mn-based competitive capacity evolution protocol, where the initial MnO₂ is layered structure and the Mn²⁺ concentration decreased during cycling. Different mechanism stages are shown at the (upper) section along with the cycle profile partitioning and pH values at 1.85 V; the (middle) section illustrates micro-mechanism schematics across four regions; the (bottom) section presents the Mn²⁺ concentration and contribution ratio. Reproduced with permission from ref. [18]. Copyright 2022, Royal Society of Chemistry.

2. Electrolytes Modifications

2.1. Acid Electrolytes

Since Mn²⁺/MnO₂ redox reactions are more favorable in acidic electrolytes, zinc metal anodes were initially deemed unsuitable for electrolytic MnO₂ batteries due to hydrogen evolution issues [22,23]. Early studies instead explored alternative anodes such as hydrogen (H₂) [24], Cu [22,23,25], Bi [22], Cd [23] and Pb [26]. For example, the MnO₂–H₂ battery utilized a 40% platinum-on-carbon catalyst to facilitate H₂ reactions and employed a cylindrical structure for H₂ storage (Figure 2a) [24]. While this system maintained a stable potential of ~1.254 V after 80 h of self-discharge, its remaining capacity was only ~71%. Metals stable in acidic electrolytes (e.g., Cu or Bi) were paired with MnO₂, achieving open-circuit voltages below 1.2 V (Figure 2b) [22]. However, prolonged cycling revealed Cu²⁺ insertion into MnO₂ at discharge voltages < 0.5 V. Then, hybrid systems integrating MnO₂–Cu and MnO₂–H₂ via a two-step process were proposed, though at the cost of Cu consumption, H₂SO₄ usage, and excess energy input [25]. Alternatively, toxic Pb- and Cd-based MnO₂ batteries achieved higher discharge voltages of 1.55 V and 1.68 V, respectively [23,26]. Notably, the MnO₂/Mn²⁺ redox exhibited superior reversibility compared to commercial PbO₂ cathodes, with higher current density and lower oxygen evolution overpotential (Figure 2c) [26].

The first Zn//electrolytic MnO₂ battery demonstrated a high discharge voltage of 1.95 V, exceeding those of Zn²⁺ or H⁺ insertion mechanisms [20]. This system employed Zn foil, carbon cloth, and an electrolyte of 1 M ZnSO₄ + 1 M MnSO₄ + 0.1 M H₂SO₄. Initial charging deposited Zn and ε-MnO₂, while subsequent cycles involved H⁺ and Zn²⁺ interactions (Figure 2d) [20]. Remarkably, the low-pH (~1) electrolyte maintained a stable pH and a 2.41 V working window for at least 50 cycles. Later studies showed that adding 0.07 mM polyvinylpyrrolidone (PVP) to this electrolyte significantly improved Coulombic efficiency and cyclability by modulating cationic solvation structures, enhancing Mn²⁺/Zn²⁺ migration and MnO₂/Zn deposition–dissolution kinetics (Figure 3a) [27].

Beyond H₂SO₄, other acids such as HCl [28], acetic acid (HAc) [29,30,31] and methanesulfonic acid (MSA) [32] have been explored. Replacing 0.1 M H₂SO₄ with 0.2 M HAc in 1 M ZnSO₄ + 1 M MnSO₄ increased capacity contributions from (de)intercalation mechanisms (Figure 3b). HAc also reduced anode corrosion and hydrogen evolution while stabilizing pH fluctuations via reversible ionization (Figure 3c), leading to higher reversible capacity and improved cycling stability [29]. The pH buffering effect of HAc is due its weak acid nature compared with H₂SO₄ [19], which may be the case for MSA. Comparative studies of Zn//α-MnO₂ batteries in 2 M ZnSO₄ with 0.1 M MnSO₄ or carboxylate additives highlighted the buffering effect of anions on interfacial H⁺ and steric hindrance of Mn²⁺ migration (Figure 3d) [33]. This mechanism may extend to cathode-free systems. Moreover, further increasing acetate content (e.g., 1 M ZnSO₄ + 0.1 M MnSO₄ + 0.1 M NaAc + 0.5 M HAc) shifted the dominant mechanism toward MnO₂ deposition–dissolution by impeding H⁺/Zn²⁺ intercalation [30]. Further modification includes replacing sulfate with acetate and addition of KBr, which tripled ionic conductivity, likely due to Br⁻ integration into Zn²⁺/Mn²⁺ solvation shells, enhancing cation mobility [31]. In addition, charging protocols also impact performance; while high-voltage constant charging enables fast charging, it promotes Mn³⁺ formation, lowering Coulombic efficiency and cycle life (Figure 3e) [32].

Figure 2. Reaction mechanisms of Mn²⁺/MnO₂ in acidic electrolytes: (a) schematic and digital photograph of the scaled-up membrane-free cylindrical Mn–H₂ cell. Reproduced with permission from ref. [24]. Copyright 2018, Springer Nature; (b) pourbaix diagram of MnO₂ cathode and Zn/Bi/Cu anode materials under varying pH conditions, to show the voltage difference. Reproduced with permission from ref. [22]. Copyright 2019, John Wiley & Son; (c) cyclic voltammetry (CV) comparison between Mn²⁺/MnO₂ and PbSO₄/PbO₂ systems, where the currents indicate the reactivity. Reproduced with permission from ref. [26]. Copyright 2020, Royal Society of Chemistry; (d) schematic illustration of reactions steps for the zinc electrolytic MnO₂ battery. Reproduced with permission from ref. [20]. Copyright 2019, John Wiley & Son: (top): Initial chronoamperometric charge process. Zn²⁺ and Mn²⁺ cations in the electrolyte are deposited onto the anode and cathode current collectors, forming metallic Zn and MnO₂, respectively. (bottom): Subsequent reversible galvanostatic discharge and chronoamperometric charge cycles. During discharge, electrolytic MnO₂ undergoes H⁺ and Zn²⁺ intercalation at the cathode, while Zn stripping and dissolution occur at the anode.

