Solvation chemistry tailored via dielectric constant engineering for stable low-temperature aqueous zinc batteries

Domain: chemistry
Published: 2026-02-25
Editorial outcome: minor_revision

Abstract
Aqueous zinc-metal batteries are promising candidates for sustainable energy storage; but their practical viability is severely limited by poor cryogenic performance caused by kinetics sluggishness and interfacial instability. Here we show a strategy for low-temperature ZMBs based on tailoring the Zn2+ solvation environment by engineering the dielectric constant (ε). By incorporating ethyl acetate, a low-ε co-solvent, into a Zn(ClO4)2 electrolyte, we strategically weaken water’s hydrogen-bond network and increase cation-anion pairing. This modified solvation structure accelerates Zn2+ transport and desolvation, promotes the formation of a protective solid electrolyte interphase rich in organic and inorganic components, and inhibits parasitic hydrogen evolution. Consequently, the optimized electrolyte enhances Zn plating/stripping stability, with Zn||Zn cells operating at 0.2 mA cm−2 for 10 months (25°C) and 1 mA cm−2 for 4,000 hours (−50°C), and Zn||PANI batteries at 1 A g−1 sustaining 10,000 cycles with negligible degradation (−50°C). This work highlights the critical importance of dielectric constant engineering in electrolyte design and paves the way for high-performance, low-temperature aqueous batteries.

Introduction
Aqueous rechargeable Zn-metal batteries (ZMBs) hold significant promise for deployment in next-generation sustainable energy storage systems1–3. However, their development is hindered by challenges related to the Zn metal electrode, including dendrite growth, the parasitic hydrogen evolution reaction (HER), and a limited operating temperature range4–7. At low temperatures, strengthened intermolecular forces in aqueous electrolytes significantly reduce ionic mobility, leading to hindered interfacial charge transfer, exacerbated Zn2+ concentration polarization at the Zn surface, and ultimately, accelerated Zn dendrite formation during plating6,8–12. Therefore, developing robust electrolyte systems that address these fundamental limitations is crucial for realizing high-performance, low-temperature ZMBs. Existing strategies for aqueous electrolyte design aim to broaden the temperature range and facilitate fast ion conduction at low temperatures6,13–18, including hydrated eutectic electrolytes17, “high-entropy” electrolytes with concentrated salts18, and water-in-salt electrolytes15,16. Although these strategies have enhanced low-temperature performance, the corrosive properties of highly concentrated electrolytes undermine long-term cycling stability, restricting the practical use of current aqueous ZMBs in cryogenic conditions.

To overcome these limitations, a fundamentally different approach addressing the root causes of low-temperature failure without relying on high salt concentrations. Here, we propose that precisely tailoring the electrolyte’s solvation chemistry, specifically through the manipulation of its dielectric constant (ε), provides an effective solution. The dielectric constant, a key property governing intermolecular and ionic interactions, has been largely overlooked in low-temperature electrolyte design19. Unlike conventional approaches that focus on lowering the freezing point, modulating ε can fundamentally change the Zn2+ solvation environment and coordination dynamics, directly influencing interfacial Zn2+ diffusion and desolvation kinetics. By introducing a low-ε co-solvent, the strong tetrahedral hydrogen-bonding network of water molecules can be disrupted, leading to enhanced ionic mobility and facilitating uniform Zn deposition at low temperatures.

