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Review

A Short Review: Comparison of Zinc–Manganese Dioxide Batteries with Different pH Aqueous Electrolytes

by
Ramona Durena
* and
Anzelms Zukuls
Institute of Materials and Surface Engineering, Faculty of Materials Science and Applied Chemistry, P. Valdena 3/7, LV-1010 Riga, Latvia
*
Author to whom correspondence should be addressed.
Batteries 2023, 9(6), 311; https://doi.org/10.3390/batteries9060311
Submission received: 12 May 2023 / Revised: 23 May 2023 / Accepted: 29 May 2023 / Published: 5 June 2023
(This article belongs to the Special Issue Advanced Materials for Zinc‐Based Battery: Development and Challenges)
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Abstract

:
As the world moves towards sustainable and renewable energy sources, there is a need for reliable energy storage systems. A good candidate for such an application could be to improve secondary aqueous zinc–manganese dioxide (Zn-MnO2) batteries. For this reason, different aqueous Zn-MnO2 battery technologies are discussed in this short review, focusing on how electrolytes with different pH affect the battery. Improvements and achievements in alkaline aqueous Zn-MnO2 batteries the recent years have been briefly reviewed. Additionally, mild to acidic aqueous electrolyte employment in Zn-MnO2 batteries has been described, acknowledging their potential success, as such a battery design can increase the potential by up to 2 V. However, we have also recognized a novel battery electrolyte type that could increase even more scientific interest in aqueous Zn-MnO2 batteries. Consisting of an alkaline electrolyte in the anode compartment and an acidic electrolyte in the cathode compartment, this dual (amphoteric) electrolyte system permits the extension of the battery cell potential above 2 V without water decomposition. In addition, papers describing pH immobilization in aqueous zinc–manganese compound batteries and the achieved results are reported and discussed.