Figure 2. Reaction mechanisms of Mn²⁺/MnO₂ in acidic electrolytes: (a) schematic and digital photograph of the scaled-up membrane-free cylindrical Mn–H₂ cell. Reproduced with permission from ref. [24]. Copyright 2018, Springer Nature; (b) pourbaix diagram of MnO₂ cathode and Zn/Bi/Cu anode materials under varying pH conditions, to show the voltage difference. Reproduced with permission from ref. [22]. Copyright 2019, John Wiley & Son; (c) cyclic voltammetry (CV) comparison between Mn²⁺/MnO₂ and PbSO₄/PbO₂ systems, where the currents indicate the reactivity. Reproduced with permission from ref. [26]. Copyright 2020, Royal Society of Chemistry; (d) schematic illustration of reactions steps for the zinc electrolytic MnO₂ battery. Reproduced with permission from ref. [20]. Copyright 2019, John Wiley & Son: (top): Initial chronoamperometric charge process. Zn²⁺ and Mn²⁺ cations in the electrolyte are deposited onto the anode and cathode current collectors, forming metallic Zn and MnO₂, respectively. (bottom): Subsequent reversible galvanostatic discharge and chronoamperometric charge cycles. During discharge, electrolytic MnO₂ undergoes H⁺ and Zn²⁺ intercalation at the cathode, while Zn stripping and dissolution occur at the anode.

Figure 3. Modifications for depositing MnO₂ in acidic electrolytes: (a) schematic of the role of the cationic accelerator (PVP) in facilitating cation transport within aqueous MnO₂–Zn batteries. Reproduced with permission from ref. [27]. Copyright 2022, John Wiley & Son; (b) galvanostatic discharge profiles and capacity contributions from MnO₂ dissolution in electrolytes containing 0.1 M H₂SO₄ and 0.2 M HAc, where the high-voltage regions correspond deposition–dissolution process; (c) in situ pH measurements and corresponding H⁺ concentrations (circles) during cycling in a battery with 0.1 M HAc and 0.2 M H₂SO₄, where less pH fluctuation was showed for HAc case Reproduced with permission from ref. [29]. Copyright 2021, John Wiley & Son; (d) evolution of the electric double layer model at the cathode–electrolyte interface during charging: Modulation of interfacial H⁺ concentration and operating potential to thermodynamically enhance capacity. Reproduced with permission from ref. [33]. Copyright 2022, John Wiley & Son. (e) The effects of different charging conditions on the deposited MnO₂. Reproduced with permission from ref. [32]. Copyright 2023, Elsevier.

Figure 3. Modifications for depositing MnO₂ in acidic electrolytes: (a) schematic of the role of the cationic accelerator (PVP) in facilitating cation transport within aqueous MnO₂–Zn batteries. Reproduced with permission from ref. [27]. Copyright 2022, John Wiley & Son; (b) galvanostatic discharge profiles and capacity contributions from MnO₂ dissolution in electrolytes containing 0.1 M H₂SO₄ and 0.2 M HAc, where the high-voltage regions correspond deposition–dissolution process; (c) in situ pH measurements and corresponding H⁺ concentrations (circles) during cycling in a battery with 0.1 M HAc and 0.2 M H₂SO₄, where less pH fluctuation was showed for HAc case Reproduced with permission from ref. [29]. Copyright 2021, John Wiley & Son; (d) evolution of the electric double layer model at the cathode–electrolyte interface during charging: Modulation of interfacial H⁺ concentration and operating potential to thermodynamically enhance capacity. Reproduced with permission from ref. [33]. Copyright 2022, John Wiley & Son. (e) The effects of different charging conditions on the deposited MnO₂. Reproduced with permission from ref. [32]. Copyright 2023, Elsevier.

2.2. Dual Electrolytes

To mitigate Zn corrosion in acidic electrolytes, one strategy involves employing selective membranes and introducing a Zn-compatible secondary electrolyte [34,35]. Unlike the reactions in Figure 2d, the overall cell reaction shifts from Equations (R1) and (R2) to Equations (R1) and (R3), increasing the theoretical potential from 1.992 V to 2.428 V (Figure 4a).

Reactions:

Cathode: Mn²⁺ + 2H₂O ⇄ MnO₂ + 4H⁺ + 2e⁻; E₀ = 1.229 V vs. SHE

(R1)

Anode (Acidic): Zn²⁺ + 2e⁻ ⇄ Zn; E₀ = −0.763 V vs. SHE

(R2)

Anode (Alkaline): Zn(OH)₄²⁻ + 2e⁻ ⇄ Zn + 4OH⁻; E₀ = −1.199 V vs. SHE

(R3)

Bipolar membranes (BPMs) were used to separate dual electrolytes, showing lower charge–discharge voltage gaps compared to anion/cation exchange membranes [36]. While BPMs allow H⁺/OH⁻ crossover to form water (Figure 4b) [36], another report adopts K⁺ and SO₄²⁻ permeation [37]. Both systems achieved average voltages of ~2.44 V. Notably, introducing Ni²⁺ into the catholyte enhanced Mn²⁺/MnO₂ electrolysis kinetics, improving the high-current (50 C) performance [36].

To prevent electrolyte neutralization, a neutral middle chamber separated by anion/cation exchange membranes was implemented (Figure 4c) [38]. Concentration gradients across the membranes generated an electric field, boosting the open-circuit voltage further to 2.83 V [38]. In this report, a scaled-up static pack powering LED panels using hybrid wind–photovoltaic energy was demonstrated. Similarly, a flow battery variant employed Bi₂O₃-modified carbon felt cathodes and Ni²⁺/Mg²⁺ catalysts in the catholyte was reported [39]. During charging process, Bi₂O₃ was dissolved and then reconstructed as nanosized crystal onto MnO₂, while NiO₂ formed from Ni²⁺. However, despite stable catholyte/anolyte pH, the middle chamber’s pH dropped from 4.8 to 2.35 after five cycles, indicating proton leakage [39].