In this work, we proposed a dielectric matching strategy, placing the aqueous electrolyte in a medium‑permittivity window by mixing a high‑ε and a low‑ε component. Wherein high-ε solvents promote the dissociation of zinc salts and low-ε solvents optimize interface stability and ion transport kinetics at low temperatures. This concept was demonstrated by incorporating ethyl acetate (EA), a low-ε co-solvent (ε=6.4), into a Zn(ClO4)2 aqueous electrolyte. The addition of EA disrupts the strong hydrogen-bonding network maintained by high dipole-dipole interactions between water molecules (ε=78.46), modifying the Zn2+ solvation structure and leading to improved electrolyte performance at cryogenic temperatures. Through a combination of in situ Fourier-transform infrared (FTIR) spectroscopy, confocal laser scanning microscopy (CLSM), and molecular dynamics (MD) simulations, we reveal that the modified solvation structure, with increased cation-anion association, promotes low-temperature desolvation and interfacial stability. This approach effectively addresses three key challenges: (1) Sluggish Zn-ion transport at low temperatures (−50°C) is significantly improved while minimizing the depletion of active Zn2+; (2) Extensive dissociation of cation-anion pairs and HER are substantially suppressed, mitigating irreversible Zn electrode corrosion; (3) The formation of a robust and stable solid electrolyte interphase (SEI) on the Zn electrode is promoted, suppressing the HER and facilitating uniform Zn2+ deposition. This electrolyte design enables improved low-temperature performance in ZMBs. Specifically, Zn||Zn cells exhibit Zn plating/stripping stability, exceeding 10 months at 25°C and 4000 h at −50°C. Furthermore, Zn||PANI full cells demonstrate cycling life of 10,000 cycles with negligible capacity fade at −50°C, representing a significant improvement in low-temperature aqueous battery technology.

Results
Design principles for anti-freezing Zn electrolyte
Under the influence of an electric field, solvated Zn ions migrate to the electrode-electrolyte interface, where they undergo reduction to metallic Zn. The Zn2+ transfer process involves the desolvation of Zn2+ at the electrolyte-electrode interface, leading to the release of solvent molecules and deposition of Zn metal onto the electrode surface20. However, this process encounters two significant challenges: (1) water molecules released during desolvation can participate in the HER and their reversible corrosion of the Zn electrode, hindering the formation of a stable SEI. (2) Solvents with high ε exhibit a strong affinity for Zn ions21, which suppresses desolvation, thus limiting Zn2+ deposition kinetics10 (letf of Fig. 1a). At low temperatures, Zn ions exhibit reduced mobility and diffusion rates, resulting in random nucleation events across the electrode surface. This process produces metallic Zn deposits with the characteristic anisotropic hexagonal close-packed crystal structure (middle of Fig. 1a)22,23. The dielectric constant of a solvent is a crucial physicochemical property that influences cation-anion interactions in electrolytes. High dielectric constant solvents like water promote effective salt dissociation and ion transport24. In contrast, low-dielectric-constant solvents enhance cation-anion pairing, which increases the reduction potential of anions at the Zn surface and promotes the formation of a stable solid electrolyte interphase25,26. Our approach focuses on adjusting solvent formulations to optimize the solvation environment, minimize corrosion from water decomposition, and trigger the formation of a thin, robust SEI layer (right of Fig. 1a).

