Introducing a Concept for Designing an Aqueous Electrolyte with pH Buffer Properties for Zn–MnO2 Batteries with Mn2þ/MnO2 Deposition/Dissolution Oliver Fitz,* Florian Wagner, Julia Pross-Brakhage, Manuel Bauer, Harald Gentischer, Kai Peter Birke, and Daniel Biro 1. Introduction The zinc-ion battery (ZIB) with aqueous electrolytes is gaining increasing signifi- cance in the scientific literature,[1–5] as well as for industrial approaches (e.g.,[6–10]). Within the great variety of active material combinations for the ZIB technology, the acidic aqueous rechargeable zinc–manga- nese dioxide battery (ARZMB) is a promis- ing technology with great attention in the literature.[1–5] The reason behind this devel- opment is the globally increasing awareness with respect to the critical raw material supply for the prominent lithium-ion battery (LIB) technology.[11,12] The need for alternative battery cell chemistries is a contemporary topic for the battery research, especially regarding the different fields of application with their respective (and increasing) demands for the battery capabil- ities.[13] In this regard, the ARZMB can play an important role for the stationary energy storage (SES), e.g., on- and off-grid, home or peak-shaving storage applications.[13–17] Still, the research for ARZMB shows an ongoing and lively debate concerning the reaction mechanism in the aqueous, acidic electrolyte media. Up to now, the reaction mechanisms presented for the ARZMB are mainly as follows: 1) the revers- ible Zn2þ (de-)intercalation;[18–43] 2) accompanied by the Hþ (de-)intercalation;[42–44] and 3) the MnO2/Mn2þ dissolution/ deposition:[29,43–53] a) driven by the proton source available through the Zn/Mn hexa–aqua complexes dissolved in the electrolyte ([Zn(H2O)6] 2þ or [Mn(H2O)6] 2þ);[42,43,53–57] and b) accompanied by the precipitation of zinc hydroxides (ZHs) in the electrode and electrolyte (e.g., zinc hydroxide sul- fate [ZHS] in the presence of ZnSO4). [42–46,58] Recent publications in the field of ARZMB emphasize the major contribution of the MnO2/Mn2þ dissolution/deposition mechanism and its influence on the electrolyte’s pH value.[42,46,53–55,59–70] Therefore, pH additives for the electrolytes were introduced to control the electrolyte’s pH, listed to the best of our knowledge in Table 1.[46,53–55,62,64–66] Interestingly, the issues regarding the stabilization of the electrolyte’s pH and the uncontrolled dissolution phenomena of the cathode active material are also a field of research within O. Fitz, F. Wagner, M. Bauer, H. Gentischer, D. Biro Battery Cell Technology Department of Electrical Energy Storage Fraunhofer Institute for Solar Energy Systems ISE 79110 Freiburg, Germany E-mail: oliver.fitz@ise.fraunhofer.de J. Pross-Brakhage, K. P. Birke Electrical Energy Storage Systems Institute for Photovoltaics (ipv) University of Stuttgart 70569 Stuttgart, Germany The ORCID identification number(s) for the author(s) of this article can be found under https://doi.org/10.1002/ente.202300723. © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH. This is an open access article under the terms of the Creative Commons Attribution-NonCommercial License, which permits use, distribution and reproduction in any medium, provided the original work is properly cited and is not used for commercial purposes. DOI: 10.1002/ente.202300723 For large-scale energy-storage systems, the aqueous rechargeable zinc–manganese dioxide battery (ARZMB) attracts increasing attention due to its excellent advan- tages such as high energy density, high safety, low material cost, and environ- mental friendliness. Still, the reaction mechanism and its influence on the electrolyte’s pH are under debate. Herein, a pH buffer concept for ARZMB electrolytes is introduced. Selection criteria for pH buffer substances are defined. Different buffered electrolytes based on a zinc salt (ZnSO4, Zn(CH3COO)2, Zn(CHOO)2), and pH buffer substances (acetic acid, propionic acid, formic acid, citric acid, 4-hydrobenzoic acid, potassium bisulfate, potassium dihydrogen citrate, and potassium hydrogen phthalate) are selected and compared to an unbuffered 2 M ZnSO4 reference electrolyte using titration, galvanostatic cycling with pH tracking, and cyclic voltammetry. By adding buffer substances, the pH changes can be reduced and controlled within the defined operating window, supporting the Mn2þ/MnO2 deposition/dissolution mechanism. Furthermore, the potential plateau during discharge can be increased from �1.3 V (ZnSO4) to �1.7 V (ZnSO4þ AA) versus Zn/Zn2þ and the energy retention from �30% after 268 cycles (ZnSO4) to �86% after 494 cycles (ZnSO4þ AA). Herein, this work can serve as a basis for the targeted design of long-term stable ARZMB electrolytes. RESEARCH ARTICLE www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (1 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH mailto:oliver.fitz@ise.fraunhofer.de https://doi.org/10.1002/ente.202300723 http://creativecommons.org/licenses/by-nc/4.0/ http://creativecommons.org/licenses/by-nc/4.0/ http://www.entechnol.de the LIB. Here, the stability of the cathode is influenced by para- sitic reactions such as the formation of corrosive acids (here, hydrofluoric acid (HF)) and the acidity of the electrolyte, which lead to the dissolution of transition metals (e.g., Mn in NCM811 cathodes). This phenomenon can be addressed by electrolyte additives (e.g., siloxanes,[71]), which can form an insoluble inter- phase film on the cathode and capture corrosive acids. This ana- logue can be compared to the herein-investigated aqueous ZIB system, where MnO2 can also uncontrollably be dissolved in the aqueous electrolyte in an acidic environment. For ARZMB, the main results regarding the electrolyte’s pH and the addition of (pH buffer) substances are summarized in the following in chronological order of their year of publication: Mateos et al.[54,55] introduced a feasible explanation approach for the functioning mechanism of the ARZMB based on a pH-dependent Nernst equation, highlighting the major role of the MnO2/Mn2þ dissolution/deposition mechanism and the Zn/Mn hexa–aqua complexes as the proton source. They showed the influence on the potential curve behavior of pH buffered elec- trolyte compositions (ECs) using acetic acid (AA) as the buffer solution in different electrolyte compositions, e.g., 1.5 M acetate acid with 0.25 M ZnCl2 and 0.1 M MnCl2, adjusted to pH 5 (pH adjustment with either 2 M KOH or 1 M HCl), and compared them to an unbuffered electrolyte with 0.25 M ZnCl2, 0.1 M MnCl2, and 0.85 M KCl. The potential curve showed a single sta- ble potential plateau at �1.55 V versus Zn/Zn2þ for buffered electrolytes compared to a steeper potential drop and different potential plateaus at �1.4 and 1.2 V versus Zn/Zn2þ, respec- tively, for unbuffered electrolytes during discharging, predicted by the pH-dependent Nernst equation.[54,55] Guo et al.[46] used AA to wash cycled MnO2 cathodes and remove precipitated ZHS from the electrodes after disassem- bling the cell, reassembled them in new cells and successfully recycled them in 2 M ZnSO4 electrolyte with recovered discharge capacity. As this treatment of the cathodes does not represent a pH buffering EC, it still shows the importance of controlling the ZHS precipitation and its active material passivation/isolation character by adjusting the electrolyte’s pH.[46] Li et al.[66] showed the influence of the electrolyte’s pH value on the discharge potential behavior by adjusting either the Mn2þ concentration (higher Mn2þ concentration reducing the pH) or adding sulfuric acid in the 1 M ZnSO4 and 1 M MnSO4 electrolyte.[66] Higher Mn2þ content (here, 3 M) and/or less pH value by adding sulfuric acid (here, pH �2) stabilized and increased the potential plateau (�1.8 V vs Zn/Zn2þ) as well as reduced its steepness. These observation can be related to the results of the publications of Mateos et al. and the formulated pH-dependent Nernst equation mentioned earlier.[54,55] Liu et al.[62] introduced a pH-buffered electrolyte with 1 M ZnSO4, 1 M MnSO4, and 0.2 M AA (pH �2). This electrolyte showed less pH changes during cycling, a stable potential plateau behavior at >1.8 V versus Zn/Zn2þ during discharging and high Coulombic efficiencies.[62] These results go in line with the results of the publications mentioned earlier. Molaei et al.[64] introduced tartaric acid as pH buffer substance for an ARZMB electrolyte with 2 M ZnSO4 and 1 M C4H6O6. The electrolyte’s pH was adjusted to pH 4 by using 0.1 M HCl or NaOH solution. This electrolyte showed a potential plateau at �1.46 V versus Zn/Zn2þ during discharge, good cycling stability, and a capacity in the range of 300–400mAh g�1.[64] Huang et al.[53] designed and characterized a high-concentrated dual-complex electrolyte based on ammonium acetate (CH3COONH4), zinc acetate (ZA) (Zn(CH3COO)2), ammonia (NH3), andMn(CH3COO)2, resulting in an ECwith the final molar ratio of 1 M CH3COONH4, 0.4M NH3, 0.2 M Zn(CH3COO)2, 0.075 M Mn(CH3COO)2, and 2 M H2O, showing an ionic conduc- tivity of 16.1mS cm�1 and an initial pH value of �7.9. This elec- trolyte showed a higher pH stability through its pH buffer capability compared to a reference electrolyte (REL) consisting of ZA and manganese acetate (initial pH value �5.6). The pH buffer behavior resulted in a potential plateau during discharge at �1.4 V versus Zn/Zn2þ, a good cycling stability, high Coulombic efficiency, and a high areal capacity of 10mAh cm�2.[53] Zhang et al.