For systems accepting discharge voltages of ~1.9 V, a membrane with solely anion or cation exchange functions were investigated. For instance, near-neutral anolytes (instead of alkaline) paired with anion exchange membranes (AEMs) were tested (Figure 4d) [40]. To protect the Zn anode from the leaked proton, ZnAc₂ was applied at anolyte to trap the leaked protons [40]. To further shield the proton leakage, hydrophobic bis(trifluoromethylsulfonyl)imide (TFSI⁻)-conducting membranes (PILG: polymer/ionic liquid/graphene) which selectively transported TFSI⁻ was applied [41]. However, both anolyte and catholyte required costly TFSI⁻ incorporation. As another case, a cation exchange membrane was applied to separate alkaline anolyte and mild acid catholyte; however, the MnO₂ was pre-coated on the cathode first, which may be the reason for its inefficiency of MnO₂ deposition in the mild acid environment [42].

Besides ion-exchange membranes, hydrogels and salt bridges were investigated to separate the dual electrolytes. As shown in Figure 4e, a Zn–MnO₂ battery based on a liquid acidic catholyte coupled with KOH hydrogel and cellophane as a separator displayed an open-circuit voltage of 2.4–2.8 V [43]. Another cathode-free battery design, shown in Figure 4f, utilized a polyacrylamide (PAM)-based catholyte containing 1 M MnSO₄ and 1 M H₂SO₄, along with an anolyte of 2 M ZnSO₄ [44]. Furthermore, a tri-layer hydrogel electrolyte was reported, comprising an acidic region (PAM + ZnSO₄ + H⁺), a buffer region (sodium polyacrylate (PAANa) + ZnSO₄), and a conservation region (PAM + ZnSO₄), as illustrated in Figure 4g [45]. Although this approach avoided the use of expensive ion-selective membranes, concerns regarding proton leakage persisted.

When a traditional salt bridge gel composed of Na₂SO₄ and agar was employed to separate the alkaline anolyte and acidic catholyte, a discharge voltage of 2.37 V was achieved, as depicted in Figure 4h [46]. For alternative salt bridge strategies, poly(vinyl alcohol) (PVA) gel was first immersed in ZnAc₂ solution and then in concentrated ZnSO₄ solution, leveraging strong Hofmeister effects to form a stable salt bridge interface (Figure 4i) [47]. Another study exploited the low solubility of K₂SO₄ in Pluronic F-127 (polyethylene–polypropylene glycol). When H₂SO₄ catholyte and KOH anolyte came into contact, a self-forming K₂SO₄ salt bridge was generated at the interface [48].

Figure 4. Battery configurations with dual electrolytes: (a) electrochemical stability window of the hybrid aqueous battery system and minimum electrolysis voltage required for stable operation; (b) schematic of a Zn//MnO₂ hybrid aqueous battery employing dual electrolytes during charge/discharge cycles, where water is formed and decomposed within membranes. Reproduced with permission from ref. [36]. Copyright 2020, John Wiley & Son; (c) cell structure and reaction mechanisms in an aqueous decoupled Zn//MnO₂ battery: Acidic and alkaline electrolytes (separated by a neutral central chamber) enable decoupled cathode and anode reactions during discharge/charge. Reproduced with permission from ref. [38]. Copyright 2020, Springer Nature; (d) pH evolution of the anolyte in an electrolyte-decoupled MnO₂/Zn battery (i.e., acid + neutral electrolytes). Reproduced with permission from ref. [40]. Copyright 2022, American Chemical Society; (e) battery design utilizing alkaline KOH gel electrolyte paired with acidic liquid electrolyte. Reproduced with permission from ref. [43]. Copyright 2019, American Chemical Society; (f) schematic of a cathode-free battery integrating mixed and decoupled gel electrolytes. Reproduced with permission from ref. [44]. Copyright 2021, American Chemical Society; (g) tri-layer hydrogel electrolyte structure in a Zn//MnO₂ battery, comprising “Acid zone”, “Buffer zone”, and “Conservation zone”. Reproduced with permission from ref. [45]. Copyright 2021, John Wiley & Son; (h) digital image of a electrolytes-decoupled glass electrolytic cell with a salt bridge. Reproduced with permission from ref. [46]. Copyright 2023, Elsevier; (i) fabrication of PVA-based salt bridge film via the Hofmeister effect. PZA sol denotes a light-white sol after addition of ZnAc₂ (zinc acetate) to the PVA solution. PZAS gel denotes the PZA sol after addition of concentrated ZnSO₄ which induces gelation. Reproduced with permission from ref. [47]. Copyright 2022, John Wiley & Son.

Figure 4. Battery configurations with dual electrolytes: (a) electrochemical stability window of the hybrid aqueous battery system and minimum electrolysis voltage required for stable operation; (b) schematic of a Zn//MnO₂ hybrid aqueous battery employing dual electrolytes during charge/discharge cycles, where water is formed and decomposed within membranes. Reproduced with permission from ref. [36]. Copyright 2020, John Wiley & Son; (c) cell structure and reaction mechanisms in an aqueous decoupled Zn//MnO₂ battery: Acidic and alkaline electrolytes (separated by a neutral central chamber) enable decoupled cathode and anode reactions during discharge/charge. Reproduced with permission from ref. [38]. Copyright 2020, Springer Nature; (d) pH evolution of the anolyte in an electrolyte-decoupled MnO₂/Zn battery (i.e., acid + neutral electrolytes). Reproduced with permission from ref. [40]. Copyright 2022, American Chemical Society; (e) battery design utilizing alkaline KOH gel electrolyte paired with acidic liquid electrolyte. Reproduced with permission from ref. [43]. Copyright 2019, American Chemical Society; (f) schematic of a cathode-free battery integrating mixed and decoupled gel electrolytes. Reproduced with permission from ref. [44]. Copyright 2021, American Chemical Society; (g) tri-layer hydrogel electrolyte structure in a Zn//MnO₂ battery, comprising “Acid zone”, “Buffer zone”, and “Conservation zone”. Reproduced with permission from ref. [45]. Copyright 2021, John Wiley & Son; (h) digital image of a electrolytes-decoupled glass electrolytic cell with a salt bridge. Reproduced with permission from ref. [46]. Copyright 2023, Elsevier; (i) fabrication of PVA-based salt bridge film via the Hofmeister effect. PZA sol denotes a light-white sol after addition of ZnAc₂ (zinc acetate) to the PVA solution. PZAS gel denotes the PZA sol after addition of concentrated ZnSO₄ which induces gelation. Reproduced with permission from ref. [47]. Copyright 2022, John Wiley & Son.