Graphical Abstract

1. Introduction

Effective and reliable battery technology is highly desired as the world leans toward sustainable energy sources, such as hydro, solar, and wind energies. However, all these options suffer from uneven energy production patterns throughout the day and year and cannot supply the energy on demand [1]. One way to overcome these shortcomings is to store the excess energy in batteries or create more everyday appliances with accumulators. This would allow us to charge up these devices when energy is available, such as in the daytime, from solar panels and use electricity-powered devices in the evening [2].
Li-ion batteries (LIBs) in recent decades have found wide applications in consumer electronics [3], such as computers, cell phones [4], and now also electric and hybrid vehicles [5]. The extensive use of LIBs can be attributed to their high energy density (up to 500 Wh/L), power density (up to 300 W/kg), high theoretical lithium storage capacity (890 mA h/g), operating voltage (>3.7 V), excellent cycling performance (over 10,000 cycles), and low self-discharge (<5% of the stored capacity over 1 month) [6,7,8,9]. However, the optimal operating range of LIBs is 15–35 °C. Furthermore, battery overheating can induce the melting of the separator, which causes an internal short circuit. In this short-circuit spot, a rapid self-heating begins at a rate of around 2592.0 to 11,860.0 °C min−1 [10], resulting in a thermal runaway [11,12]. This means that once the process has started, no cooling application can stop it. Unfortunately, this process results in a fire or an explosion, thus creating a considerable safety risk for consumers [4,13,14]. Furthermore, another significant drawback has emerged: a material shortage of commonly used metals in LIBs could be on the rise [15,16,17]; thus, this suggests that new or improved battery technology is needed.
One such technology could be improved aqueous Zn-MnO2 batteries. They do not pose large safety risks; the two-electron reaction of Zn2+/Zn has a high theoretical capacity of 820 mAh/g, Zn has low toxicity, and is earth-abundant [18,19]. Moreover, MnO2 as a cathode has an attractive total theoretical capacity of 617 mAh/g. Thus, making Zn-MnO2 cells with an energy density of >400 W h/L and >500 Wh/kg [20], which is comparable to some LIBs, is an interesting technology to improve large-scale rechargeable battery applications.
The origin of aqueous Zn-MnO2 batteries can be traced back to 1866 when the French scientist Georges-Lionel Leclanché created and patented the first Zn-MnO2 battery, which was called the Leclanché cell [21]. Improving the electrolyte composition and design of this battery cell has led to the creation of commercialized dry cells. Further development of Zn-MnO2 batteries took place in the 1950s when the Canadian engineer Lewis Urry improved and rebuilt the battery cell known at that time. In 1960, a patent was granted for aqueous alkaline Zn-MnO2 batteries [22,23,24]. These batteries have continued their development until now, improving the composition and shelf life, and replacing harmful elements inside the batteries, such as mercury. This has led aqueous alkaline Zn-MnO2 batteries to be the longest and most widely used primary battery technology that still dominates the global market as the primary battery source nowadays [25].
Aqueous alkaline Zn-MnO2 batteries consist of a zinc metal anode, a manganese dioxide powder mixture with finely dispersed carbon as the cathode, and concentrated potassium hydroxide (>30 wt%) aqueous solution as the electrolyte. In this electrode pairing (as depicted in Figure 1), Zn acts as an anode with a standard potential of −1.199 V vs. SHE (standard hydrogen electrode) and undergoes reaction (1), whereas MnO2 acts as an active material in the cathode with a standard potential of +0.250 V vs. SHE and undergoes reactions (2) and (3) in the primary cells. As MnO2 is non-conductive, usually a small amount of finely dispersed carbon is added to ensure the conductivity of the cathode. This results in an overall reaction (4) [1,26,27,28]. The calculated open circuit potential (OCP) for this system of 1.45 V is in discrepancy with the actual commercially available battery OCP of 1.55–1.6 V. This effect can be attributed to the fact that reactions (1)–(4) are very general and the actual system is quite complex [29,30,31,32]. Additionally, the total theoretical capacity of the MnO2 cathode consists of a two-stage (two-electron) reaction. The first electron reaction (2) of the MnO2 cathode in an alkaline medium corresponds to the intercalation of hydrogen in the solid phase with a theoretical capacity of 309 mAh/g. As the discharge continues, the second electron reaction (3) can also occur. A dissolution–precipitation heterogeneous process with a theoretical capacity of 308 mAh/g continues to enter the deep discharge stage of the cathode [20,28,33].