[65] used ammonium dihydrogen phosphate (NHP, NH4H2PO4) in an EC consisting of 1 M ZnSO4 and 0.1 M MnSO4 with 25mMNHP (initial pH 2.84). This electrolyte showed a stable pH behavior (in Zn||Zn symmetrical cells), a potential plateau at �1.4 V versus Zn/Zn2þ, and a good cycling stability.[65] Still, the previously mentioned publications dealing with pH buffer substances are not providing a prior theoretical estimation regarding the stoichiometric ratios of the buffer substances and their interaction, which is influencing the initial pH value and the pH buffer capacity following the Henderson–Hasselbalch equation for pH buffers. Altogether, we could not find a common solution approach in the literature to design the ARZMB electro- lytes with pH buffer addition, meeting all the demands for the ARZMB, such as, e.g., the initial pH of the electrolyte and its pH range during operation, the ion concentration for the major elec- trolyte components such as Zn/Mn salts and the pH buffer sub- stance, as well as their ratio, their electrochemical stability window, and the ionic conductivity of the electrolyte. Therefore, this work aims at defining the demands for an ARZMB electrolyte with pH buffer properties, based on our previously published studies.[42,70] Based on these demands, selection criteria for finding appropriate pH buffer substances are formulated and respective substances are selected. Furthermore, the pH buffer capacity of these substances is analyzed by the titration technique with acid and basic titrants. Additionally, theoretical calculations by applying the Henderson–Hasselbalch equation are performed, under consid- eration and discussion of the underlying assumptions for this equation. Subsequently, an RE consisting of 2 M ZnSO4 is tested Table 1. Overview over pH additives for ARZMB electrolytes found in the literature so far. Chemical substance Chemical formula Source Ammonium dihydrogen phosphate Phosphoric acid NH4 þ H2PO4 �H3PO4 [65] Ammonium acetate NH4 þ CH3COO� [53] Zinc acetate Zn2þ (CH3COO�)2 Ammonia NH3 Acetic acid CH3COOH [46,54,55,62] Sulfuric acid H2SO4 [66] Tartaric acid C4H6O6 or C2H2(OH)2(COOH)2 [64] www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (2 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de in cyclic voltammetry (CV) and (long-term) cycling experiments and compared to different EC with pH buffer additives. Finally, recommendations for designing an ARZMB electrolyte with pH buffer capabilities are summarized. This work can be seen as an important contribution on the way to the solution of the ARZMB electrolyte pH issue to enable long- term stable ARZMB systems for stationary storage applications. 2. Results and Discussion Reviewing the literature for ARZMB, the standard electrolyte can be defined as an aqueous 2 M ZnSO4 solution as the electrolyte basis.[2,42,43] Therefore, this EC will be set as the reference system for the following experiments to, first, clarify the pH challenge within ARZMB electrolytes by showing the acidic and alkaline pH behavior of this aqueous electrolyte solution containingmetal ions, and, second, compare modified EC regarding their pH and cycling behavior. 2.1. Reference System To demonstrate the pH behavior of the (unbuffered) RE 2 M ZnSO4, this RE was examined in response to a pH change (see Figure 1). The increasing OH� concentration in the cell dur- ing cycling was imitated by adding NaOH solution (0.1/1/10 M NaOH, see Experimental Section) and the formation of Hþ by adding 0.1 M HCl in a titration experiment. This experiment con- firms the results of a previous publication by Lee et al.,[45] which is one of the first publications in the field of ARZMB to highlight the importance of the pH value during operation of the cell.[45] To obtain a general understanding of the bufferingmechanisms of the electrolyte, the pH curve in Figure 1 can initially be inter- preted qualitatively (a quantitative interpretation will follow in Section 2.2.2): titrating small amounts of OH� ions (NaOH solu- tion), the pH significantly increases. Rapidly, the initially transpar- ent solution turns milky and opaque, indicating the precipitation of a ZH species (here, ZHS, see reaction Equation (1)).[42,45] 6OH� þ SO4 2� þ 4Zn2þ þ nH2O ⇌ Zn4ðOHÞ6SO4 ⋅ nH2O with n ¼ ½4, 5� (1) This electrically isolating species buffers the change in pH in the electrolyte, but can lead to the clogging of the porous cathodes and to Zn2þ loss within the electrolyte, increasing the inner resis- tance of the cell, and finally lead to the degradation of the ARZMB cell.[44,72] Therefore, the control of the pH value in the ARZMB electrolyte is a major task for the long-term stability of the ARZMB and will be addressed in the further scope of this work. 2.2. Identification of Buffer Substances 2.2.1. Selection Criteria Due to the large choice of potential buffer substances, a definition of selection criteria is required. Finally, the pH buffer substances in question can be selected within the scope of this work based on the defined selection criteria using a funnel approach. Based on patent and literature research on buffer materials, some zinc salts and additives were first collected, which are shown in Figure S2, Supporting Information. Substances or systems that are suitable as pH buffers can usually be described as weak Brønsted acids and their corresponding bases. Due to the—comparatively weak—acid/base character (shown in higher pKa values in the range of, e.g., pKa> 0 or rather >3), the substances can be protonated or deprotonated, depending on the pH value, and thus compensate for a pH change in the aqueous medium. As can be seen from the Henderson–Hasselbalch equation (see Equation (S9), Supporting Information), the initial pH value of the EC and the pH buffering behavior can be adjusted by the ratio of the weak acid and a weak base. Due to the specific cell chemistry, there are several requirements for the electrolyte of the ARZMB. Therefore, the following selection criteria were taken into account for a meaningful selection of materials for the following experiments. Solubility in Water: A zinc salt should be present as the stan- dard component of the electrolyte. The addition of a manganese salt is often shown in the literature to improve performance, but it is deliberately avoided for this work. The renunciation is based on the guarantee of an electrolyte system that is as binary as possible consisting of preferably two substances to come as close as possible to the mathematical descriptions of buffer solutions according to the Henderson–Hasselbalch equation (see Section 3.4 for further details). In addition, a sufficient ionic con- ductivity (>10mS cm�1) of the electrolyte is required to reduce Ohmic losses due to resistance in the electrolyte. Supporting electrolytes (e.g., KCl) to increase the ionic conductivity should be avoided, as this reduces the transport of the active ZnSO4 -0.5 0.0 0.5 1.0 1.5 2.0 0 1 2 3 4 5 6 7 pH c(OH-) / mol.l-1 pH 5.5 c(H+) / mol.l-1 -0.01 0.00 0.01 3 4 5 6 ~0.003 mol.l-1 A B C Figure 1. A) Titration curve of the RE 2 M ZnSO4 titrating it with 0.1 M NaOH and 0.1 M HCl, respectively, as well as the optical comparison of B) the initial look of the solution before and C) the look after the addition of NaOH. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (3 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de components through migration and further increases the ionic strength. Furthermore, the concentration of the Zn2þ ions influences the energy density of the cell.[66] Consequently, the solubility of the zinc salt and—at the same time—the concentration of available Zn2þ ions must be sufficiently high. To ensure a buff- ering effect of the electrolyte, the existence of a weak acid and base is necessary and a sufficient solubility of the buffer substan- ces in water is also necessary. Therefore, the solubility is set at >0.3 mol L�1 within the scope of this work. Electrochemical Stability: In principle, the electrochemical voltage series can be used to get a first impression of the elec- trochemical stability of the respective substance based on the possible chemical reactions in an aqueous solution. For this pur- pose, the considered potential range must be defined: consider- ing the standard electrode potentials of the relevant chemical reactions of Zn/Zn2þ, MnO2/Mn2þ and H2/H þ with respect to the respective reference point such as the standard hydrogen electrode (SHE) at a pH value of 0 (s. standard conditions in the electrochemical series) or the Zn/Zn2þ half-cell reaction at pH 0 and pH 4, respectively, potential shifts in the chemical reactions’ result, which are visualized in Figure 2.[49,73] Since zinc is the potential reference in the electrochemical system investigated here, the electrochemical potential versus SHE is transformed into the potentials versus Zn/Zn2þ (s. Figure 2B). To enable a dissolution/deposition reaction of MnO2/Mn2þ in the system, it is necessary to ensure a sufficient electrochemical potential for this reaction. In a slightly acidic electrolyte at pH� 4, the potential for MnO2/Mn2þ dissolution/deposition is �1.518 V versus Zn/Zn2þ (s. Figure 2C). Therefore, the end-of-charge potential must be at least 1.518 V versus Zn/Zn2þ or more. Based on the literature, end-of-charge voltages between 1.7 and 1.9 V ver- sus Zn/Zn2þ and end-of-discharge voltages between 0.7 and 1 V versus Zn/Zn2þ are given. Within the scope of this work, the battery cells are to be cycled in the potential range of 0.7 to 1.8 V versus Zn/Zn2þ. To ensure electrochemical stability even at increased voltage ranges (conceivable, e.g., in the context of regeneration cycles or maintenance tasks), a potential range of 0.5–2.1 V versus Zn/Zn2þ for the EC should be used as a selec- tion criterion.