EA emerged as a promising co-solvent candidate (Fig. 1b) due to its low dielectric constant, low viscosity (1.9 mPas), low freezing point (−84°C), and cost-effectiveness (Supplementary Tables 1–3). Solvent freezing point and viscosity are crucial for ion migration rates under low-temperature conditions. However, the considerable difference in dielectric constants between EA and water presents a significant challenge, as phase separation occurs when these solvents are mixed. To tackle this incompatibility, we systematically investigated common Zn salts (Supplementary Fig. 1). Our screening indicated that Zn(ClO4)2·6H2O enhances the miscibility of EA and water by altering the solution’s overall dielectric response (details of salt-induced dielectric response regulation in incompatible solvent systems are given in Supplementary Discussion 2). The dielectric constant values were calculated through MD simulation27,28. The dielectric constant of the Zn(ClO4)2/H2O electrolyte is 42.48, categorizing it as a high-dielectric-constant electrolyte (HDCE). In contrast, the Zn(ClO4)2/EA electrolyte exhibits a much lower dielectric constant of 5.47. When Zn(ClO4)2 is dissolved in a water-EA mixture, the resulting electrolyte has an intermediate dielectric constant of 24.13 (Fig. 1c) and is thus designated as a medium-dielectric-constant electrolyte (MDCE). It is widely accepted that the form of ion association in high-dielectric-constant electrolytes consists of solvent-separated ion pairs (SSIPs) and is beneficial for achieving high ionic conductivity, but impedes reaction kinetics due to suppressed desolvation processes. In contrast, low-dielectric-constant electrolytes primarily feature contact ion pairs (CIPs). The resulting lower concentration of SSIP reduces overall ion transport21,29. This suggests that the Zn(ClO4)2-EA/H2O solution exhibits a moderate dielectric constant, leading to a balanced ion association structure that enhances ion transport efficiency (Fig. 1c). To evaluate the viability of these hybrid electrolytes, we dissolved Zn(ClO4)2 in a solvent with various EA:H2O molar ratios (0:1, 1:20, 1:16, 1:12, 1:8, and 1:6) (Supplementary Figs. 3 and 4). Differential scanning calorimetry (DSC) measurements revealed phase transitions (Supplementary Fig. 5). The electrolyte without EA has a crystallization temperature (Tc) of −38.5°C, and adding EA at a molar ratio of EA:H2O=1:20 reduced the crystallization temperature to −51.3°C. Increasing the EA content to a 1:16 ratio eliminated the crystallization peak, resulting in only a glass transition temperature (Tg) below −112.8°C. This decrease in Tg confirms that optimal EA content significantly improves the electrolyte’s supercooling capability, indicating that these formulations may be particularly suitable for low-temperature applications (<−40°C). We measured the ionic conductivity of Zn(ClO4)2-based solutions across various EA/H2O ratios and temperatures (Supplementary Fig. 6). Among the tested EA:H2O ratios, 1:16 delivered the highest ionic conductivity at −50°C (24.92 mS cm⁻1). This finding suggests that precise control of the EA/H2O ratio is critical for maintaining high ionic conductivity at low temperatures. The addition of low-dielectric-constant solvents (EA) likely disrupted the hydrogen-bonding network and suppressed the low‑temperature ordering of water, thereby facilitating ion transport. To assess zinc plating/stripping reversibility across compositions, we assembled Zn||Cu and Zn||Zn cells at 25°C and −50°C; both experiments identified EA:H2O=1:16 as the molar ratio for low‑temperature operation (Supplementary Fig. 7).

Spectroscopic analysis revealed significant changes in molecular interactions within the hybrid electrolyte system. To comprehensively understand the molecular mechanisms of different segments in MDCE, the coordination of Zn2+ with ClO4− anions was investigated by FTIR spectroscopy (Supplementary Fig. 8c, d). The results show peaks assigned to SSIPs (noncoordinated ClO4−, 1076 cm−1), CIPs (Zn(ClO4−)n≥1, 1047 cm−1), indicating strong adsorption for ClO4−, free anions (1114 cm−1) and free EA (1037 cm−1). The results indicate that upon the introduction of EA, the content of SSIP and free anions in the electrolyte decreases, while the concentration of CIP increases. Additionally, a fraction of EA exists as a free solvent. Furthermore, we examined O–H stretching vibrations between 3000 and 4000 cm−1 using FTIR and Raman spectroscopy (Fig. 2a, Supplementary Fig. 9). This analysis revealed progressive changes in water solvation as EA concentration increased. We observed a shift to higher wavenumber in the peaks for cluster water (3335 cm−1) and bulk water (3229 cm−1), along with a shift to lower wavenumber in the peaks for isolated water (3552 cm−1). These spectral shifts indicate progressive disruption of water’s hydrogen‑bond network: both weakly and strongly H‑bonded populations convert to non‑H‑bonded water30,31, thereby suppressing water crystallization. These results suggest an increased disruption of the hydrogen-bonding network of water and promote the conversion of weak and strong-strength hydrogen bond networks to non-hydrogen bond networks in water, which potentially inhibits the ordered arrangement of water molecules induced by a decrease in temperature.