Anode:
Zn + 4(OH)− → Zn(OH)42− + 2e−              E0 = −1.199 V vs. SHE
Cathode:
MnO2 + H2O + e− → MnOOH + OH−        (1st e−)        E0 = +0.250 V vs. SHE
MnOOH + H2O + e− → Mn(OH)2 + OH−       (2nd e−)             
Overall:
Zn + 4(OH)− + MnO2 + 2H2O → Zn(OH)42− + Mn(OH)2       E = 1.449 V
Although researched for a very long time now, alkaline Zn-MnO2 batteries are still of great interest to scientists. Aqueous alkaline electrolyte Zn-MnO2 batteries have proven to be a great source for primary batteries; however, their secondary battery performance is very limited due to the formation of insoluble compounds (usually in the second electron reaction of the cathode). For this reason, many scientists are devoting their research activities to developing rechargeable aqueous Zn-MnO2 secondary batteries with higher specific capacities and OCP. In this short review, we discuss improvements and achievements in recent years for different aqueous Zn-MnO2 battery types, as shown in Figure 2. We have discussed alkaline and acidic/mild aqueous electrolyte rechargeable Zn-MnO2 battery types and also recognized a dual (amphoteric) aqueous electrolyte battery concept that could increase research interest in aqueous Zn-MnO2 batteries.

2. Alkaline Electrolyte Zn-MnO2 Batteries

Many authors have attributed the poor rechargeability of alkaline Zn-MnO2 batteries to the formation of electrochemically inactive manganese oxides such as Mn3O4 and ZnMn2O4. These oxides form during deep discharge that takes place while accessing the second electron capacity of MnO2 [22,34]. After several charge/discharge cycles, the irreversible manganese oxides coated the active MnO2 particles, preventing them from undergoing further reactions. N. D. Ingale with co-authors [35] investigated this problem and provided proof for the connection between insoluble compound formation and the depth of discharge (DOD). They experimentally obtained 3000 cycles for alkaline (45 wt% KOH water solution) Zn-MnO2 batteries at 10% DOD. However, at 20% DOD, only ~500 cycles could be obtained. For these experiments, the DOD was based on the first theoretical electron capacitance of the MnO2 cathode (308 mAh/g).
Other scientists have tried to improve the rechargeability of aqueous alkaline Zn-MnO2 batteries by using additives to electrolytes and/or MnO2 cathodes [7,34]. A.M. Bruck with co-workers [36] investigated the effects of Bi2O3 addition to the MnO2 electrode. They found that Bi3+ occupies an interstitial position in the layered MnO2 phase—birnessite—limiting the Mn3+ ion diffusion in the lattice. Thus, Bi3+ promotes the conversion from the MnO2 birnessite phase directly to Mn(OH)2 and prevents Mn2O3 phase formation during the battery charge/discharge process. A similar technique has been employed by G.G. Yadav with co-authors [28]. They used Cu-intercalated birnessite mixed with Bi2O3 as a cathode in alkaline Zn batteries and obtained capacity reversibility (from the second electron capacity of the MnO2 cathode—617 mAh/g) for more than 6000 cycles. However, the tested cathode material contained only ~19% of the active MnO2 material.
Other authors have investigated the possibility of improving the alkaline aqueous Zn-MnO2 battery system by modifying the electrolytes [34,37,38,39]. For example, B.J. Hertzberg with co-authors [34] researched the effects of mixing KOH electrolyte with LiOH addition. They found that the insertion of lithium or a combined lithium proton is the dominant reduction mechanism for MnO2. This prevents Zn poisoning reactions and increases the battery cycle lifetime up to 60 cycles with negligible capacity loss (cycling around 360 mAh/g). However, despite their rechargeability improvements, aqueous alkaline Zn-MnO2 batteries still have limited applications due to their low cell potential of ~1.4 V. For this reason, other scientists have gone one step ahead and changed the electrolyte even further, excluding hydroxides entirely.