2.3. Near Neutral Electrolytes

To eliminate the need for ion-selective membranes, near-neutral electrolytes have been investigated for depositing MnO₂ while mitigating severe Zn anode corrosion. As illustrated in Figure 5a, a membrane-free flow battery was constructed using zinc foil and carbon felt electrodes with a 1 M MnSO₄ + 1 M ZnSO₄ electrolyte [17]. Additionally, another novel cell design further enhanced electrolyte utilization by rotating deposited MnO₂ out of the electrolyte. A subsequent study revealed the detailed experimental parameters, including carbon current collector selection, electrolyte salt concentration/ratio, electrolyte volume, and battery configuration [49]. The results demonstrated that lower Zn²⁺ proportions improved discharge curve stability, while higher Mn²⁺ concentrations reduced pH (2~4) and elevated discharge voltage. However, excessively high Mn²⁺ concentrations (3 M) hindered MnO₂ dissolution and induced salt precipitation, impairing long-term cycling performance [49].

To enhance the cycling stability of a cathode-free battery (control electrolyte: 0.5 M MnSO₄ + 1 M ZnSO₄), 7 M methylurea (Mu) was introduced. Mu, containing hydrophilic and hydrophobic moieties, transformed the electrolyte into a nanomicellar system (Figure 5b). These nanomicelles disrupted the continuous aqueous phase and its hydrogen-bonding network, restructuring local hydrogen bonds within micelles or at micelle–water interfaces. Additionally, Mu participated in Zn²⁺/Mn²⁺ solvation, displacing water molecules to lower desolvation energy barriers and regulate nanograin formation [50]. The restructured hydrogen-bonding network suppressed water decomposition and parasitic side reactions on the Zn anode. A protective solid–electrolyte interface (SEI) layer (Zn_x(Mu)ᵧSO₄·nH₂O) also formed, preventing anode corrosion. While the enlarged Zn²⁺ solvation radius ([Zn(Mu)_x(H₂O)ᵧ]²⁺) may hinder Zn²⁺ intercalation, the elevated pH (5.26) still favored MnO₂ deposition [50]. In another report, glycine (Gly) was explored as an additive for 0.1 M MnSO₄ + 1 M ZnSO₄ electrolyte [56]. Gly complexes with Mn²⁺, releasing H⁺ to promote MnO₂ dissolution at the cathode, while adsorbing on the Zn anode surface to shield it from detrimental reactions [56].

The cathode-free battery can also serve as a control sample for normal batteries where MnO₂ is pre-coated on the cathode. For example, the addition of acetate additives to sulfate electrolytes significantly increased the battery capacity, as shown in Figure 5c. This enhanced capacity was validated using stainless steel as the cathode-free electrode, confirming the additional involvement of a dissolution–deposition mechanism [51]. Another example involves the incorporation of bentonite colloid (Ben-colloid) into the electrolyte (Figure 5d). The Ben colloid was hypothesized to self-concentrate Mn²⁺ ions and reduce the deposition potential, where a stainless steel was also applied as the cathode-free electrode [52]. Both studies demonstrated the pH buffering effect of these additives with weak acid nature, as illustrated in Figure 5e. Specifically, when excess 0.2 g/mL ZnO was introduced, a gel-like electrolyte formed, maintaining a stable pH of 6.4 throughout the charge–discharge cycles (Figure 5f) [53]. In this gel electrolyte, H⁺ ions were effectively eliminated, thereby promoting manganese deposition. Further analysis revealed that Mn²⁺ ions and ZBS synergistically generated Zn₂Mn₃O₈·H₂O nanorod arrays as active material during charging [53]. In a separate study, ZnO was incorporated onto electrodes, but its primary effect appeared to be limited to modifying the electrolyte’s pH [57]. Additionally, Mn(H₂PO₄)₂ was employed as a pH buffering agent in chloride-based electrolytes [58]. The buffering mechanism for Mn²⁺/MnO₂ conversion was further extended to weak Brønsted acids [59]. Without pH buffering additives, the chloride electrolyte was compared with other anion salts, where 3 M ZnCl₂ + 0.1 M MnCl₂ was recommended [60].

Acetate-based (Ac) electrolytes have been extensively studied. Cyclic voltammetry (CV) of acetate-, sulfate-, and sulfonate-based electrolytes (1 M Zn²⁺ + 0.4 M Mn²⁺, Figure 5g) revealed distinct reaction mechanisms: acetate electrolytes exhibited a two-electron MnO₂/Mn²⁺ reaction (reduction peak at 1.5 V), whereas sulfate/sulfonate systems showed Zn²⁺/H⁺ insertion into the MnO₂ cathode, characterized by dual reduction peaks at 1.35 V and 1.2 V [54]. Complete MnO₂ dissolution after discharge was confirmed in acetate electrolytes. Density functional theory (DFT) calculations indicated that acetate anions adsorb onto MnO₂ surface sites, weakening H₂O binding and reducing dissolution energy barriers [54]. Thus, acetate anions enhanced Zn anode compatibility. A flow battery utilizing 0.5 M MnAc₂ + 0.5 M ZnCl₂ + 2 M KCl as electrolyte achieved an area capacity of 10 mAh/cm² [61]. After applying carbon black coating and higher electrolyte concentrations, the capacity further increased to 20 mAh/cm². Subsequent optimization of electrolyte to 1 M MnAc₂ + 1 M ZnAc₂ + 2 M KCl fully coordinated generated H⁺, achieving an energy density of 42 Wh/L in static batteries [61]. Although Mn³⁺ disproportionation was not detected with Na₄P₂O₇ as an indicator, ethylenediaminetetraacetate (EDTA) confirmed Mn³⁺ presence due to its stronger complexation capability than Ac⁻ [55,61]. Citrate and EDTA were further compared with acetate as complexing agents (Figure 5h). However, citrate electrolytes exhibited lower pH, while EDTA-based batteries suffered from self-discharge [55]. The criteria for selecting optimal complexing agents were proposed as follows [55]: (i) strong Mn²⁺ binding affinity to limit free Mn²⁺ concentration; (ii) minimal H⁺ consumption to maintain higher reaction potentials; (iii) weak interactions with Zn²⁺ and Mn³⁺ to avoid competitive binding (Zn²⁺ vs. Mn²⁺) or MnO₂ suspension (Mn³⁺-induced instability), thereby preserving Coulombic efficiency (CE).