[4,67,74–77] The electrochemical stability of the zinc salt and the other electrolyte constituents in the mentioned potential range versus Zn/Zn2þ is relevant for the service life of the battery, so that the substances are not involved in undesirable side reactions. Therefore, for a selection of the buffer substances in this work, their electrochemical stability in the potential range investigated here was evaluated on the basis of the electrochemical series.[73] Dissociation Constant: As many reactions in the ARZMB influ- ence the pH value of the electrolyte, certain pH thresholds must be observed for cycling: 1) The potential pH diagram for zinc (s. Figure S8, Supporting Information) indicates the beginning of a precipitation reaction of a ZH species by means of a vertical line at a pH value of �5.5 or more[49,70]; 2) The adapted potential pH diagram for zinc, taking into account the overvoltages for the hydrogen evolution reaction (HER), shows increased corrosion of zinc with the formation of hydrogen from a pH below �3[78]; 3) The potential-pH diagram for manganese (s. Figure S9, Supporting Information) shows the potential overlapping of the MnO2/Mn2þ dissolution/deposition and the oxygen evolution reac- tion (OER) due to the proximity of the two equilibrium lines to each other. The lower the pH value, the higher the necessary potential for triggering the OER and the potential gap between the OER and the MnO2 deposition [49,70]; and 4) As the potential-pH diagram for manganese shows, the electrodeposition of MnO2 is highly depen- dent on the potential and pH conditions within the aqueous solu- tion. By this, and especially by the pH conditions, the electrodepositedmanganese oxide species can vary, and the dispro- portionation of Mn2þ to Mn2þ and Mn4þ as well as the formation of MnOOH can take place (latter at higher pH values).[59,79,80] The initial pH value of the electrolyte solution should be set within these limits to reduce the OER on the one hand and to avoid the HER with zinc corrosion on the other. Therefore, an initial pH value of �4� 1 is sought as a compromise to provide a sufficient pH buffer range. It should be noted that during the initial material investiga- tions of the selected buffer systems (BSs), the concentration of weak acid is initially kept constant at 0.1 M, which is based on the specifications for standard pH buffer solutions.[81] By taking into account the dissolved zinc and manganese salts, the BS can be defined based on the Henderson–Hasselbalch equation (s. Equation (S9), Supporting Information).[73] Safety and Toxicity: The ARZMB aims at providing an environ- mentally friendly and harmless storage technology for SES. For this reason, attention should be paid to an environmentally and health-friendly choice of materials during the ARZMB develop- ment. Within the framework of these selection criteria, the choice of material is made in compliance with the hazardous sub- stance labeling, whereby highly toxic and oxidizing substances such as sulfides and peroxides or cancerogen mutagen reprotoxic (CMR) substances are to be excluded. Based on these specifications, different electrolyte systems with three concepts for the formation of an EC with pH buffer properties are formulated within the scope of this work: an addi- tive system (AS), a BS, and a hybrid system (HS) (Figure 3). The AS represents an aqueous system consisting of the standard zinc salt 2 M ZnSO4 with the addition of a single weak Brønsted acid with the molar concentration of 0.1 M. Therefore, this EC only shows one difference compared to the RE with the aim of buffering the pH and differs mainly in the pKa values of the added buffer acids. The BS is also a binary system but consisting of a zinc salt with its conjugate acid. The conceptual design of the BS corresponds Zn/Zn2+ -0.762 0.762 1.986 E 0/ V vs. Zn/Zn2+ pH 0 E 0/ V vs. Zn/Zn2+ pH 4 0 φ1 = 0.7 φ2 = 1.8 SHE MnO2/Mn2+ 0 1.224 E 0/ V vs. SHE pH 0 0 1.518 C B A Figure 2. A) Standard potentials of the electrochemical series versus SHE (pH 0) and compared to the visualization versus Zn/Zn2þ B) at pH 0 and C) at pH 4.[55,73,99] www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (4 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de to the characteristic of a pH buffer described in Section S1, Supporting Information. Instead of ZnSO4, an alternative zinc salt is intended to be used. The advantage of the BS is the possibility of varying the concentration of its buffer base and buffer acid components, which in principle allows the setting of an initial pH value of the electrolyte solution based on the Henderson–Hasselbalch equation. The main challenge in the formulation of the EC of the BS for ARZMB is the sufficient sol- ubility of the zinc salt. The HS is a ternary system consisting of the established zinc salt ZnSO4 with the addition of acids with different pKa values. This allows the targeted adjustment of an initial pH value based on the dissociation constant of the acids, regardless of the con- centration of the added amount. All substances which meet the selection criteria in Table 2 are further classified. The classification was performed on the basis of the systems presented in Table 3. Detailed properties of the selected acids are listed in Table S1–S5, Supporting Information. 2.2.2. Titration of Buffer Electrolyte Compositions In the previous chapter, the EC were categorized into three sys- tems. Within this section, the different EC are investigated and compared with the RE 2 M ZnSO4 in titration experiments to gain insights in the buffer behavior.[54,55,57] The results are shown in Figure 4 by plotting the pH value over the standardized hydroxide concentration, which was added as NaOH solution to simulate the pH increase by MnO2 disso- lution when discharging the ARZMB. First, evaluating the titration curves in Figure 4A, the initial pH value of the different EC significantly differs due to the vary- ing pKa values of the additives and their constant molar concen- trations of 0.1 M, respectively, which influences the buffer capacity. All EC show a pH buffer behavior compared to the RE 2 M ZnSO4, which can be referred to the significantly delayed pH increase. Also, the EC in Figure 4A shows a similar buffer behavior compared to the titration curves of the different addi- tives without ZnSO4 addition to the electrolyte (s. Figure S2, Additive System (AS) ZnSO4 + weak Brønsted acid Hybrid System (HS) ZnSO4 + weaker & stronger Brønsted acid Buffer System (BS) zinc salt + conjugated Brønsted acid electrolyte � � � � � � specifications solubility >0.3 mol.l-1 ionic conductivity >10 mS.cm-1 potential window 0.5-2.1 V vs. Zn/Zn2+ initial pH value pH 4.5 ± 1 dissociation constant pKa 4-6 safety & toxicity: no CMR/peroxides/sulfides Figure 3. Specifications at the electrolyte of the aqueous rechargeable zinc–manganese dioxide battery (ARZMB) and formulation of the systems for the classification of the different electrolyte composition (EC). Table 2. Summary of electrolyte specifications for the pH buffer substances within this work. Parameter Criterion Solubility >0.3 mol L�1 Ionic conductivity >10mS cm�1 Electrochemical stability in the potential window 0.5–2.1 V versus Zn/Zn2þ Initial pH value pH 4.5� 1 Dissociation constant pKa� 4–6 Safety & toxicity No CMR substances, peroxides, sulfides Table 3. Overview of development cooperation divided into the additive system (AS), the buffer system (BS), and the hybrid system (HS). Additive system (AS) Buffer system (BS) Hybrid system (HS) 2 M zinc sulfateþ 0.1 M Brønsted acid pH 3–4 2 M zinc sulfateþ 0, 1 M Brønsted acids in combination Inorganic acid Organic acid Zinc salt Conjugated acid Stronger acid Weaker acid Potassium bisulfate (KHS) 4-Hydrobenzoic acid (HBA) Zinc acetate (ZA, 0.8 M) Acetic acid (AA, 4.56 M) Formic acid (FA) Acetic acid (AA) Potassium dihydrogen citrate (KDHC) Propionic acid (PPA) Zinc formate (ZF, 0.3 M) Formic acid (FA, 1.69 M) Potassium hydrogen phthalate (KHP) Acetic acid (AA) Citric acid (CA) Formic acid (FA) www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (5 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de Supporting Information). Furthermore, it is noticeable that the initial pH value of all EC (in addition to potassium bisulfate [KHS]) is significantly lower when ZnSO4 is included in the solu- tion. This behavior can be referred to the formation of zinc hexa– aqua complexes [Zn(H2O)6] 2þ in an aqueous solution, which have an acidic character (pKa= 9).