3. Mild to Acidic Electrolyte Zn-MnO2 Batteries

To overcome the low cell potential of 1.4 V and improve the cyclability of aqueous Zn-MnO2 batteries, other scientists have changed the pH of the electrolyte, thus altering the reaction mechanism of the cell (as depicted in Figure 3). This was achieved by replacing standard hydroxide solutions such as KOH, NaOH, and LiOH [27,34,35,40,41] with ZnSO4, MnSO4 (K2SO4; Na2SO4) [42,43,44,45,46,47] salts, and/or diluted acid (H2SO4) [48,49]. As shown in the Pourbaix diagram of manganese oxide compounds (Figure 4a), at higher pH values, MnO2 goes through a low potential reaction (0.250 V vs. SHE) with the possibility of transforming into electrochemically inactive manganese oxides such as Mn3O4. However, by lowering the pH value well into the acidic medium, the oxygen atom from the MnO2 cathode can combine with the hydrogen ions to create water; thus, Mn4+ ions can be reduced to Mn2+ and dissolved in an acidic medium without electrochemically inactive manganese oxide creation. Additionally, the reaction potential of the cathode is significantly improved above 1.2 V vs. SHE by placing MnO2 in an acidic medium. On the other hand, this change in the pH of the electrolyte also increases the reaction potential of the Zn anode, which lowers the overall potential of the cell. As seen in the Pourbaix diagram of zinc oxide compounds (Figure 4b), this potential increase is quite notable, from −1.199 V vs. SHE in an alkaline medium to −0.762 V vs. SHE in an acidic medium.
Despite the increase in the reaction potential of the Zn anode, this cell configuration allows the cathode to undergo a higher operating potential reaction, resulting in an overall potential increase up to ~2 V, according to reactions (5)–(7) [52]:
Anode:
Zn → Zn2+ + 2e−                 E0 = −0.762 V vs. SHE
Cathode:
MnO2 + 4H+ + 2e− → Mn2+ + 2H2O             E0 = +1.224 V vs. SHE
Overall:
Zn + MnO2 + 4H+ → Zn2+ + Mn2+ + 2H2O              E = 1.986 V
However, this increased cell potential comes with a considerable drawback. In an acidic environment, a parasitic reaction can take place at the Zn anode. A corrosion reaction can occur, the Zn anode dissociates into the electrolyte as Zn2+, and a hydrogen gas H2 is produced. This corrosion reaction can not only decrease the capacity but also lead to cell rupture [53]. A. Mitha with co-workers [54] designed Zn-LiMn2O4 batteries with mildly acidic (pH 4) aqueous electrolyte. To overcome gas production on the anode, fumed silica and polyethylene glycol (MW 300 g/mol) were added. Both additives acted as Zn anode corrosion inhibitors and dendrite suppressors. Thus, this led to achieving a 27–40% lower corrosion current density on the Zn anode and performing 1000 charge–discharge cycles, cycling at a capacity of ~140 mAh/g and retaining ~65% of its initial capacity after 1000 cycles.
H. Pan with co-authors [55] also reported mild electrolyte aqueous Zn-MnO2 batteries with 5000 charge–discharge cycles and a capacity retention of 92% (based on the first electron capacity of the MnO2 cathode—308 mAh/g). The high battery cyclability and capacity retention were attributed to the addition of 0.1 M MnSO4 to 2 M ZnSO4 electrolyte. The Mn2+ ion suppresses MnO2 dissolution into the electrolyte, thus improving the stability of the cathode and the utilization of the MnO2 active material. This idea has been well adapted in the literature [43,44,45,46]. However, C. Qiu with co-authors [56] addressed the possibility of adding Mn2+ deposition to the MnO2 cathode while cycling and adding extra capacity to the cathode. Cells with different Mn2+ concentrations were tested, and as expected, all the cells showed a stable capacity for a certain number of cycles, followed by fast capacity fading due to the consumption of Mn2+ ions from the electrolyte. Although effective for a number of cycles, the Mn2+ additive may not be an effective strategy for improving the charge–discharge stability. Furthermore, C. Qiu with co-authors again provided proof for the formation of ZnMn2O4 at the cathode during discharge, which is attributed to the battery capacity fading and poor cyclability. This indicates that the same rechargeability issues exist in mild aqueous Zn-MnO2 cells under alkaline conditions. Nonetheless, there could be other methods to improve aqueous Zn-MnO2 batteries and extend the water decomposition potential window over 1.23 V [57] by combining scientific advances made in aqueous alkaline and acid electrolyte systems.