X-ray absorption fine structure (XAFS) is effective for analyzing the solvent sheath structure of Zn ions in the electrolyte32. The Zn K-edge absorption spectra were collected for both the Zn foil and the HDCE (aqueous Zn(ClO4)2 without EA), as well as the MDCE (Zn(ClO4)2·H2O/EA electrolyte, molar ratio of EA:H2O=1:16). The MDCE displayed a lower front peak energy of Zn in the X-ray absorption near-edge structure (XANES) spectrum compared to the HDCE (Fig. 2b), suggesting differences in the local coordination environment of the Zn ions32. To resolve the Zn2+ solvation structure, we compared experimental XAFS with spectra simulated from density functional theory (DFT)‑optimized models of [Zn(ClO4)(H2O)5]+ and [Zn(ClO4)EA(H2O)4]+ (Supplementary Fig. 10 and Supplementary Data 1). In the EXAFS region, a pronounced peak at 1.65 Å corresponds to the Zn–O distance in [Zn(ClO4)(H2O)5]+ complex33. By contrast, MDCE exhibits enhanced multi‑shell features, indicating a more heterogeneous coordination environment. Wavelet‑transform EXAFS further separates these environments and reveals additional high‑k components in MDCE attributable to EA/ClO4− coordination (Fig. 2c; Supplementary Fig. 10g). The extensive regions observed in the wavelet transform indicate the coordination distance, demonstrating that EA is involved in the first solvent sheath of the Zn ion and influences its solvent structure34.

MD simulations were conducted to examine Zn2+ solvation structure and the hydrogen-bonding network at various temperatures (Supplementary Fig. 11a, b). Figure 2d and Supplementary Fig. 12 present the analysis of bond lengths, bond angles, and variations in the average hydrogen bond numbers (nHBs) among solvent molecules at 25°C and −50°C. At 25°C, the nHBs for HDCE and MDCE were calculated to be 2.32 and 2.03, respectively (Supplementary Fig. 12). Upon cooling to −50°C, HDCE exhibited a substantial increase in hydrogen bonding, with nHBs rising to 2.69, as visualized by the color change in Fig. 2e. In contrast, MDCE showed a more moderate increase, with nHBs reaching only 2.28 (Fig. 2d). This indicates that HDCE forms a significantly more hydrogen-bonded network at low temperatures compared to MDCE. This suggests that the drop in temperature enhances molecular interactions within the solvent, which could negatively impact ion transport efficiency. Furthermore, the addition of EA in MDCE significantly lowered the number of HBs, indicating a disruption of the hydrogen-bonding network compared to HDCE. A thorough analysis was carried out on the effects of introducing low-dielectric-constant solvents on solvation structure. Radial distribution functions (RDFs) and coordination numbers (CNs) are shown in Fig. 2f and Supplementary Fig. 11c, d. For MDCE, as the temperature decreased from 25 to −50°C, the average CN of Zn2+ with oxygen in EA and ClO4− increased slightly, from 0.19 to 0.21 and from 1.35 to 2.12, respectively (Fig. 2f and Supplementary Fig. 11d). HDCE exhibited a pattern of modest CN changes with reducing temperature (Supplementary Fig. 11c). The solvation structure analysis revealed differences between the electrolyte systems (Fig. 2g). The MDCE retained most of the mobile Zn2+ carriers (solvent-surrounded Zn2+, SSZ2+) and Zn2+-ClO4− (partial contact ion pairs, pCIP+) across varying temperatures, thus ensuring high ionic conductivity and contributing to lower overpotentials in symmetric cells. The number of Zn2+-ClO4− (full contact ion pair) structures increased significantly in the Zn2+-ClO4− (Zn2+-anion cluster, ZAC−). This transition may be linked to the introduction of a low-dielectric-constant solvent; the reduced dielectric constant of EA diminishes electrostatic screening between ions. This change permits ClO4− to approach Zn2+ more readily, fostering the formation of additional anion–cation–anion interactions21. DFT frontier‑orbital analysis shows that EA‑containing CIP complexes have a smaller HOMO–LUMO gap than CIPs without EA, implying greater electron‑transfer propensity and facilitating anion‑derived SEI formation (Supplementary Fig. 11e)35.

The ionic conductivity results at various temperatures demonstrate that incorporating the low-dielectric-constant solvent EA improves ion transport at low temperatures. Although EA reduces the extent of ion pair dissociation, it reduces the viscosity of the electrolyte at low temperatures and enhances ion mobility, resulting in a net improvement in conductivity. Additionally, the moderate level of ion pairing promotes the formation of a stable SEI layer, ultimately enhancing battery performance.