[55] Additionally, in addition to potassium dihydrogen citrate (KDHC) and citric acid (CA), the titration curves do not exceed a pH of �5.5, where the pre- cipitation of ZHS begins (s. Figure 1). Considering this, since both KDHC and CA contain the citrate group C6H5O7, a delayed precipitation of ZHS could be attributed to the complex forma- tion of zinc–citrate complexes in aqueous solution. Nevertheless, in consequence of the multiple dissociation stages in reference to their multiple pKa values, CA and KDHC have the highest buffer capacities. The BS provides the possibility to set an initial pH value by using a combination of a base and its conjugate acid using the Henderson–Hasselbalch equation. As for the BS a relatively high acid concentration was added to adjust the initial pH value (s. Table 3), the buffer capacity of the BS is significantly higher than in the AS or HS and the titration results cannot be directly compared to the AS or HS. Furthermore, the BS does not contain any sulfate ions, so that in consequence the precipitation of ZHS cannot take place. Though, another ZH species precipitation can occur which is dependent on the concentration of metal ions in the solution according to reaction (S1, Supporting Information). In Figure 4C, where the titration curves of the BS with the pH over the added hydroxide concentration are displayed, this can be seen at by a shift of the tipping point to �pH 7. The precipitation of ZH during the titration of ZAþ acetic acid (ZAþ AA) started KHP KDHC KHS AA PPA CA HBA FA RE: ZnSO4 0.00 0.02 0.04 0.06 0.08 0.10 0 1 2 3 4 5 6 7 8 pH c(OH-) / mol.l-1 pH 5.5 3.5 4.0 4.5 5.0 5.5 0.00 0.01 0.02 0.03 0.04 0.05 l.lo m/ -1 pH KHP KDHC KHS AA PPA CA HBA FA RE: ZnSO4 ZA + AA ZF + FA 0.0 0.5 1.0 1.5 2.0 2.5 3.0 0 1 2 3 4 5 6 7 8 pH c(OH-) / mol.l-1 pH 5.5 3.5 4.0 4.5 5.0 5.5 0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 l.lo m/ -1 pH ZA+AA ZF+FA RE: ZnSO4 AA + FA comparison: AA comparison: FA 0.00 0.02 0.04 0.06 0.08 0 1 2 3 4 5 6 7 8 pH c(OH-) / mol.l-1 pH 5.5 3.5 4.0 4.5 5.0 5.5 0.000 0.005 0.010 0.015 0.020 l.lo m/ -1 pH AA+FA RE: ZnSO4 A B C D E F Figure 4. Comparison of the titration curves of A) the additive system (AS), C) the buffer system (BS), and E) the hybrid system (HS), for comparison, the results of FA and AA from the AS are added by the titration with NaOH solution. Furthermore, the pH buffer capacity curves are added for B) the AS, D) BS, and F) HS, respectively, based on the analysis of the titration curves by their fitting and derivation, for a detailed interpretation of the pH buffer behavior within the defined pH working window for each EC. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (6 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de at �pH 6 in the form of small flakes in the electrolyte (s. Figure 5). Interestingly, ZH exceptionally dissolved again after some time, and only after pH �7 was reached, strong turbidity could be recognized indicating the final tipping point of the ZH precipitation. The delayed precipitation of ZH can also be attrib- uted to the complex formation of zinc–acetate complexes in aque- ous solution. Looking at the titration curve of zinc formateþ formic acid (ZFþ FA) (Figure 4E), a comparable behavior compared to ZAþ AA (s. Figure 4A) can be noticed, but with a reduced pH buffer behavior and a slightly lower initial pH value, which can be referred to the lower overall molar concentration of ZF and FA, as well as to the lower pKa value of FA. Within the HS, by the combination of different buffer acids, it is possible to influence the buffer capacity and the initial pH value within the EC. In Figure 4E, the EC with ZnSO4þ AAþ FA as well as the single components þFA and þAA from the AS (s. Figure 4A) for comparison reasons are visu- alized. This EC combines two Brønsted acids with different pKa values (here, AA with pKa 4.76 and FA with pKa 3.75) with a molar concentration of 0.1 M each. Comparing the HS to the sin- gle components from the AS leads to the following interpreta- tions. 1) Looking at the buffer capacities, the HS containing both AA and FA with a total acid concentration of 0.2 M shows a lower buffer capacity than the sum of the single components AA and FA in 0.1 M each (s. Table S7, Supporting Information). Therefore, even though the buffer capacity is dependent on the acid concentration, it seems not linear proportional; 2) Investigating the titration curves characteristic, the curve of the HS shows a comparable behavior than that of the single com- ponents of the AS with ZnSO4þ AA and þFA, respectively. As for the HS, the combination of two acids with different pKa val- ues in fact results in a pH behavior with a multistage dissociation behavior for each of the pKa value. As the pKa values of FA and AA are close to each other, the dissociation stages should be over- lapping, resulting in an extended buffer behavior. A quantification of the buffer capacities is given in Table S7, Supporting Information. The calculation method is shown in Section S4, Supporting Information, based on Equation (S16), Supporting Information. Here, the linear range (criterion: R2> 0.99) of the pH buffer range and its OH� concentration change is determined. Regarding the results for the pH buffer capacity (either normalized to one pH step, or for the total linear range of the titration curve), within the AS with 0.1 M acid con- centration each, the EC ZnSO4þ AA shows the highest pH buffer capacity. The calculation method can be used for deter- mining the pH buffer capacity regarding the normalized buffer capacity for one pH step, as well as for the linear range of the titration, which follows the literature for the determination of the pH buffer capacity (e.g., ref. [73]). Still, this calculation method does not give all the necessary information of the pH buffer behavior (e.g., transgression of pH limits such as ZHS precipitation/HER, total pH buffer capac- ity within pH operation window, etc.) for a specific aqueous sys- tem such as the ARZMB. Here, for a quantitative analysis of the buffer capacity behavior within the selected pH operation win- dow of pH 3.5–5.5, the titration curves in Figure 4A,C,D were fitted with a polynom of 6th grade (s. Figure S4, S5, and S7, Supporting Information, note: x- and y-axis swapped before fit- ting). Subsequently, the fitting result was differentiated to obtain the behavior of the buffer capacity within the selected pH opera- tion window. The results are shown in Figure 4B,D,F for the AS, BS, and HS, respectively. By this method, the pH buffer behavior in regard of its EC can be analyzed and thereby modified to receive the pH buffer behavior which is needed for the aqueous system (here, ARZMB). The following conclusions can be drawn by interpreting the pH buffer capacity curves. 1) The RE ZnSO4 shows in fact a pH buffer capacity within the selected pH opera- tion window, but on a very slight extent. The pH buffer behavior can mainly be referred to the weak acid strength of the Zn2þ ions and their hexa–aqua complex (with a molar concentration of 2 M and low pKa value of �9, s. provided earlier), enabling a (de-)protonation to a certain degree; 2) The AS generally show low buffer capacities compared to the RE, but only regarding the selected pH operation window of pH 3.5–5.5. In contrast, for the pH buffer range with a linear pH increase of each buffer acid of the AS, the buffer capacity generally significantly exceeds the buffer capacity of the RE (s. Table S7, Supporting Information). This shows the significance of the different pH buffer calculation methods in accordance to the linear range of the titration, or to the defined pH operation window; 3) KDHC (s. Figure 4A) shows high pH buffer capacities within the defined pH operation window for pH values between 3.5 and 4, and again for pH 5.0–5.5, which can be visually analyzed with this introduced evaluation method. This EC property can be used, for example, to specifically shift the maximum pH buffer capacity to these pH regions (e.g., to maximize the pH buffer capacity for the avoidance of the ZHS precipitation at pH �5.5, or to reduce HER at lower pH values �3.5); and 4) The BS of ZAþ AA contains 0.8 M ZA and 4.56 M AA, leading to an increased molar concentration and ionic strength of a buffer base and acid in total, and therefore shows an improved pH buffer characteristic. This approach is proved by the significantly higher pH buffer capacity compared to the AS and HS (with lower acid concentrations) and the pH buffer capacity max- imum at pH� 4,4, which is in the range of the pKa value for AA of pKa= 4,76. The deviation can be explained by the overall ionic strength of the particular EC of the BS. The application-oriented approach by defining the pH opera- tion window of the EC for the aqueous system and analyzing the Figure 5. Zinc hydroxide (ZH) precipitation during the titration of 0.1 M NaOH into the EC of the BS. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (7 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de pH buffer behavior within this operation window (s. Figure 4B, D,F) further helps to design an EC with targeted pH properties. 