4. Dual/Amphoteric Electrolyte Zn-MnO2 Batteries

Recently, another type of Zn-MnO2 battery has been adapted, which takes advantage of the increased MnO2 potential in acidic electrolytes without creating drawbacks to the Zn anode. This has been realized by placing the Zn anode in an alkaline electrolyte and the MnO2 cathode in an acidic electrolyte, as shown in Figure 5. By creating such a unique battery cell type, it is possible to increase the OCP to 2.45 V according to reactions (8)–(10) [32,58,59].
Anode:
Zn + 4(OH)− → Zn(OH)42− + 2e−              E0 = −1.199 V vs. SHE
Cathode:
MnO2 + 4H+ + 2e− → Mn2+ + 2H2O              E0 = +1.224 V vs. SHE
Overall:
Zn + 4(OH)− + MnO2 + 4H+ → Zn(OH)42− + Mn2+ + 2H2O      E = 2.423 V
The increased theoretical battery potential and widened water decomposition window can also be explained with Pourbaix diagrams. As shown in Figure 6, by operating the battery in a constant pH electrolyte, the maximum water–electrolyte operation potential window is ~1.23 V. By operating a water-based electrolyte battery in a wider potential window, oxygen evolution reaction (OER) and hydrogen evolution reaction (HER) can take place. If a dual-electrolyte system is used and the cathode is operated at lower pH values and the anode at higher pH values, as shown in Figure 6, the stable electrochemical window of water could be increased to ~2 V. However, the HER has a high overpotential for zinc, which can be explained by the Tafel Equation (11):
η = α + b log i
where η is the HER overpotential, α is the transfer coefficient (overpotential when current density i is equal to the unit current density), b is the Tafel slope (a constant), and i is the current density [60,61]. In addition, OER has a significant intrinsic overpotential of at least 0.37 V due to the complex reaction mechanism [62] and slow reaction kinetics [63]. This way, the water–electrolyte operating potential significantly improves up to 3 V, and the OER and HER reactions are prohibited, thus paving the way for high-voltage aqueous batteries.
Overall, this idea of creating electrolytes with different pH values and increasing the supposed half-reaction potentials as well as the origin of the water decomposition potential origin can be found described in 2005 by three non-related articles. S. Hasegawa with co-authors [64] demonstrated an H2O2 fuel cell with acidic (H2SO4) catholyte and alkaline (NaOH) anolyte. J.L. Cohen with co-authors [65] published findings on H2/O2 fuel cells, where acidic (0.1 M H2SO4) and alkaline (0.1 M KOH) electrolytes were used. Additionally, in the same year, E.R. Choban with co-authors [66] presented a methanol/O2 fuel cell with acidic (0.5 M H2SO4) and alkaline (1 M KOH) electrolytes. Later this dual-electrolyte (acidic and alkaline) idea was adapted to different electrode batteries such as Al-air [67,68,69,70], Zn-air [68,71,72], Mg-air [68,73], Zn-PbO2 [74,75,76], Zn-Br2 [77], etc. However, the battery working time and efficiency of this battery cell type are limited by both the electrolyte neutralization reaction and ion diffusion. Scientists have tried to overcome this limitation by using membranes and/or gelation of electrolytes.
The first zinc–manganese compound-based battery using an acid-alkaline dual electrolyte was presented by L. Chen with co-authors [78]. The battery cell consisted of a Zn metal sheet placed in an alkaline anolyte (KOH + LiOH solution) against Ti mesh (current collector) placed in an acidic KMnO4 (active cathode material) catholyte. The battery exhibited an astonishing OCP of 2.8 V and a capacity of up to 750 mAh/g (discharge current density of 375 mA/g). However, no charge–discharge cycles were observed. Later, G. G. Yadav with co-authors [79] showed a two-compartment Zn-MnO2 cell in 2019. The Zn anode compartment electrolyte consisted of a 45% KOH aqueous solution gelled in potassium polyacrylate. The MnO2 cathode compartment electrolyte consisted of either KMnO4 or MnSO4 in a diluted H2SO4 solution. Both compartments were separated using a cellophane membrane, as shown in Figure 7a. This approach enabled achieving cell potential up to 2.45 V (MnSO4) and 2.8 V (KMnO4). Battery cells with KMnO4 in the catholyte were stable, with capacity retention for 35 charge/discharge cycles when utilized at 20% (~62 mAh/g) of MnO2 one-electron capacity. Moreover, the battery cells with MnSO4 in the catholyte were stable for 120 charge/discharge cycles when utilized at a capacity of 20% (~62 mAh/g). Unfortunately, this article does not present measurements that would exceed the 35 h mark. A similar approach to Zn-MnO2 cells has been shown by C. Liu with co-authors [80] in 2020. As depicted in Figure 7b, the cell consisted of two electrolyte compartments: 2.4 M KOH + 0.1 M Zn(CH3COO)2 anolyte and 0.5 M H2SO4 + 1.0 M MnSO4 catholyte. A bipolar membrane was used to prevent ion migration. With this design, a cell potential of ~2.4 V was demonstrated, a Coulombic efficiency of 98.4% was reached, and 1500 cycles were performed while utilizing 0.4 mAh/cm−2 areal capacity.
Other authors have tried to improve this concept by adding one more (neutral) compartment between the alkaline and acidic electrolytes. C. Zhong with co-authors [32] in 2020 showed a rechargeable Zn-MnO2 battery as shown in Figure 7c. The battery cell consisted of three electrolyte compartments: (1) acidic 3M H2SO4 + 0.1 M MnSO4 catholyte; (2) neutral 0.1 M K2SO4 electrolyte; (3) alkaline 6 M KOH + 0.2 M ZnO + 5 mM vanillin anolyte. The acidic and alkaline compartments were separated from the neutral compartment using anion- and cation-exchange membranes. In this system, an astonishing OCP of 2.83 V was recorded. Additionally, deep discharge (616 mAh/g) cycling was performed for over 200 h, with 2% capacity fading. However, it should be noted that ion-selective membranes are quite expensive.
More recently, in 2022, A. Zukuls with co-authors [59] demonstrated a rechargeable aqueous Zn-MnO2 battery in which the mixing of anolyte and catholyte was prevented by gelling both electrolytes. As shown in Figure 7d, the battery consisted of an acidic (0.5 M H2SO4) Pluronic® F-127 aqueous catholyte and an alkaline (1.0 M KOH) Pluronic® F-127 aqueous anolyte. No specific membranes were used in this design, only a porous paper separator was used for visual purposes. With this design, 200 charge–discharge cycles were performed (cycling ~10% of MnO2 one-electron capacity) while maintaining a high and stable OCP of ~2.4 V. Additionally, the impact of a neutral layer (0.5 M K2SO4 + Pluronic® F-127) between the catholyte and anolyte was tested. With this configuration, 200 cycles (cycling ~10% of MnO2 one-electron capacity) were also performed while maintaining a high and stable OCP of ~2.4 V. However, the discharge and charge potentials were 0.5 V lower and higher, respectively. Although in the design with the neutral layer, a larger internal resistance was observed due to an increase and decrease in the charge–discharge potential, the battery was overall more stable. In a two-layer electrolyte, the battery was stable only for 25 h; however, in a three-layer electrolyte, the battery was stable for 30 h. After this time, a rapid neutralization reaction took place and battery failure was observed. Even though this battery design did not use expensive ion-exchange membranes, thus reducing the expected price, the design did not provide long-term use. Other studies describing zinc–manganese compound-based batteries with dual (amphoteric) electrolytes are listed in Table 1.