Solvation kinetics of interfacial chemistries
To better understand the solvation structure and examine ion pairing in both electrolytes, we employed a potentiometric method, with details provided in Supplementary Fig. 13 (the derivation of cell potential is outlined in Discussion 1). In brief, the measured open circuit potential (OCP) across the cell is directly related to the difference in free energy of solvation between the two electrolytes, influenced by variations in ion pairing and ion-solvent interactions. This method allowed us to determine the relative degree of cation-anion association in MDCE compared to HDCE. Because stronger binding yields more negative solvation free energies, the more negative ΔGsolv measured for MDCE (Fig. 3a; Supplementary Table 6) indicates tighter Zn²⁺ binding and, correspondingly, a higher fraction of ion pairs in the liquid36.

At the molecular level, ion desolvation disrupts both ion-solvent and ion-ion interactions. The restricted thermal movement of ions and molecules at low temperatures makes the desolvation process slower. For zinc-based systems, Zn2+ desolvation is recognized as the rate-limiting step, and lower desolvation energies promote more efficient ion transport37. We determined the activation energy (Ea) for Zn2+ desolvation kinetics in both electrolyte systems using Arrhenius temperature dependence analysis38 (Fig. 3b, Supplementary Fig. 14). The resulting Ea values were 22.1 for HDCE and 20.9 kJ mol−1 for MDCE, with the latter being lower than values reported for most comparable electrolyte systems39–42. To provide molecular-level insights into these differences, we calculated desolvation energies of various Zn2+-solvent complexes using DFT (Supplementary Fig. 15). The analysis revealed that [Zn(H2O)x(EA)]2+ complexes (where x=1–5) exhibit sequential water removal energies ranging from 0.77 to 3.11 eV, substantially lower than conventional [Zn(H2O)x]2+ (x=1–6) (1.03–4.90 eV). Similarly, in [Zn(H2O)5−x(ClO4−)]+ complexes (x=1–5), desolvation energies (0.61–2.53 eV) are markedly higher than those for [Zn(H2O)4−x(ClO4−)(EA)]+ (x=1–4) (0.48–1.0 eV).

Zn2+ desolvation at the electrode interface involves multiple proton-electron transfer processes (2H2O→H2+2OH−, H+ + 2e−→H2) that can negatively impact battery performance. Suppressing interfacial water activity is crucial for minimizing these parasitic reactions. In situ FTIR was utilized to investigate the evolving solvation structure during Zn plating (Zn2+ desolvation) and Zn stripping (Zn2+ solvation) (Fig. 3c and Supplementary Fig. 16a). In MDCE, significant spectral changes occurred during the initial discharge/charge cycle, including bands related to decomposition and species formation for interfacial H2O (νOH), EA (νCH3, νOCH2), and ClO4− (νClO) (Fig. 3e). These changes illustrate the dynamics of weak hydrogen bonds (W-H) and strong hydrogen bonds (S-H) of interfacial water, that is water at the interface between the electrode surface and the electrolyte solution (Fig. 3f). In HDCE, the increase in W-H signal during plating indicates higher water activity and stronger HER (Supplementary Fig. 16b). Conversely, MDCE exhibited weaker W-H signals and reduced νClO intensity during desolvation (Fig. 3g, Supplementary Fig. 16c, d), suggesting suppressed water activity and a weaker dissociation of Zn2+-ClO4− ion pairs. This implies that ClO4− ions likely form a passivation layer on the Zn electrode surface in the MDCE system, thereby inhibiting dendrite growth and mitigating structural degradation of the electrode caused by Zn corrosion. Consistent with the FTIR transients pointing to early interphase formation, linear sweep voltammetry (LSV) differentiates the two electrolytes (Supplementary Fig. 17). The full-scale LSV curves in Supplementary Fig. 16a show the behavior of both electrolytes. As can be seen from the enlarged curves in Supplementary Fig. 17b, the first reduction scan in the HDCE reveals HER starting at −0.86 V. In contrast, the first reduction scan in the MDCE has a reduction peak starting at −0.71 V and ending at −0.82 V, which reaches its maximum at −0.76 V (Supplementary Fig. 17c). The reduction peak can only be observed in the first cycle, suggesting it is linked to SEI formation43.