2.4. Mediator Additives

The disproportionation reaction of oxidized Mn³⁺ (2Mn³⁺ + 2H₂O → MnO₂ + 4H⁺ + Mn²⁺) leads to the formation of “dead” MnO₂, which detaches from current collectors [32,55,61,62]. To recycle the dead MnO₂, mediator additives are introduced into the electrolyte to reduce it back to Mn²⁺. As shown in Figure 6a,b, the addition of iodide (I⁻) enables the oxidation of iodide to I₃⁻, facilitating the conversion of solid MnO₂ into soluble Mn²⁺ [63]. Figure 6c demonstrates that dead MnO₂ dispersing in the electrolyte after 18 cycles and can be dissolved completely upon adding KI, resulting in a transparent solution. Moreover, the generated I₃⁻ is also found to contribute to capacity as zinc–iodine batteries [63,64,65]. Similarly, bromide (Br⁻) is proved to enhance capacity (Figure 6d), but its oxidation requires a high charging voltage (e.g., 2.2 V as the case in acidic sulfate electrolytes) [66]. When the charging voltage is fixed below 1.8 V, no Br₂ forms, and Br⁻ primarily functions to improve electrolyte ion conductivity [31,65]. When chromium cations (Cr³⁺, with Cr²⁺/Cr³⁺ redox potential at −0.407 V vs. SHE, higher than Zn/Zn²⁺ at −0.763 V vs. SHE) act as mediators, uniform MnO₂ deposition is promoted (Figure 6e) [55]. Additionally, ultrasonic treatment was proposed to recover batteries by dislodging over-thick MnO₂ layers (Figure 6f) [55].

Although mediator additives improve capacity and MnO₂ recovery, they also induce shuttle reactions with both electrodes [21,67,68]. These reactions become severe in thin-film batteries with limited electrolyte volume and narrow electrode gaps (Figure 6g). To mitigate this, starch was added to coordinate with I₃⁻, significantly extending battery shelf life (Figure 6h) [21,68]. Another strategy to mitigate shuttling employs activated carbon electrodes [67]. Compared to graphite, activated carbon’s high surface area provides more sites for MnO₂ anchoring. Furthermore, positive charges accumulating at the cathode (supercapacitor-like effect) electrostatically attract I₃⁻. Combined with reduced MnO₂ detachment and electrostatic interactions, this approach enhances both Coulombic efficiency and shelf life (Figure 6i,j) [67].

3. Electrode Modifications

3.1. Anode Modifications

For Zn anodes, constructing a solid electrolyte interface (SEI) is a common strategy to combat hydrogen evolution corrosion (HEC) caused by acidic electrolytes. A proton-resistant Pb-containing interface (composed of Pb and Pb(OH)₂) was formed on Zn anodes (denoted as Zn@Pb) via a facile displacement reaction (i.e., immersing Zn foil in PbAc₂ solution). This interface further transformed in situ into PbSO₄ upon H₂SO₄ corrosion [69]. When PbAc₂ was further added to the working electrolyte, Zn anode stability was significantly enhanced (Figure 7a,b). In situ XRD analysis revealed that Pb is plated and stripped on the anode surface during cycling. The Pb layer protects the Zn plating layer from HEC throughout the plating/stripping process due to Pb’s low affinity for H⁺ and strong Pb–Zn/Pb–Pb bonding [69].

In another report, a phytic acid (PA)-modified layer was fabricated by soaking Zn foil in a 1% PA solution [70]. Instead of H₂SO₄, H₃PO₄ was used as an acidic additive in the electrolyte. As shown in Figure 7c,d, incorporating H₃PO₄ suppressed hydrogen evolution, with an optimal electrolyte pH of 2.2. However, even with the PA-modified layer, HEC occurred when H₂SO₄ was introduced [70]. This occurred because zinc phosphate formation at the interface is necessary. During the Zn stripping, Zn²⁺ ions accumulate at the interface to form zinc phosphate (Zn₃(PO₄)₂·4H₂O, ZPO) and zinc hydroxide sulfate (Zn₄SO₄(OH)₆·5H₂O, ZHS) at the PA-modified layer. Subsequently, during Zn plating, H⁺ ions migrate alongside Zn²⁺ to the anode, where they are adsorbed by the PA-modified layer and absorbed by ZPO/ZHS. Throughout repeated plating/stripping cycles, H⁺ ions are reversibly adsorbed within the interphase, maintaining a stable proton concentration in the electrolyte [70].

3.2. Cathode Modifications

For most studied pre-coated MnO₂ batteries, constructing a cathode electrolyte interface (CEI) is a common strategy to suppress MnO₂ dissolution. Here, the CEI layer prevents dead MnO₂ dispersion into the electrolyte via limiting Mn²⁺ diffusion. However, though we need the CEI to capture Mn³⁺ and H⁺, the CEI layer must also allow MnO₂/Mn²⁺ redox reactions, as illustrated in Figure 7e [71]. After evaluating the electrostatic potential distributions of monomers, including acrylamide (AAm), 2-acrylamido-2-methylpropanesulfonic acid (AMPS), 2-chloroacrylic acid (ClAA), acrylic acid (AA), methacrylic acid (MAA), and vinylphosphonic acid (VPA), MAA was selected for in situ gel polymerization on carbon cloth (PMAA@CC) in the presence of Zn²⁺ and Mn²⁺ [71]. By minimizing dead MnO₂ in the electrolyte, H⁺ generation is reduced, which in turn mitigates Zn anode corrosion and enhancing the reversibility of MnO₂/Mn²⁺ reactions. Another approach to suppress Mn³⁺ diffusion and disproportionation involves manganese polyacrylate (PAAMn), which was coated on stainless steel foil as a cathode active material [72]. Due to the relatively low ion conductivity, the diffusion of Mn³⁺ may also be limited in hydrogels, as reported in refs. [44,45,73,74].