5. Challenges and Outlook

Primary aqueous alkaline Zn-MnO2 battery technology already dominates the primary battery market. This technology has proven itself as safe, affordable, and reliable. As the world moves towards renewable energy sources, we are in need of safe, affordable, and reliable secondary battery technology with a high-power density and excellent capacity retention during charge/discharge cycles. A good candidate for this application could be aqueous Zn-MnO2 battery technology. However, several improvements and new technological approaches are needed. A very new approach has been demonstrated by G. G. Yadav with co-authors [79] and by others [32,59,80]. In the literature, dual (amphoteric) electrolyte aqueous cells have been demonstrated that result in a significant increase in cell potential over 2 V. This voltage increase has been achieved in these new aqueous cell types by placing a Zn anode in an alkaline medium and an MnO2 cathode in an acidic medium. To prevent neutralization reactions from occurring, different approaches have been used, such as the usage of ion exchange membranes or other types of membranes, gelation of electrolytes, or both. However, when put to the test, this cell design presents several problems that should be addressed in the near future by researchers. Different recyclability problems continue to exist for the Zn anode in an alkaline medium and the MnO2 cathode in an acidic medium. For aqueous Zn-MnO2 batteries to compete with LIBs, full utilization of the MnO2 cathode capacity should be achieved with high-capacity retention over a large number of charge–discharge cycles. Additionally, the economic aspects should be taken into account when considering new types of chemical energy storage systems. Although ion exchange membranes provide excellent immiscibility with the electrolyte pH, they are quite expensive. Different hydrogels such as Pluronic® F-127 are economically disadvantageous as well. Thus, effective methods for pH immobilization should be more thoughtfully considered.

6. Conclusions

In this short review, aqueous Zn-MnO2 batteries with different pH electrolytes have been acknowledged. The historical origins of alkaline Zn-MnO2 batteries can be traced back to the middle of the 20th century. Although they have been successfully utilized in the primary battery market and research has been carried out to tackle the rechargeability challenges, little to no secondary aqueous Zn-MnO2 batteries are seen in the market. Another type of Zn-MnO2 battery discussed is a mild to acidic aqueous electrolyte battery. If an alkaline-type battery suffers from electrochemically inactive and insoluble compound formations, such as Mn3O4 and ZnMn2O4, then, in an acidic electrolyte, this problem is partially solved since these compounds are soluble in acid. Moreover, by decreasing the pH value of the electrolyte, the cathode goes through a higher potential reaction and increases the overall potential of the cell from 1.4–1.6 V in the alkaline electrolyte to 1.9–2.0 V in the mild/acidic electrolyte. However, more recently, the scientific community has shown a new and innovative battery cell type that combines the positive effects of alkaline and acidic electrolytes and advances the possibility of creating a water-electrolyte battery with a potential of up to 3 V. This has been accomplished by operating the anode in an alkaline medium and cathode in an acidic medium. Thus, increasing the potential of this aqueous dual/amphoteric electrolyte Zn-MnO2 battery up to 2.4 V improves the reachability of the cell. However, this unique battery cell imposes new challenges, such as neutralization reactions between both electrolytes.

Author Contributions

Conceptualization, R.D. and A.Z.; methodology, R.D. and A.Z.; validation, A.Z.; formal analysis, R.D. and A.Z.; investigation, R.D.; writing—original draft preparation, R.D.; writing—review and editing, A.Z.; visualization, R.D.; supervision, A.Z.; funding acquisition, R.D. All authors have read and agreed to the published version of the manuscript.

Funding

Ramona Durena acknowledges that this work has been supported by the European Social Fund within Project No 8.2.2.0/20/I/008 «Strengthening of PhD students and academic personnel of Riga Technical University and BA School of Business and Finance in the strategic fields of specialization» of the Specific Objective 8.2.2 «To Strengthen Academic Staff of Higher Education Institutions in Strategic Specialization Areas» of the Operational Programme «Growth and Employment».

Institutional Review Board Statement

Not applicable.

Informed Consent Statement

Not applicable.

Data Availability Statement

Not applicable.

Conflicts of Interest

The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper.