Zn2+ transport and charge transfer kinetics at low temperatures
The Coulombic efficiency (CE) was evaluated for both the HDCE and MDCE systems44. The MDCE showed high CE values of 97.1% (Supplementary Fig. 18). In contrast, the HDCE demonstrated a CE of only 87.1%, with performance worsening due to the loss of active Zn to electrochemically inactive Zn0 during cycling. The decline in CE of the HDCE system is attributed to the continuous depletion of Zn, which is exacerbated by the occurrence of HER. HER produced Zn5(ClO4)2(OH)8 as a byproduct17, further intensifying Zn depletion. In situ two-dimensional pH characterization was conducted using CLSM to monitor the HER. In aqueous electrolytes, the HER consumes protons (H+), potentially leading to a localized increase in pH. By introducing a trace fluorescent pH-sensitive probe into the electrolyte, real-time changes in fluorescence intensity near the Zn electrode indicate pH shifts (Fig. 4a, with detailed images shown in Supplementary Fig. 19a–d). In the HDCE system, an increase in fluorescence intensity was observed near the electrode as discharge time progressed, while the higher pH region gradually diffused into the bulk electrolyte. Conversely, the MDCE system exhibited only a slight change in the fluorescence intensity field, indicating effective suppression of HER (Fig. 4b). LSV measurements (Supplementary Fig. 19e) supported these findings. At a potential of −0.252 V, the HDCE system demonstrated a current response of 100 μA, indicative of water decomposition, while the MDCE maintained a current response of 0 μA. This confirms the MDCE system’s ability to effectively inhibit water decomposition across a broad electrochemical window. Specifically, the MDCE has a wide potential window of 2.82 V, significantly exceeding the narrower stable potential window of 2.34 V in the HDCE.

The kinetics of Zn2+ movement and charge transfer at low temperatures were analyzed. Zn||Cu cells were constructed and tested for high-capacity deposition at 25°C, −40°C, and −50°C, with 5 mAh cm−2 applied at a current density of 0.5 mA cm−2 (Fig. 4c, d). When voltage is applied, the electrode potential shifts swiftly, leading to the formation of crystal nuclei on the initial Zn substrate as the electrode potential reaches its maximum; this peak is linked to the nucleation process known as nucleation overpotential (Vn). Following this peak, Vn decreases because crystal growth demands less driving force than that for the formation of crystal nuclei. When the electrode potential stabilizes, the potential difference during this phase is denoted as ΔV. Initial tests at 25°C indicated that the HDCE exhibited a slightly lower Vn (46 mV) and ΔV (29 mV) than the MDCE (Vn=58 mV, ΔV=34 mV). This minor discrepancy is due to a slight decrease in ion transport speed within the MDCE system, attributed to the inclusion of EA into the solvation structure. However, upon reaching the nucleation overpotential at −50°C, the MDCE system showed a substantial reduction in both Vn and ΔV values compared to the HDCE system. This suggests a significantly lower energy barrier for Zn nucleation and growth at lower temperatures when utilizing the MDCE electrolyte. As the temperature drops, the sluggish movement of Zn2+ in HDCE does not compensate for the depletion of Zn ions during the deposition, resulting in a diminished concentration of Zn ions on the electrode surface and an increased polarization potential.

Subsequent examination of the disassembled cells revealed a significant difference in the morphology of the deposited Zn (Supplementary Fig. 20). In the HDCE system, Zn deposition on copper electrodes was irregular, with a considerable decrease in deposited Zn from 25 to −50°C; at −50°C, a trace amount of deposits was observed. This irregular morphology, characterized by flaky Zn deposits of various sizes across all temperatures, could lead to battery failure due to the rapid formation of Zn dendrites. The high η values suggest that a higher energy barrier hinders Zn ion deposition on the Cu current collector at low temperatures, causing many Zn ions to become disconnected and electrochemically inactive, referred to as dead Zn. In contrast, the MDCE system consistently produced dense, dendrite-free Zn structures throughout the entire temperature range studied, demonstrating enhanced performance in promoting uniform and dense Zn deposition even under cryogenic conditions.