Doping deposited MnO₂ is an effective method to improve conductivity. In dual-electrolyte systems, Ni³⁺ and Bi³⁺ doping catalyzes Mn²⁺/MnO₂ reactions and enhances MnO₂ conductivity without concerns about cation crossover due to anolyte/catholyte separation [36,39,75]. For single-electrolyte batteries, bismuth (III) pyridine-3,5-dicarboxylate (BiMOF) was employed as a Bi³⁺ reservoir to stabilize MnO₂ structure by controlled Bi³⁺ release during Mn²⁺ deposition [76]. Similarly, adding Al³⁺ to the electrolyte was reported to form a proton–donor reservoir to maintain acidic conditions and to introduce oxygen vacancies into MnO₂ [77].

As different phases were reported for the deposited MnO₂, the crystal structure may vary with factors, such as anions, pH, and charging protocols. Notably, incorporating 0.1 mM surfactant molecules into the electrolyte induces favorable c-axis orientation in hexagonal Zn and MnO₂ deposition (Figure 7f) [78]. The surfactant (e.g., t-Oct-C₆H₄-(OCH₂CH₂)_nOH, n ≈ 9–10) with hydrophobic tail and hydrophilic head first form an aligned bilayer structure, followed by a gradient liquid crystal interphase during deposition. This interphase comprises lamellar liquid crystals (near the electrode), hexagonal liquid crystals (intermediate), and micelle clusters (near the electrolyte). Consequently, the layered δ-phase MnO₂ was deposited instead of the ε-phase. (Figure 7g). Additionally, the surfactant promotes (002) plane growth of Zn, yielding a flat, compact anode morphology with enhanced corrosion resistance against hydrogen evolution [78].

Figure 7. Anti-corrosion strategies for Zn anodes and modifications to mitigate dissolution and crystal structures of MnO₂ cathode: (a) cycling performance of Zn@Pb-Ad (Ad denotes PbAc₂ additive in electrolyte), Zn@Pb, and bare Zn symmetric batteries in 0.2 M H₂SO₄ (4 mA cm⁻², 0.5 mAh cm⁻²); (b) post-cycling photographs of Zn symmetric batteries, where the control bare Zn anode fell apart. Reproduced with permission from ref. [69]. Copyright 2023, John Wiley & Son; (c) optical images of Zn deposition on PAZn anodes with H₃PO₄ (PS) or H₂SO₄ (SS) additives at 30 mA cm⁻² for 6 min. Hydrogen bubble formed for the H₂SO₄ case; (d) linear sweep voltammetry (LSV) curves of PAZn anodes in 1 M Na₂SO₄ at varying pH adjusted by H₃PO₄/H₂SO₄. The S shape curve for the H₃PO₄ case indicates the existence of optimal pH value. Reproduced with permission from ref. [70]. Copyright 2025, John Wiley & Son; (e) in situ hydrogel ion-anchoring strategy to form CEI for suppressing Mn³⁺ and H⁺ diffusion. Reproduced with permission from ref. [71]. Copyright 2024, American Chemical Society; (f) dynamic phase transition during deposition: from aligned surfactant bilayers to gradient liquid crystal interphases (lamellar near the electrode, hexagonal intermediate, and micelle clusters near the electrolyte); (g) TEM images of MnO₂ deposited in pristine electrolyte (ε-phase, (top)) versus surfactant-modified electrolyte (δ-phase, (bottom)). Reproduced with permission from ref. [78]. Copyright 2024, Springer Nature.

Figure 7. Anti-corrosion strategies for Zn anodes and modifications to mitigate dissolution and crystal structures of MnO₂ cathode: (a) cycling performance of Zn@Pb-Ad (Ad denotes PbAc₂ additive in electrolyte), Zn@Pb, and bare Zn symmetric batteries in 0.2 M H₂SO₄ (4 mA cm⁻², 0.5 mAh cm⁻²); (b) post-cycling photographs of Zn symmetric batteries, where the control bare Zn anode fell apart. Reproduced with permission from ref. [69]. Copyright 2023, John Wiley & Son; (c) optical images of Zn deposition on PAZn anodes with H₃PO₄ (PS) or H₂SO₄ (SS) additives at 30 mA cm⁻² for 6 min. Hydrogen bubble formed for the H₂SO₄ case; (d) linear sweep voltammetry (LSV) curves of PAZn anodes in 1 M Na₂SO₄ at varying pH adjusted by H₃PO₄/H₂SO₄. The S shape curve for the H₃PO₄ case indicates the existence of optimal pH value. Reproduced with permission from ref. [70]. Copyright 2025, John Wiley & Son; (e) in situ hydrogel ion-anchoring strategy to form CEI for suppressing Mn³⁺ and H⁺ diffusion. Reproduced with permission from ref. [71]. Copyright 2024, American Chemical Society; (f) dynamic phase transition during deposition: from aligned surfactant bilayers to gradient liquid crystal interphases (lamellar near the electrode, hexagonal intermediate, and micelle clusters near the electrolyte); (g) TEM images of MnO₂ deposited in pristine electrolyte (ε-phase, (top)) versus surfactant-modified electrolyte (δ-phase, (bottom)). Reproduced with permission from ref. [78]. Copyright 2024, Springer Nature.

3.3. Current Collectors

For the cathode-free structure, the initial charging process deposits MnO₂ from the electrolyte onto current collectors. During this step, metallic zinc is also plated, suggesting that a conventional zinc foil anode may be unnecessary. A dual electrode-free structure—using current collectors as initial electrodes—was demonstrated in MnO₂–Cd/Zn/Cu batteries (Figure 8a) [21,23,78,79]. Furthermore, the cycling performance of cuvette cells and thin-film batteries with/without zinc anodes was compared using an electrolyte containing 1 M ZnAc₂ + 1 M MnAc₂ + 2 M KCl + 0.1 M KI. Results showed that electrode-free thin-film batteries (carbon//carbon configuration) exhibited lower Coulombic efficiencies (CEs) but longer cycle life than Zn//carbon systems (Figure 8b) [21]. Thus, although excess zinc compensates for deposition inefficiency and corrosion losses, it intensifies hydrogen evolution side reactions. Hydrogen generation caused battery inflation and failure, particularly sensitive in sealed thin-film structures with lean electrolyte volume and confined space, which are closer to real-world applications [21].