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Figure 1. The schematic working mechanism of alkaline Zn-MnO2 batteries.
Figure 1. The schematic working mechanism of alkaline Zn-MnO2 batteries.
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Figure 2. Different types of Zn-MnO2 batteries with their possible potentials and the most commonly used electrolytes.
Figure 2. Different types of Zn-MnO2 batteries with their possible potentials and the most commonly used electrolytes.
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Figure 3. The schematic working mechanism of mild/acidic Zn-MnO2 batteries.
Figure 3. The schematic working mechanism of mild/acidic Zn-MnO2 batteries.
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Figure 4. Pourbaix diagram of (a) manganese compounds (Reprinted with permission from Ref. [50]; 2015, published by Springer Nature Switzerland AG) and (b) zinc compounds in water (Reprinted with permission from Ref. [51]; 2016, published by Elsevier B.V).
Figure 4. Pourbaix diagram of (a) manganese compounds (Reprinted with permission from Ref. [50]; 2015, published by Springer Nature Switzerland AG) and (b) zinc compounds in water (Reprinted with permission from Ref. [51]; 2016, published by Elsevier B.V).
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Figure 5. The schematic working mechanism of dual/amphoteric electrolyte Zn-MnO2 batteries.
Figure 5. The schematic working mechanism of dual/amphoteric electrolyte Zn-MnO2 batteries.
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Figure 6. Pourbaix diagram of water.
Figure 6. Pourbaix diagram of water.
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Figure 7. Design schemes of different aqueous amphoteric (dual) electrolyte Zn-MnO2 batteries: (a) experimental setup of G. G. Yadav (Reprinted with permission from Ref. [79]. Copyright 2019 American Chemical Society.); (b) experimental setup of C. Liu (Reproduced with permission from Ref. [80]; published by John Wiley & Sons, Inc., Hoboken, NJ, USA, 2020); (c) experimental setup of C. Zhong (Reproduced with permission from Ref. [32]; published by Springer Nature Limited, 2020); (d) experimental setup of A. Zukuls (Reproduced with permission from Ref. [59]; published by Elsevier B.V, 2022).
Figure 7. Design schemes of different aqueous amphoteric (dual) electrolyte Zn-MnO2 batteries: (a) experimental setup of G. G. Yadav (Reprinted with permission from Ref. [79]. Copyright 2019 American Chemical Society.); (b) experimental setup of C. Liu (Reproduced with permission from Ref. [80]; published by John Wiley & Sons, Inc., Hoboken, NJ, USA, 2020); (c) experimental setup of C. Zhong (Reproduced with permission from Ref. [32]; published by Springer Nature Limited, 2020); (d) experimental setup of A. Zukuls (Reproduced with permission from Ref. [59]; published by Elsevier B.V, 2022).
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Table
Table 1. List of papers investigating zinc–manganese compound-based batteries with dual electrolytes.
Table 1. List of papers investigating zinc–manganese compound-based batteries with dual electrolytes.
Ref.Year of PublishingAnode MaterialCathode MaterialAnolyteCatholyteMembranePotentialCyclingMaximum CapacityCapacity RetentionStability
[78]2013Zn-plateKMnO4 dissolved in catholyte (Ti mesh as a current collector)2 M KOH + 2 M LiOH1 M H2SO4 + KMnO4Li-ion exchange membrane (Li1+x+yAlxTi2−xSiyP3−yO12)2.8 V OCP-from 510 mAh/g (discharge current density of 37.5 mA/g)
to 750 mAh/g (discharge current density of 375 mA/g)		-~110 h (with an intermittent supply of KMnO4)
[79]2019Zn-foil85 wt% electrolytic manganese dioxide + 15 wt% multiwalled carbon nanotubes on carbon felt~45 wt% KOH solution in potassium polyacrylate1 M MnSO4 + 0.5 M H2SO4 mixed in ratio of 4:1Four-layers of cellophane2.45 OCP120 cycles (capacity 62 mAh/g) at C/2308 mAh/g at C/2100%35 h