Battery stability was evaluated under rapid cooling conditions. After 27 h of operation at 25°C, the temperature was lowered to −50°C within 60 min. The MDCE system maintained consistent Zn deposition and stripping profiles with stable voltage outputs throughout this transition, while the HDCE system exhibited noticeable soft-shorting (Supplementary Fig. 21a). Furthermore, MDCE enabled the stable Zn plating/stripping of symmetric Zn||Zn cells, while those with HDCE experienced rapid failure at a Zn plating specific capacity of 0.2 mAh cm−2 and a current density of 0.2 mA cm−2 (Fig. 4e). In contrast, MDCE-based cells demonstrated improved stability under identical conditions. These cells showed an initial gradual decrease in voltage polarization during the first 130 h, attributable to electrode activation and SEI formation processes. Following this initialization phase, the MDCE-based cells exhibited stable performance for over 10 months. MDCE provided improved performance under even more demanding conditions at −50°C. Zn||Zn cells with MDCE maintained stable operation for over 2800 h at a specific capacity of 1 mAh cm−2 and a current density of 1 mA cm−2 (Fig. 4f). At a specific capacity of 0.2 mAh cm−2 and a current density of 0.2 mA cm−2, the Zn||Zn cell incorporating MDCE exhibited improved stability for 4000 h (Supplementary Fig. 21b). These results clearly demonstrate MDCE’s improved efficacy in stabilizing Zn ion deposition and stripping processes under thermal conditions.

The long Zn plating/stripping performance of the Zn||Zn symmetric cells containing MDCE was improved under deep charging and discharging (with both charge and discharge durations of 10 h), high areal current density (20 mA cm−2), and low temperatures (−50°C). After Zn plating/stripping for 100 h at an areal capacity of 2 mAh cm−2 (with an areal current density of 0.2 mA cm−2), an irreversible increase in polarization voltage emerged in cells with conventional electrolyte, while the cell with MDCE operated stably for 500 h (Supplementary Fig. 21c). When subjected to an high current density of 20 mA cm−2, the robustness of the MDCE system became even more evident. The Zn||Zn cell with MDCE maintained stable performance for 1400 h, whereas the Zn|HDCE|Zn battery experienced significant voltage polarization by the 9th hour (Supplementary Fig. 21d). This level of long-term operational stability, both at ambient and low temperatures, rivals existing reports (Supplementary Table 5 Supplementary Fig. 22), and additional data further supports this long-cycle stability claim (Supplementary Fig. 21e). The MDCE-based Zn||Cu cell achieved 1200 cycles with a high CE (98.4%), while its HDCE counterpart displayed substantial voltage fluctuations after only 32 cycles (Supplementary Fig. 23a, b). Notably, cells utilizing the MDCE achieved a high CE over 1200 cycles at −50°C, with an average CE of 99.3% (Supplementary Fig. 23d, e).

Role of the solvation structure in determining SEI properties
The formation of a stable SEI in aqueous systems presents significantly greater challenges compared to nonaqueous electrolytes. In aqueous ZMBs, the HER competes vigorously with anion reduction at the Zn surface, thereby limiting the electrons available for anion reduction and subsequent SEI formation25,45. Notably, our findings indicate that the incorporation of EA into the electrolyte significantly promotes the establishment of a robust SEI.