2.2.3. Cycling of Buffer Electrolyte Compositions The investigated ECs of the titration experiments are further examined in cycling experiments. The experimental setup is based on our previous study (s. ref. [42]) with the combination of full cell cycling and in-operando pH tracking with a pH sensor (for details, s. Experimental Section). It is noticeable that the initial pH values of the buffered electrolytes (pH 1.5–2.9) are significantly below the initial pH value of the RE 2 M ZnSO4 (pH 3.35). Regarding the potential-pH diagram of zinc (s. Figure S8, Supporting Information), the Zn corrosion in par- allel to the HER is amplified with lower pH values. To prevent this phenomenon and its effect on cycling and pH behavior, an uniform initial pH value of the EC would be necessary. However, this can only be achieved by varying the concentration of the buffer substance or adding an additional substance to adjust the pH, which in turn would change the EC and the molar con- centration of the buffer. Due to this trade-off, the uniform buffer acid concentration of 0.1 M is used as a basis for the following cycling experiments for the sake of comparability, and the cycling and pH behavior is evaluated taking into account the influences of a varying initial pH value. AS: For the AS in Figure 6A1–H1, the potential curves versus Zn/Zn2þ and in parallel the pH curves are displayed over time. In addition, in Figure 6A2–H2, the areal discharge capacity per cycle (QDC) for both the RE and the investigated EC is visualized. The results are compared to the RE ZnSO4 in Figure 6Aþ B. All the pH curves show a pH buffer behavior in terms of the delay of the pH increase within the first few cycles, compared to the RE. Still, differences in terms of the intensity of the pH increase during cycling are observable regarding the buffer sub- stance: For example, KHS (s. Figure 6C1) and FA (s. Figure 6H1) show a stronger pH increase in comparison to the other investi- gated EC. This is congruent with the titration experiments (s. Figure 4) as KHS and FA show the least system relevant buffer capacities, which can be related to the lower pKa value of the respective acid and therefore a lower pH buffer capacity in the pH range under consideration. Also, it is conspicuous that the pH curves quickly approach the pH value of �5.2 when the electrolyte is at rest after the initial 10 cycles. Considering local pH phenomena within the pores of the porous positive carbon fiber electrode by MnO2 dissolution, during the rest phase, the local phenomena can be measured globally and furthermore, the pH level at pH�5.2 can be attributed to the pH level of the ZHS precipitation. Observing the pH curves of AA and propionic acid (PPA), but also HBA (note: HBA was not completely soluble in the concentration of 0.1 M), they show the least trend over the 10 cycles compared among all EC. Regarding the cyclability of the modified EC with buffer acids, it needs to be mentioned that both KDHC and CA could not be cycled properly. After the first few cycles, the capacity already reached zero and no more cycling was possible. Looking at the electrodes after cycling (s. Figure S10, Supporting Information), the cathodes of CA (s. Figure S10A, Supporting Information) and KDHC (s. Figure S10B, Supporting Information) are fully discharged so that no MnO2 is visible on the carbon fiber substrate compared to the pristine state after the initial MnO2 electrodeposition. This is referred to the citrate group which could form stable Mn complexes (high stability constant of complex for [MnHCit]�: lg K= 3.54–3.67)[82,83] and impede proper Mn2þ/MnO2 deposition/dissolution or even the loss of active Mn2þ. Regarding the cycle stability compared to the unbuffered RE, the EC of potassium hydrogen phthalate (KHP), KHS, AA, and PPA shows a very similar capacity curve, whereas FA shows significantly lower capacities. For the latter, this behavior can be attributed to the deposition/dissolution behavior of the Mn2þ/MnO2 loading of the electrolyte/electrode, which is assumed to be the essential reaction mechanism regarding the latest literature for ARZMB with acid electrolytes.[29,43–53] The different characteristics of the capacity curves should be examined within long-term stability tests (LSTs), which is per- formed within Section 3.3. BS: In Figure 7, the cycling results of the BS electrolytes for A1, A2) ZFþ FA and B1, B2) ZAþ AA in comparison to A,B) the RE are shown. Here, zinc salts with their conjugated acid were considered. For the EC-containing ZFþ FA (s. Figure 7A1,B1), after the first cycle, the capacity drops to nearly zero and no further cycling can be performed indicating an inactive cell. This behavior can be related to the decomposition of the zinc anode during the first cycle as a result of zinc corrosion/HER (s. Figure S11A, Supporting Information), which was already discussed for the AS with ZnSO4þ FA. The EC-containing ZAþ AA (s. Figure 7A2,B2), although the capacity is significantly lower compared to the RE, shows a higher capacity stability with even a slight increase in capacity. Furthermore, the pH curve after the 10 cycles stays constant when the cell is at rest. The reduced capacity of the cell can partly be attrib- uted to the reduced ionic conductivity of the EC (s. Table 7) in respect to the RE. An important observation is the formation of salt precipitations around the electrical contacts of the cell cuvette, which only occurs for the EC with higher AA concentration and therefore can be attributed to the AA content of the electrolyte and its evaporation (s. Figure S11B, Supporting Information). This phenomenon could be inhibited by using closed cell formats. HS: The cycling results for the HS are shown in Figure 8. For these EC, the pH buffering effect and initial pH value can be adjusted by combining two different buffer acids with different pKa values each. For the combination of AAþ FA with ZnSO4, after a slight initial pH increase, a significant pH buffer behavior is shown with a periodic pH oscillating around pH �3.25 (s. Figure 8A), which is of higher degree compared to ZnSO4þ FA. Still, the buffer and cycling behavior both has a comparable characteristic with the EC ZnSO4þ FA of the AS. This again leads to the conclusion that within the HS, the stron- ger acid with the lower pKa value dominates the pH buffer behav- ior of the EC. The cycling of the HS confirms the possibility of combining two buffer acids with different chemical properties, whereby tar- geted buffering effects of the EC can be adjusted. This principle can be transferred to other possible EC and can enable an elec- trolyte design specifically according to the requirements of the battery cell chemistry. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (8 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de E pH ZnSO4 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH ZnSO4 0 5 10 0 1 2 3 q D CH mc. hA m/ -2 cycle number E pH ZnSO4 ZnSO4+KHP 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH ZnSO4 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number E pH ZnSO4 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number A A1 A2 C2C1 D1 D2 B2B1 B E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E pH ZnSO4 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number ZnSO4 + KHP ZnSO4 ZnSO4 + KDHC ZnSO4 ZnSO4 + KHS ZnSO4 ZnSO4 + AA ZnSO4 + KHS ZnSO4+AA ZnSO4+KDHC Figure 6. A1–H1) Cycling results of the different EC of the AS showing the potential curve together with the pH behavior, as well as A2–H2) the cycle stability curve for the initial ten cycles, A,B) both compared to the RE 2 M ZnSO4 without pH buffer capability. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (9 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de In general, after analyzing the experiments, it must be noted that the pH value can only be measured directly in front of the cathode, but not inside the porous electrode. Therefore, the pH results only indirectly represent the local pH behavior at the exact location of the electrochemical reaction (electrode–electrolyte interface). Nevertheless, the reversibility of the pH fluctuations as well as the results of the cells shows that a transferability of the pH fluctuations within the porous electrode is appropriate. Summary: The potential curves together with the pH behavior of the electrolyte generally show a pH buffering behavior within the first 10 cycles compared to the RE 2 M ZnSO4. The buffering behavior is indicated by a delayed increase of the pH value in the electrolyte. In principle, the general pH increase of the EC can be attributed to irreversible gassing reactions (e.g., zinc corrosion with HER). For the AS, in regard to the limited buffer capacity of the buffered EC due to the addition of only 0.1 M buffer acid in the context of this comparative study, the pH buffer can be con- sumed quickly. Overall, the discharge capacities for all EC are observed to be either at or below those of the RE ZnSO4. Among the properly cycled electrolytes, FA in particular shows a significantly lower capacity than ZnSO4. Various factors must be considered: due to the addition of the buffer substances, the initial pH value in the EC generally drops below the initial pH value of the RE ZnSO4, which influences zinc corrosion and MnO2 dissolution as well as the gassing behavior during cycling. Looking at the E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number F1 E2E1 F2 G2G1 E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number H2H1 E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ ZnSO4 + PPA ZnSO4 ZnSO4 + PPA ZnSO4 ZnSO4 + CA ZnSO4 + CA ZnSO4 ZnSO4 + HBA ZnSO4 ZnSO4 + FA ZnSO4 + FA ZnSO4 + HBA Figure 6. Continued. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (10 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc .hA m/ -2 cycle number E pH ZF+FA 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH ZF + FA 0 5 10 0 1 2 3 q D CH mc .hA m/ -2 cycle number E pH ZA+AA 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH ZA + AA 0 5 10 0 1 2 3 q D CH mc .hA m/ -2 cycle number A B B2B1 A2A1 ZnSO4 ZnSO4 ZnSO4 ZnSO4 E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ Figure 7. A1,B1) Cycling results of the different EC of the BS showing the potential curve together with the pH behavior, as well as A2,B2) the cycle stability curve for the initial 10 cycles, A,B) both compared to the RE 2 M ZnSO4 without pH buffer capability. E pH 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number E pH ZnSO4 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 t / h 2 3 4 5 6 7 pH 0 5 10 0 1 2 3 q D CH mc.hA m/ -2 cycle number DC BA E / V v s Zn /Z n2+ E / V v s Zn /Z n2+ ZnSO4 ZnSO4 ZnSO4 ZnSO4 + FA + AA ZnSO4+FA+AA Figure 8. C) Potential and pH curve over time, and D) the areal-specific discharge capacity per cycle for A,B) the cycling experiment within the HS in comparison to the RE. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (11 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de ionic conductivities (s. Table 7), all EC with 2 M ZnSO4 are basi- cally in a similar range of �55mS cm�1 —with the exception of KHS, which is significantly higher at 62.4 mS cm�1. This may be due to the additional Kþ ions and/or the additional sulfate anions SO4 2�. 2.3. LSTs To validate the previous results of the buffered EC, the effects on the long-term cycle stability, of stirring as well as of an increased acid concentration, are investigated in comparison to the RE ZnSO4. Hereto, the pH behavior and the potential plateaus are examined. Based on the results of the previous cycling experiments of the different EC in Section 3.2.3, the best-performing EC were selected for the following LSTs and compared to the RE, in particular ZnSO4þAA, ZnSO4þ PPA, and ZnSO4þ FAþ AA. In advance to the LST, the redox behavior of the selected EC was examined performing CV experiments. 2.3.1. CV for Selected Electrolytes For a better understanding of the electrochemical behavior of the EC used in the LST, they are further characterized in CV experi- ments. Thereby, the redox behavior of the different EC is inves- tigated regarding different acid concentration and electrolyte systems. The results of the CV experiments are displayed in Figure 9, whereby on the ordinate, the current is related to the area of active material on the electrode (Aelectrode= 3,75 cm2). In Figure 9A, the RE shows a typical reduction peak for the first cycle at �1.25 V versus Zn/Zn2þ and an oxidation peak at �1.6 V and 1.8 V versus Zn/Zn2þ, which refers to the MnO2 oxidation and OER, respectively.[42] For the second cycle, the peak characteristic stays the same, but with slight changes in the peak appearance (e.g., widening of the peaks) which can be referred to changes in the ionic composition of the electrolyte through the Zn2þ and Mn2þ deposition/dissolution reactions. The other EC based on ZnSO4 with the addition of acids such as AA (s. Figure 9B), PPA (s. Figure 9C), and AAþ FA (s. Figure 9D) show comparable peak characteristics for the first cycle with a distinct reduction and oxidation peak for the Mn2þ/MnO2 deposition/dissolution, and a (smaller) second peak, which can be referred to the formation of MnOOH (s. Section 3.2.1). Still, for the second cycle, the peak character- istics significantly change especially for the ZnSO4þAA and ZnSO4þ AAþ FA electrolyte. For þAA, the peak characteristic changes to very low current flows. This behavior cannot be deeply explained within this study and needs to be investigated in fur- ther studies. A comment on this phenomenon is added to Supporting Information (s. S9). Nevertheless, as the following section shows, this electrolyte still shows a cyclability in full cells. For þAAþ FA, the formation of new reduction peaks can be referred to the zinc corrosion behavior, which was already discov- ered for the pure ZFþ FA electrolyte (s. Figure S11A, Supporting Information). Altogether, the CV shows the redox behavior of the electrolytes and a comparable peak characteristic for the different EC com- pared to the RE ZnSO4, proving their cyclability. Still, the changes in the peak characteristics for the second cycle need to be related to the cycling behavior within the following section. 2.3.2. Long-Term Cycling To ensure a homogeneous distribution of the different compo- nents of the EC within the electrolyte–electrode interface as well 0.8 1.0 1.2 1.4 1.6 1.8 -3 -2 -1 0 1 A B C D 1st cycle 2nd cycle mc.A m/I -2 0.8 1.0 1.2 1.4 1.6 1.8 -5 -4 -3 -2 -1 0 1 mc.A m/I -2 1st cycle 2nd cycle 0.8 1.0 1.2 1.4 1.6 1.8 -3 -2 -1 0 1 mc.A m/I -2 1st cycle 2nd cycle 0.8 1.0 1.2 1.4 1.6 1.8 -2 -1 0 1 mc.A m/I -2 1st cycle 2nd cycle ZnSO4 ZnSO4 + AA ZnSO4+AA+FA ZnSO4+PPA E / V vs Zn/Zn2+ E / V vs Zn/Zn2+ E / V vs Zn/Zn2+ E / V vs Zn/Zn2+ Figure 9. Graphs of the cyclic voltammetry of the EC A) ZnSO4 (2 M), B) ZnSO4þ AA (2 Mþ 1M), C) ZnSO4þ PPA (2 Mþ0.5 M), and D) ZnSO4þ AAþ FA (2 Mþ 0.1 Mþ 0.1 M) for the first two cycles each. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (12 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de as inside the porous cathode structure, the effect of stirred elec- trolytes was tested for the LST. This was based on the assumed major reaction mechanism of the Mn2þ/MnO2 deposition/ dissolution, which leads to the importance of a stable pH value in front of and inside the porous cathode structure. Furthermore, the acid concentration within the previous tests, which was set to a fixed value of 0.1 M for a comparative pH buffer study, was modified and increased for AA (from 0.1 to 1 M) and PPA (from 0.1 to 0.5 M) to evaluate the influence of a higher pH buffer concentration on the overall pH buffer capacity and long-term stability. In Figure 10, the stirred electrolytes show a comparable characteristic of periodic pH changes during cycling, as it was observed for the EC in the previous EC buffer cycling tests. Interestingly, looking at the stirred RE ZnSO4 compared to the cycling experiment without stirring, it becomes visible that the pH oscillations are more pronounced and the pH level for them is slightly lower, even though the critical pH value of �5.5 is still reached at the end of each discharge cycle. This indi- cates that due to convection from stirring the electrolyte, the local pH value changes within the porous electrode better become measurable in the bulk electrolyte by stirring or rather mixing the local electrolytic domains. Still, the pH stability for the other EC with AA, PPA, and FAþ AA acids addition can be increased and the mass-transport limitation decreased by the stirring, which is shown in a more stable pH level within the first 10 cycles. In addition, for the increased concentration of the acid, a higher pH buffer capacity and therefore a more stable pH level with smaller periodic pH amplitudes during cycling can be observed. Generally, observing the pH curve characteristics of the EC with buffer additives in Figure 10, after the initial rest phase and initial discharge, the pH curves oscillate around a constant pH plateau without a visible increasing trend. Thereby, regarding ZnSO4þ AA (s. Figure 10D), very little oscillations around a very constant pH value are visible which might be a combined effect related to the combination of an increased concentration as well as stirring. For a closer look at the effect on the EC with buffer acids on the potential curves in Figure 11Aþ B, the first discharge cycle is visualized by plotting the potential versus Zn/Zn2þ over the areal discharge capacity qDC. All EC-containing buffer additives show lower discharge capacities compared to the RE. Nevertheless, ECs with increased acid concentrations (1 M AA, 0.5 M PPA) show significantly higher potential plateaus (þAA at �1.7 V and þPPA at �1.6 V vs Zn/Zn2þ) than the RE (at �1.2–1.4 V vs Zn/Zn2þ). Additionally, for the EC ZnSO4þAA, a stable sin- gle potential plateau can be observed, which can be explained by the pH-dependent Nernst equation (introduced by Mateos et al. in ref. [55]) and the high pH buffer capacity of the 1 M AA addi- tion. The þAAþ FA electrolyte shows two different plateaus comparable to the RE ZnSO4, which can be related to the reduc- tion peaks in Figure 9D, and which can be explained by zinc cor- rosion phenomena, as already observed in Figure S6, Supporting Information, for FA-containing EC. The effect of the addition of the buffer acids on the potential plateau gets even more 0 12 24 36 48 0.6 0.9 1.2 1.5 1.8 2.1 time / h E pH 2 3 4 5 6 7 pH 0 12 24 0.6 0.9 1.2 1.5 1.8 2.1 time / h E pH 2 3 4 5 6 7 pH 0 12 24 0.6 0.9 1.2 1.5 1.8 2.1 time / h E pH ZnSO4 + FA + AA 1 2 3 4 5 6 7 pH A B C D E / V v s. Z n/ Zn 2+ E / V v s. Z n/ Zn 2+ E / V v s. Z n/ Zn 2+ ZnSO4 ZnSO4 + AA 0 12 24 36 48 60 0.6 0.9 1.2 1.5 1.8 2.1 E /V vs .Z n/ Zn 2+ time / h E pH ZnSO4 + PPA 2 3 4 5 6 7 pH Figure 10. Potential curves in parallel to the pH curves dependent on the time for the first 10 cycles of the long-term stability tests considering stirring for A) the RE ZnSO4 (2 M), B) ZnSO4þ AA (2 Mþ 1 M), C) ZnSO4þ PPA (2 Mþ 0.5 M), and D) ZnSO4þ FAþ AA (2 Mþ 0.1 Mþ 0.1 M). Note: C) The open circuit voltage (OCV) phases for the þPPA electrolyte can be ascribed to a test plan bug, but should not affect the cycling behavior, as the pH value also stays constant within the OCV phases. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (13 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de pronounced by visualizing the potential over the state of dis- charge (normalized to the first discharge cycle) in Figure 11B. To consider the effect of higher potential plateaus, for the LSTs in Figure 11C, the areal-specific discharge energy eDC instead of the areal-specific discharge capacity qDC is regarded based on the equation for the energy of a battery (E ¼ QU), which better addresses the demands for battery applications. Therefore, in Figure 11C, eDC is shown over the cycle number. Compared to the RE, the energy levels for the EC are generally lower. Still, for the RE in Figure 11C, within the first 30 cycles a strong energy fading can be seen, which is followed by a low energy plateau at �1.0 mWh cm�2, and which can be attributed to the ZHS precipitation and accumulation (and its pH buffer effect at �5.5, s. ref. [42]). In comparison, the initial energy fading of ZnSO4þ AA and ZnSO4þ FAþAA within the first 30 cycles is lower compared to the first discharge cycle. Furthermore, after the initial energy drop, the energy increases again approaching a constant energy plateau of �1mWh cm�2 after �150 cycles. This behavior gets especially visible in plotting of the energy retention (normalized to the first discharge cycle) over the cycle number (s. Figure 11D). The reason for the strong initial capacity fading with pH buffer additives could be related to the lack of pre-dissolved Mn2þ ions, which deteriorates the reversibility of the Mn2þ/MnO2 deposition/dissolution reaction within the EC, which needs to be investigated in further studies using EC with Mn2þ preloading. Still, as for SES, the energy density is less important than a stable energy and capacity output over many cycles and the lifetime, the lower specific energy level of the EC with buffer additives is acceptable. In contrast, the higher potential plateaus and the stable energy curve is an important finding, which emphasizes the importance of modified EC, especially stabilizing the pH changes during cycling. In addition, these investigations should serve the further selection of buffer substances and the addition of Mn2þ to the EC, which is supported by the herein presented selection criteria for ARZMB applications. Finally, although stirring showed a positive effect regarding the buffer behavior, the implementation of stirring in a battery stack could be complicated. Therefore, an alternative electrode substrate that is highly porous could help to improve the natural convection in the cell, which is part of further studies. The most important challenge in the future will be to improve the significantly higher initial energy and capacity fading of electrolyte compositions containing pH buffer additives. Hereto, latest literature reported that the reason for the initial capacity fading might be the lack of Mn2þ ions in the electrolyte since no manganese was dissolved in it in advance.[21,27–29,33,36,46,51,52,60,70,72,84–95] Nevertheless, the significant difference between pure ZnSO4 and electrolyte com- positions containing pH buffer additives regarding the initial energy and capacity loss needs to be further investigated, which will be discussed in the following section. 2.4. Recommendations for Designing Buffer Electrolytes for ARZMB For the investigations in this study, different EC based on a zinc salt and pH buffering agents are presented. As discussed previously, the preloading of the EC with Mn2þ ions leads to a better overall cell performance and cycle stability. Within this study, we deliberately avoided the addition of Mn2þ ions to the EC to keep the electrolyte solutions as simple as possible, which is due to the applicability of the Henderson–Hasselbalch equa- tion (s. Equation (S9), Supporting Information). To this equation, 0.0 0.5 1.0 1.5 2.0 2.5 0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0 qDC / mAh.cm-2 0 20 40 60 80 100 0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0 state of discharge / % 0 100 200 300 400 500 0 1 2 3 4 e D C mc.h W m/ -2 cycle number 0 100 200 300 400 500 0 50 100 150 200 %/ noitneter ygrene cycle number X A B C D ZnSO4 ZnSO4+AA ZnSO4+PPA ZnSO4+AA+FA ZnSO4 ZnSO4+AA ZnSO4+PPA ZnSO4+AA+FA ZnSO4 ZnSO4+AA ZnSO4+PPA ZnSO4+AA+FA ZnSO4 ZnSO4+AA ZnSO4+PPA ZnSO4+AA+FA E / V v s. Z n/ Zn 2+ E / V v s. Z n/ Zn 2+ Figure 11. A) Potential curves versus Zn/Zn2þ for the selected EC ZnSO4þ AA, þPPA, and þAAþ FA compared to the RE ZnSO4 over the areal dis- charge capacity, and B) the same potential curves normalized to the state of discharge. C) The energy curves for the selected EC over the cycle number, and D) the energy retention normalized to the first discharge cycle. Note: The cycling of the þAAþ FA cell stopped due to zinc corrosion phenomena, in accordance to Figure S6, Supporting Information. The cycling of the RE cell was stopped intentionally. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (14 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de the following assupmtions apply (s. Table 4). Despite the assumptions for the Henderson–Hasselbalch equation do not fully apply to the herein investigated aqueous system, the equa- tion still helps for demonstrating the concept of designing pH- buffered aqueous electrolytes in acidic environment for ZMB (Table 4). Based on the equation, the following conclusions can be drawn for the application in the system considered here (Figure 12). According to Equation (S13, Supporting Information), the buffer capacity maximizes at the point pH= pKa (s. Figure 4D). Furthermore, for the ARZMB with the dissolved Zn2þ and Mn2þ hexa–aqua complexes, an aqueous system with an acidic character is present. Therefore, buffer acids with pKa values in the range of the desired pH value of the EC (or slightly above for less acid strength to meet the acidic character of the EC) are necessary to minimize the further reduc- tion of the pH through their addition to the EC. Furthermore, due to their high pKa value, the acid strength is low, and the dis- sociation only takes place to a certain degree, leading to a good (de-)protonation behavior for these acids. To map speciation in more complex systems, the resulting nonlinear system of equa- tions must be solved numerically, as implemented for example in common programs such as PHREEQC and MINTEQ.[96,97] This type of model could be used in future studies to investigate addi- tional influences such as complexation behavior and the use of multiple electrolyte salts but is beyond the scope of this article. In addition to the design parameters followed by the Henderson–Hasselbalch equation, the ratio between the Zn2þ and Mn2þ salt concentration within the EC needs to be set accord- ing to the reaction mechanism: as the major reaction mechanism should be based on the Mn2þ/MnO2 deposition/dissolution, a Zn2þ/Mn2þ ratio of 1 should consequentially be selected for the EC. By this, the Zn2þ/Mn2þ deposition during charging and the Zn/MnO2 dissolution during discharging can symmetri- cally be initiated for the anode and cathode, respectively. Altogether, the following summary for the key parameters to design a pH–buffer electrolyte for ARZMB in acidic environment can be formulated (Table 5). 3. Conclusion This study is intended to introduce a general conception for designing pH-buffered electrolytes for the zinc–manganese dioxide battery chemistry with the Mn2þ/MnO2 deposition/ dissolution. Therefore, in a first step, selection criteria for the buffer substances are defined and different buffer concepts based on a zinc salt as the main component are introduced, such as an AS (ZnSO4þ acid), a BS (zinc saltþ conjugated acid), and an HS (ZnSO4þ combination of two acids for targeted initial pH value). The contemplable buffer substances are selected based on the selection criteria and the following acids are chosen within this study: AA, PPA, FA, CA, 4-hydrobenzoic acid, KHS, KDHC, and KHP. For evaluating the buffer behavior of the selected acids, titra- tion experiments are performed for the buffered electrolytes and compared to the RE ZnSO4. A pH buffer behavior can be proved for all the acids and quantified using titration curves and by introducing the pH buffer capacity behavior by using fitting and derivation techniques of the titration results, respectively. To validate this finding in full cells, the electrolyte composi- tions are cycled, and the pH value of the electrolyte is tracked in operando. All the selected electrolytes are cyclable and the Table 4. Comparison of the assumptions for the validity of the Henderson–Hasselbalch equation and the conditions in the context of the experiments within this study. The green shading indicates a good validity, the yellow shading indicates an approximation, and the red shading indicates no agreement of the conditions for the Henderson– Hasselbalch equation and the experiment.[100,101] Conditions Henderson–Hasselbalch equation Experimental conditions 1. Monobasic acids that release an Hþ ion upon dissociation Valid within the context of the herein considered dissociation degree for all the acids 2. Self-ionization of water negligible (mostly for pH values <10) Valid in the context of this study 3. Low ion concentrations <1mol L�1, influence of ionic strength neglected Ion concentrations partly >1mol L�1 4. Weak acids with pKa value between 5 and 9 Mainly valid in the context of this study 5. Binary electrolyte consisting of (weak) acid and corresponding (weak) base Not valid in the context of the experiments (systems consisting of mixtures of substances) 6. Ultrapure water with pH= 7a) Not given due to open systems, for pH< 5.5, the dissociation is negligible a)By contact with the atmosphere, CO2 is absorbed and the pH can drop to�5.8.[100] Figure 12. Conclusions for designing pH-buffered electrolytes for ARZMB on the basis of the Henderson–Hasselbalch equation as an aid to under- stand the pH buffer concept. Table 5. Summary of the key parameters recommended for EC for ARZMB as a result of this study. Parameter Criterion Molar concentration ratio c(Zn2þ)/c(Mn2þ) 1 Molar concentration c(Mn2þ) bzw. c(Zn2þ) >0.5 mol L�1 Initial pH value pH 4.5� 1 pH–operation window �3.5–5.5 Ionic conductivity >10 mS cm�1 Electrochemical stability in the potential window 0.5–2.1 V versus Zn/Zn2þ Safety & toxicity No CMR substances, peroxides, sulfides www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (15 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de pH buffer behavior can again be proved. Still, for the different acids, impacts on the specific capacity/energy of the cell as well as the capacity/energy curve can be observed. For testing the impact of the pH-buffered electrolytes on the long-term stability, cycle tests are performed for selected electro- lytes with AA, PPA, and the combination of AA and FA. The potential plateau during discharge can be increased from �1.3 V (ZnSO4) to �1.7 V (ZnSO4þ AA) versus Zn/Zn2þ and the energy retention from �30% after 268 cycles (ZnSO4) to �86% after 494 cycles (ZnSO4þAA). Altogether, the results can serve as a basis for future electro- lyte designing, where the concepts of this study can be used as a fundamental basis to understand, describe and quantify the buffer behavior of ARZMB electrolytes and its impact of the cycle stability for the investigated reaction mechanism. The evaluation of the pH buffer capacity curves within this study can serve as a powerful method for analyzing the pH buffer capacity character- istic of an EC within the defined pH operating window. By this, a targeted electrolyte design for the respective aqueous system, e.g., the ARZMB within this study can be enabled. 