Zn-foil85 wt% electrolytic manganese dioxide + 15 wt% multiwalled carbon nanotubes on carbon felt~45 wt% KOH solution in potassium polyacrylate1 M MnSO4 + 0.5 M H2SO4 mixed in ratio of 1:3Four-layers of cellophane2.8 OCP35 cycles (capacity 62 mAh/g) at C/2308 mAh/g at C/2100%35 h
[80]2020Zn-foilpotentiostatic electrodeposited MnO2 on carbon cloth2.4 M KOH + 0.1 M Zn(CH3COO)0.5 M H2SO4 + 1.0 M MnSO4bipolar membrane (Fumasep FBM)2.42 V at 2 mA/cm21500 cycles (capacity of 0.4 mAh/cm2) at the current density of 2 mA/cm20.4 mAh/cm2 at the current density of 2 mA/cm2~90%200 h
[58]2020Zn-foilcarbon-felt as a cathode-less current-collector3 M NaOH + 0.3 M ZnO3 M MnSO4 + 0.3 M H2SO4 + 0.06 M NiSO4Selective membrane (no more information provided)2.44 at 1 mA/cm (1C)450 cycles (capacity of 1.0 mAh/cm2) at 2C1.0 mAh/cm2 (or ~270 mAh/g) at 1C99.9%140 h
[81]2020Zn-foilMnO2 + acetylene black + PTFE (70:20:10)1 M NaOH + 0.01 M Zn(CH3CO2)22 M ZnSO4 + 0.1 M MnSO4Nafion 1152.2 V90 cycles (capacity 282.2–428.6 mAh/g) at 200 mA/g300 mAh/g at (not specified)~100%Not shown
[82]2021Zn-foilMnO2 powder coated on Ti foil2 M ZnSO4 + H2SO4 gelled in polyacrylamide2 M ZnSO4 (with adjusted pH = 7) gelled in polyacrylamideNeutral layer: 2 M ZnSO4 gelled in sodium polyacrylate1.8 V at 5 A/g5000 cycles (capacity of 150 mAh/g) at 5 A/g516 mAh/g at 0.05 A/g93.18%>6 months
[59]2022Zn-foilMnO2 + carbon black (13:5) on carbon felt1 M KOH Pluronic® F-127 hydrogel0.5 M H2SO4 Pluronic® F-127 hydrogel-2.4 V OCP200 cycles (discharge for 120 s) at 1 mA/s25 mAh/g at 20 mA/g->25 h
[83]2022Zn-foilMoS-MnO2 heterostructure + conductive carbon + PVDF (80:15:5) on carbon cloth1 M Zn(ClO4)2 + acetonitrile + water + H2SO41 M Zn(ClO4)2 + acetonitrile + water + KOHPVA+PVP crosslinked membrane1.9 V (at the start of discharge at 0.2 A/g)5000 cycles (160 mAh/g) at 10 A/g464 mAh/g at 0.2 A/g74%Not shown
[84]2022Zn-foilpotentiostatic electrodeposited MnO2 from catholyte on carbon cloth0.3 M ZnO + 3 M NaOH3 M MnSO4 + 0.3 M H2SO4agar with Na2SO42.5 V at 0.1 C350 cycles (capacity ~550 mAh/g) at 1 C577.8 mAh/g at 1C94%>350 h
[85]2022Zn-foilelectrochemically
deposited MnO2 on carbon paper		0.5 M ZnSO4 + 1 M LiTFSI (bis(trifluoromethane)sulfonimide
lithium salt)		1 M MnSO4 + 1 M HTFSI (bis(trifluoromethanesulfonyl)imide)dual-hydrophobic-induced membrane (consisting of polymer/ionic liquid/graphene mixture)2.05 V at 1 mA/cm2000 cycles (fixed area capacity of 1 mAh/cm2) at 1 mA/cm218 mAh/cm2 at 1 mA/cm298%2275 h
[86]2023Zn-foilcarbon cloth2.4 M KOH + 0.1 M Zn(CH3COO)21.0 M H2SO4 + 0.3 M CuSO4 + 1.0 M MnSO4Cu foil1.84 V at 2.5 mA/cm23500 (capacity ~0.5 mAh/cm2) at 2.5 mA/cm20.5 mAh/cm2 at 2.5 mA/cm2~125% (Coulombic efficiency after 3500 cycles)>107 h
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Durena, R.; Zukuls, A. A Short Review: Comparison of Zinc–Manganese Dioxide Batteries with Different pH Aqueous Electrolytes. Batteries 2023, 9, 311. https://doi.org/10.3390/batteries9060311

AMA Style

Durena R, Zukuls A. A Short Review: Comparison of Zinc–Manganese Dioxide Batteries with Different pH Aqueous Electrolytes. Batteries. 2023; 9(6):311. https://doi.org/10.3390/batteries9060311

Chicago/Turabian Style

Durena, Ramona, and Anzelms Zukuls. 2023. "A Short Review: Comparison of Zinc–Manganese Dioxide Batteries with Different pH Aqueous Electrolytes" Batteries 9, no. 6: 311. https://doi.org/10.3390/batteries9060311

APA Style

Durena, R., & Zukuls, A. (2023). A Short Review: Comparison of Zinc–Manganese Dioxide Batteries with Different pH Aqueous Electrolytes. Batteries, 9(6), 311. https://doi.org/10.3390/batteries9060311

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