Interestingly, the electrode-free design enables a facile assembly of batteries with complex geometries. For example, a cable-shaped battery was fabricated by embedding two commercial carbon fiber ropes in a starch gel electrolyte (Figure 8c) [21,68]. During initial charging, Zn and MnO₂ were directly deposited on the fibers, eliminating troublesome pre-coating steps. Here, the starch additive not only stabilized the gel electrolyte but also suppressed iodide-mediated shuttle reactions [21,68]. Another example is a micro-battery (Figure 8c), where polyacrylamide (PAM) hydrogel was applied to microelectrodes [79]. After full discharge, the battery reverts to its initial state (only current collectors), enabling polarity switching (Figure 8d,e). This feature enhances safety during over-discharge and expands application flexibility [21,79].

Integrating supercapacitor electrodes with the electrolyte of a zinc electrolytic MnO₂ battery yielded a hybrid “supercapattery” (Figure 8f) [67]. For conventional supercapacitors, the stored charge quantity Q is only proportional to charged voltage V, if the capacitance C is fixed, i.e., Q = C × V. However, the electrochemical working window of the aqueous electrolytes is limited by the water decomposition. Thus, we integrated the two types of devices by leveraging Zn–I₂ and Zn–MnO₂ charge/discharge platforms, circumventing water decomposition limits. Furthermore, the iodide additive can recover dead MnO₂. Overall, a flexible thin-film device that combines supercapacitor, Zn–I₂ and Zn–MnO₂ functionalities was demonstrated (Figure 8g). This integration outperformed individual components in capacity, shelf life, and cycle stability [67].

4. Outlooks and Perspectives

To reveal detailed information about the reviewed battery structures,

Table 1. Reported cathode-free zinc electrolytic MnO₂ batteries: anode and cathode materials, electrolyte composition, device structure, and highest area capacities. The additives are marked in bold.