4. Experimental Section Materials: The RE consisted of 2 M ZnSO4 (Zinc sulfate heptahydrate EMSURE ACS, ISO, Reag. Ph. Eur. For analysis, Supelco, Merck KGaA, Darmstadt, Germany). The electrolytes were prepared depending on their classification and solubilities: For the electrolytes of the AS, a 2 M ZnSO4 with 0.1 M of the AS buffer substances was prepared based on the RE. The buffer substances used are listed in Table 6. For the electrolytes of the BS, a solution consisting of 0.8 M ZA (ZA, AnalaR NORMAPUR, VWR Chemicals, Radnor, Pennsylvania, USA) with 4.56 M AA was prepared. In addition, a solution consisting of 0.3 M ZF (ZF 98%, Alfa Aesar, Kandel, Germany) with 1.69 M FA was prepared. Both solutions were adjusted with respect to their mixing ratios of zinc salt with conjugate acid in such a way that an initial pH of 4 was obtained (s. Henderson–Hasselbalch equation). For the electrolytes of the HS, an electrolyte consisting of 2 M ZnSO4 with 0.1 M FA and AA each was prepared. Additionally, for the cycling experiments, to reduce the surface tension of water at the porous cathode, 0.04 M sodium dodecyl sulfate (SDS, USP-NF, BP, Ph. Eur. Pure, pharma grade, ITW Reagents, Darmstadt, Germany) was added to the electrolyte solution which corre- sponds to the critical micelle concentration of SDS. The EC are summa- rized in Table 7. Characterization: The titration tests were carried out using a pH meter (SevenExcellence pH/Cond S470, Mettler Toledo (MT), Gießen, Germany) with the pH electrode MT InLabPro. The titrations took place in a temperature-controlled laboratory at 21� 1 °C. In each case, 40mL of the EC was stirred continuously in a beaker on a magnetic stirrer at 250 rpm. The initial pH of the solution was determined when no pH change could be measured for at least 300 s. Then, the NaOH and HCl solutions (first few titration steps with 0.1 M, then 1 M, and finally 10 M to deplete the buffer capacity) were each added to the solution in reasonable μl increments using a micropipette and the pH and tempera- ture for each step were determined and recorded. The titration data were then combined to map the pH buffer behavior into the basic and acidic ranges. The added volumes of titrant were converted to added concentra- tions c(Hþ) or c(OH�). In addition, the ionic conductivity was measured at standard conditions for each EC with the conductivity electrode MT InLab710 using the same pH meter MT SevenExcellence pH/Cond S470. For the comparison of the EC, 10 charge/discharge cycles were initially performed. The following test plan was used for this purpose (s. Table 8). The indication of the current rate refers to the amount of MnO2 pre-depos- ited by electrodeposition with 37.5 mg per cathode. (Table 8) Cell Assembly: The cell setup for these experiments was designed in such a way that pH measuring during cycling with the electrode MT InLab Semi-Micro was possible (Figure 13). Therefore, a cell housing with internal dimensions of 52� 29.5� 10.5mm was fabricated using a 3D Table 6. List of acids for the AS, BS, and HS. Designation Chemical formula Abbreviation Substance description of the manufacturer 4-hydrobenzoic acid C7H6O3 HBA 4-Hydrobenzoic acid 99%, Alfa Aesar, Kandel, Germany Acetic acid C2H4O2 AA Acetic acid puriss. p.a., ACS reagent, reag. ISO, reag. Ph. Eur., ≥99.8%, Sigma Aldrich, Darmstadt, Germany Citric acid C6H8O7 CA Citric acid 99þ%, Alfa Aesar, Kandel, Germany Formic acid CH2O2 FA Formic acid 98–100% for analysis, Merck KGaA, Darmstadt, Germany Potassium dihydrogen citrate KH2C6H5O7 KDHC Potassium dihydrogen citrate hydrate (�15.2% H2O), 99% (dry basis), Alfa Aesar, Kandel, Germany Potassium hydrogen phthalate KHC8H4O4 KHP Technical buffer solution pH 4.01, Mettler Toledo, Gießen, Germany Potassium hydrogen sulfate KHSO4 KHS Potassium hydrogen sulfate, VWR Chemicals, Radnor, Pennsylvania Propionic acid C3H6O2 PPA Propionic acid 99%, Alfa Aesar, Kandel, Germany Table 7. Summary of the EC within this work. Notation Initial pH Ionic conductivity [mS cm�1] Ref. ZnSO4 3.55 55.5 AS 2 M ZnSO4þ 0.1 M PPA 2.59 54.4 2 M ZnSO4þ 0.1 M AA 2.47 55.2 2 M ZnSO4þ 0.1 M FA 1.92 56.6 2 M ZnSO4þ 0.1 M KHS 1.50 62.4 2 M ZnSO4þ 0.1 M KHP 2.78 56.0 2 M ZnSO4þ 0.1 M KDHC 1.92 55.1 2 M ZnSO4þ 0.1 M CA 1.57 56.9 2 M ZnSO4þ 0.1 M HBAa) 2.90 54.7 BS 0.8 M ZAþ 4.56 M AA 3.36 10.7 0.3 M ZFþ 1.69 M FA 2.62 22.1 HS 2 M ZnSO4þ 0.1 M FAþ 0.1 M AA 1.95 55.5 a)HBA not completely soluble. www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (16 of 19) © 2023 The Authors. Energy Technology published by Wiley-VCH GmbH http://www.advancedsciencenews.com http://www.entechnol.de printer (Anycubic Photon, Hong Kong Anycubic Technology Co. Ltd., Tsim Sha Tsui, Hong Kong). The contact rails for the electrodes with the dimen- sions 25� 8� 2 mm as well as screws and contact shoes were made of stainless steel (material no. 1.4301) to avoid undesired corrosion reac- tions. The cathode was prepared by electrodeposition of Mn2þ on a carbon fiber substrate with the dimensions of 50� 15mm whereby 30� 15mm were coated with MnO2 (mass loading 37.5mg based on a specific mass loading of 10mg cm�2.[29,98] Zinc was used both as anode (zinc foil 99.95%, 25 μm, Goodfellow GmbH, Hamburg, Germany, dimension 50� 15� 0.025mm) and as ref- erence electrode (RE) in the form of a zinc wire (99.9% zinc, Goodfellow GmbH, Hamburg, Germany). To ensure a constant electrode gap (�10mm) in the cell, a cellulose/polyester separator (Spec-Wipe 3, VWR International GmbH, Darmstadt, Germany) was added to the bottom of the cell in a U-shape. The cycling data was reported with the VMP3multichannel potentiostat (BioLogic Science Instruments, Seyssinet-Pariset, France). The calibrated pH electrodes were then inserted and aligned in front of the cathode. The electrolyte was added to a level of �30mm (electrolyte volume �9mL) so that the cathode was completely immersed in it. Finally, the cells were sealed with parafilm. CV: For the CV experiments, the cell setup mentioned earlier in Cell assembly and Characterization (s. also Figure 13) was used but without pH tracking. The parameters used for the potentiostat are summarized in Table 9. Supporting Information Supporting Information is available from the Wiley Online Library or from the author. Acknowledgements This work was funded by the German Federal Fellowship (Deutsche Bundesstiftung Umwelt, DBU). Open Access funding enabled and organized by Projekt DEAL. Conflict of Interest The authors declare no conflict of interest. Data Availability Statement The data that support the findings of this study are available from the corresponding author upon reasonable request. Keywords aqueous batteries, electrochemistry, pH buffers, pH buffer capacities, titrations, zinc–manganese dioxide batteries Received: June 29, 2023 Revised: August 10, 2023 Published online: September 3, 2023 [1] U. Fegade, G. Jethave, F. Khan, A. Al-Ahmed, R. Karmouch, M. Shariq, Inamuddin, M. F. Ahmer, Int. J. Energy Res. 2022, 46, 13152. [2] Y. Wang, Z. Wang, F. Yang, S. Liu, S. Zhang, J. Mao, Z. Guo, Small 2022, 18, 2107033. [3] J. Song, K. Xu, N. Liu, D. Reed, X. Li, Mater. Today 2021, 45, 191. [4] N. Wang, X. Qiu, J. Xu, J. Huang, Y. Cao, Y. Wang, ACS Mater. Lett. 2022, 4, 190. [5] Z. Tie, Z. Niu, Angew. Chem. Int. Ed. 2020, 59, 21293. [6] E. Bellini, Salient Energy Develops Zinc-Ion Battery for Residential Applications, pv Magazine, Berlin, Germany 2022. [7] B. D. Adams, US20200176198A1, 2018. [8] Enerpoly. Sustainable Batteries for Energy Storage, https://enerpoly. com/ (accessed: May 2022). [9] Eos Energy Enterprises, https://eosenergystorage.com/ (accessed: May 2022). [10] Urban Electric Power, https://urbanelectricpower.com/ (accessed: May 2021). [11] P. Ferro, F. Bonollo, Sustainable Mater. Technol. 2023, 35, e00543. Figure 13. Cell assembly with pH electrode. Table 8. Test plan for the cycling experiments with pH measurement for the first 10 cycles of the investigated EC. # Designation Parameter 0 Rest phase t= 6 h 1 Initial discharge (CC) i= 100mA g�1, EDCH< 0.7 V versus Zn/Zn2þ 2 Charge (CC) i= 100 mA g�1, ECH> 1.8 V versus Zn/Zn2þ 3 Charge (CV) ECH= 1.8 V versus Zn/Zn2þ, t= 1.5 h 4 Discharge (CC) i= 100mA g�1, EDCH< 0.7 V versus Zn/Zn2þ 5 Loop #2–4 for 10 cycles Table 9. Test plan for the cyclic voltammetry experiments for the selected EC filtered out for the long-term stability tests. # Designation Parameter 0 Rest phase t= 1 h 1 Lower potential limit E1≤ 0.7 V versus Zn/Zn2þ 2 Upper potential limit E2≥ 1.8 V versus Zn/Zn2þ 3 Voltage scan rate dE/dt=0.1 mV s�1 4 No. of cycles n= 2 5 Rest phase t= 10 h www.advancedsciencenews.com www.entechnol.de Energy Technol. 2023, 11, 2300723 2300723 (17 of 19) © 2023 The Authors. 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