Anode	Cathode	Electrolyte	Device Structure	Capacity (mAh/cm2)	Ref.
Pt/C	Carbon felt	1 M MnSO4 + 0.05 M H2SO4	Swagelok cell	4	[24]
			Cylindrical cell	2
Cu plate	Carbon cloth	0.3 M CuSO4 + 0.3 M MnSO4 + 0.5M H2SO4	Pouch cell	0.8	[22]
Carbon cloth		0.1 M Bi(NO3)3 + 0.1 M MnSO4 + 1M HNO3
Zn plate		0.3 M ZnSO4 + 0.3 M MnSO4
PbSO4/Pb	Carbon felt	0.5 M H2SO4 + 1 M MnSO4	Pouch cell	24.4	[26]
			Flow cell	4.88
Carbon cloth	Carbon cloth	0.5 M CdSO4 + 0.5 M MnSO4 + 0.5 M H2SO4	Home-made cell	10	[23]
		0.5 M CuSO4 + 0.5 M MnSO4 + 0.5 M H2SO4		0.5
		1 M ZnSO4 + 1 M MnSO4 + 0.1 M H2SO4		0.5
Zn foam	Carbon cloth	1 M ZnSO4 + 1 M MnSO4 + 0.1 M H2SO4	Home-made cell	10	[20]
Zn foil	Carbon felt	1 M ZnSO4 + 1 M MnSO4 + 0.1 M H2SO4 + 0.07 mM polvinylpyrrolidone (PVP)	Beaker cell	4	[27]
Zn foil	Carbon felt	1 M ZnCl2 + 1 M MnCl2 + 0.2 M HCl	Home-made cell	0.5	[28]
Zn foil	Carbon felt	1 M ZnSO4 + 1 M MnSO4 + 0.2 M HAc	Pouch cell	2	[29]
Zn foil	Carbon felt	1 M ZnAc2 + 0.5 M HAc+0.5 M MnAc2 + 1 M KBr	Beaker cell	17	[31]
Zn plate	Carbon felt	1 M MnAc2 + 1 M ZnO + 2.5 M MSA + 1 M KCl	Beaker cell	2	[32]
3D porous Zn @carbon-felt	Carbon felt	3 M NaOH + 0.3 M ZnO/bipolar membrane/ 3 M MnSO4 + 0.3 M H2SO4 + 0.06 M NiSO4	Home-made cell	13	[36]
Zn foil	Carbon cloth	2.4 M KOH + 0.1 M ZnAc2/bipolar membrane/ 1 M MnSO4 + 0.5 M H2SO4	Home-made cell Flow cell	15	[37]
Zn foil	MnO2 loaded carbon felt	6 M KOH + 0.2 M ZnO + 5 mM vanillin //0.1 M K2SO4//3 M H2SO4 + 0.1 M MnSO4	Pack	66.3	[38]
Zn metal on carbon felt	Bi2O3 coated carbon felt	4 M NaOH + 0.1 M ZnO + 5 mM vanillin //3 M Li2SO4//0.1 M MnSO4 + 2.5 M H2SO4 + 0.03 M NiSO4 + 0.01 M MgSO4	Flow cell	22.6	[39]
Zn foil	Carbon felt	0.8 M ZnSO4 + 0.2 M ZnAc2 + 0.1 M NaCl //1 M MnSO4 + 0.15 M H2SO4 + 0.1 M NaCl	Beaker cell	7.3	[40]
Zn foil	Carbon paper	0.5 M ZnSO4 + 1 M LiTFSI /PILG/1 M MnSO4 + 1 M HTFSI	Home-made cell	18	[41]
Zn foil	MnO2 loaded carbon felt	1 M NaOH + 0.01 M ZnAc2/Na+-form Nafion membrane/2 M ZnSO4 + 0.1 M MnSO4	Glass cylinders	-	[42]
Zn foil	Carbon cloth	2 M ZnSO4 in PAM /12 M MnSO4 + 1 M H2SO4 in PAM	Simple pouch	0.7	[44]
Zn foil	Carbon cloth	3 M KOH + 0.3 M ZnO/NaSO4 agar /3 M MnSO4 + 0.3 M H2SO4	H-shaped electrolytic cell	1.35	[46]
Zn foil	Carbon felt	1 M ZnSO4 + 1 M MnSO4	Flow cell	2	[17]
Zn foil	Carbon cloth	1 M ZnSO4 + 2 M MnSO4	Beaker cell	2	[49]
Zn foil	Carbon cloth	1 M ZnSO4 + 0.5 M MnSO4 + 7 M methylurea	Open thin-film cell	0.5	[50]
Zn foil	Carbon cloth	1 M ZnSO4 + 0.1 M MnSO4 + 3 M Glycine	Coin cell	0.5	[56]
Zn foil	Carbon nanotube film	1 M ZnSO4 + 0.1 M MnSO4 + 0.2 g/mL ZnO	Coin cell	2.5	[53]
		+0.25 g/mL ZnO	Pouch cell	0.46
ZnO@C	Carbon cloth	ZnSO4+PVA/1 M MnCl2 + 1 M ZnSO4 + H2SO4	-	0.5	[57]
Zn foil	Graphite foil	3 M ZnCl2 + 0.1 M MnCl2	Home-made cell	10	[60]
Zn foil	Carbon cloth	1 M ZnAc2 + 0.4 M MnAc2	Coin cell	0.5	[54]
			Swagelok cell	1
Zn on graphite felt	Carbon black on graphite	1.5 M MnAc2 + 1.5 M ZnCl2 + 3 M KCl	Flow cell	20	[61]
Zn foil	Carbon felt	1 M ZnAc2 + 1 M MnAc2 + 2 M KCl + 0.1 M KI	Flow cell	20	[63]
Zn foil	Carbon felt	1 M MnSO4 + 1 M ZnSO4 + 0.03 M ZnI2	Home-made cell	20	[64]
Zn foil	Carbon felt	0.5 M HAc + 0.5 M MnAc2 + 1 M ZnAc2 + 1 M KBr + 0.1 M KI	Home-made cell	20	[65]
Zn foil	Carbon felt	1 M ZnSO4 + 1 M MnSO4 + 0.2 M H2SO4 + 0.05 M ZnBr2 + 0.2 M Br2	Home-made cell	13.3	[66]
Zn foil	Carbon cloth Carbon felt	0.5 M ZnCl2 + 0.5 M MnAc2 + 2 M KCl + 1.75 M HAc + 0.05 M CrCl3	Pouch cell	5	[55]
Printed zinc	Printed graphite	1 M ZnAc2 + 1 M MnAc2 + 2 M KCl + 0.1 M KI	Thin-film cell	0.1	[68]
Zn@Pb	Graphite felt	0.2/0.1 M H2SO4 + 1 M ZnSO4 + 1 M MnSO4 + PbAc2	Home-made cell	5	[69]
Phytic acid modified Zn	Carbon cloth	1 M ZnSO4 + 0.5 M MnSO4 + H3PO4	Pouch cell	10	[70]
Zn foil	PMAA@CC	1 M ZnAc2 + 0.4 M MnAc2	Pouch cell	1	[71]
Zn foil	PAAMn@SS	1 M ZnSO4 + 0.5 M MnSO4	-	0.1	[72]
Zn foil	BiMOF@Carbon paper	1 M ZnSO4 + 0.2 M MnSO4	-	0.17	[76]
Zn foil	Graphite foil	0.25 M Al2(SO4)3 + 2 M ZnSO4 + 0.5 M MnSO4	Home-made cell	2	[77]
Cu foil	Carbon felt	1 M ZnSO4 + 0.2 M MnSO4 + 0.1 mM surfactant	Coin cell	3	[78]
Printed graphite	Printed graphite	1 M ZnAc2 + 1 M MnAc2 + 2 M KCl + 0.1 M KI	Thin-film cell	0.1	[21]
Printed zinc	Printed activated carbon	1 M ZnAc2 + 1 M MnAc2 + 2 M KCl + 0.1 M KI	Thin-film cell	0.5	[67]

lists the electrode materials and electrolyte formulations. By categorizing recent strategies for performance and stability enhancement into electrolyte and electrode modifications, we propose the following possible research directions:

Electrolyte pH: Although initial electrolytes may be neutral, the charging process will generate H⁺, which inevitably acidifies the electrolyte. This acidic environment accelerates Zn anode corrosion and hydrogen evolution, leading to gas-induced battery inflation.

Electrolyte Volume: While electrolytes serve as material sources for Zn//MnO₂ deposition, excessive volumes (taking 30 µL/cm² for thin-film battery as a reference) reduce material utilization efficiency. Optimization should balance electrolyte quantity with deposition kinetics.

Zn Anode: Despite advances in Zn alloys and SEI layers for widely studied zinc–ion batteries (mostly based on mild acid sulphate electrolytes), their application in electrolytic MnO₂ systems (mostly based on in acid or acetate electrolytes) remains unexplored. Translating these innovations may mitigate corrosion and hydrogen evolution.

MnO₂ Cathode: Industrial electrolytic manganese dioxide (EMD) dominates the primary market. Studying doping and phase transitions during in situ MnO₂ electrodeposition may help to refine EMD production processes.

Current Collectors: Excess Zn addition compensates for anode loss but obscures the reaction mechanisms. Single or dual electrode-free configurations are recommended as control experiment samples. Additionally, alignment between real-world application conditions and laboratory battery designs should be considered. For instance, carbon cloth/felt’s high surface area differs significantly from printed electrodes, potentially limiting practical relevance.

Battery Design: Thin-film batteries with lean electrolytes and compact dimensions closely mimic real-world applications. These structures are highly sensitive to hydrogen accumulation and have been commercialized (e.g., Zinergy’s products).

Shelf Life: Long-term stability tests are critical, especially for systems employing mediator additives prone to shuttle reactions.

Integrated Devices: Activated carbon electrodes enhance MnO₂ anchoring and contribute pseudocapacitive capacity. Hybrid supercapattery designs (combining supercapacitors with Zn–MnO₂/I₂ chemistries) offer superior performance and novel mechanistic pathways.