Scanning electron microscopy (SEM) was employed to examine the morphology of Zn deposits at various applied current densities. As the current density increased, Zn deposition in the HDCE system became more dendritic and disordered, whereas the MDCE system produced dense and uniform Zn deposits on the electrode surfaces. The corresponding X-ray diffraction (XRD) results further confirm that Zn deposition was dominated by (002) crystallographic surfaces, with the redeposited Zn forming uniform and compact plate-like structures instead of random orientations, which enabled the batteries to exhibit improved reversibility46 (Supplementary Fig. 24). Measurements of deposition thickness after 50 cycles revealed differences between the electrolyte systems. While MDCE produced a thin (3 μm) and dense Zn deposition layer, HDCE resulted in a thick (30 ± 5 μm), porous deposit characterized by a corroded surface (Supplementary Fig. 25). In situ electrochemical optical microscopy demonstrated progressive surface roughening and dendrite formation in the HDCE over 120 min, while MDCE maintained smooth, compact deposits throughout the plating process (Supplementary Fig. 26). Furthermore, the increased thickness of Zn plating in HDCE indicated a higher degree of porosity, which rapidly accelerated the depletion of both the electrolyte and Zn ion47,48. The surface of the plated Zn in HDCE and MDCE was examined by cryo-transmission electron microscopy (cryo-TEM). Energy-dispersive X-ray spectroscopy elemental mappings, combined with high-angle annular dark field scanning, revealed the uniform dispersion of Zn, Cl, and O in this layer (Supplementary Fig. 27). The region shown in Fig. 5a was identified by high-resolution cryo-TEM (HRTEM) as a random distribution of ZnCO3 (0.281 nm) and ZnO (0.275 nm), characteristic of a mosaic-type SEI, in the system of MDCE. The polycrystalline nature of this SEI was confirmed by fast Fourier transform (FFT) analysis. In contrast, HDCE surfaces exhibited only metallic Zn with a (101) plane orientation (d-spacing 0.209 nm) and showed no evidence of a protective SEI layer49,50.

XPS depth profiling was used to examine the evolution of SEI compositions (Supplementary Fig. 28). Figure 5c, d present the corresponding C 1s, Cl 2p, and Zn 2p3/2 spectra for MDCE, respectively. As shown in Fig. 5c, the organic species (C–C/C–H, C=O, and CH3COO−) are observed in MDCE, which is attributed to the reduction of EA. The Cl 2p spectra indicate the presence of inorganic species, including Zn–Cl and Cl–C5, both associated with anion-derived SEI. The Zn–Cl signal remained dominant deep within the SEI layer (Fig. 5d). Furthermore, the Zn–Cl (1023.1 eV) and Zn–O (1022.5 eV) were detected from the SEI layer. These results suggest that the SEI layer contains organic and inorganic species (Fig. 5e). Time-of-flight secondary ion mass spectrometry (TOF-SIMS) depth profiling was also employed to analyze depth dependence, with additional information provided by the three-dimensional structure and composition of the SEI. The analysis tracked the depth profiles of both organic (CH3COO−, C−) and inorganic (Cl−) fragments, revealing the presence of decomposition products from Zn salt and EA solvent breakdown. As sputtering time increased, the oxygen content remained relatively stable; however, the Cl intensity in MDCE (1.19 × 105) was significantly lower than in HDCE (3.05 × 105) systems. This variation indicates decreased f-CIP and ZAC− species decomposition in the MDCE (Supplementary Fig. 29a, b).

Depth profiling revealed distinct distribution patterns for inorganic and organic species. Inorganic species (Cl−) maintained stable, high intensities throughout the SEI layer, while organic fragments (C−, CH3COO−) showed a noticeable decline in intensity with increasing depth (Supplementary Fig. 29d). Three-dimensional visualization of the ion fragments (Fig. 5f) provided spatial insights. In the MDCE, the SEI consisted of organic components whose concentration diminished with depth. In contrast, the Zn electrode in the HDCE system showed no detectable organic components (Supplementary Fig. 29c, e). Notably, no SEI was present in HDCE. However, numerous ClO4− and Cl− fragments were found, indicating the occurrence of side reactions in HDCE that resulted in the formation of byproducts such as Zn5(OH)8Cl2 and Zn4ClO4(OH)717. DFT indicates that EA undergoes single‑electron reduction; the resulting radical/carbanion then deprotonates to an alkene, while acetate (CH3COO−) coordinates Zn2⁺ and contributes to forming an organic‑rich SEI. The computed reaction free energy for the overall decomposition reaction EA→CH3COO−+CH2CH3 is −1.12 eV at 298 K and −1.06 eV at 233 K, indicating a thermodynamically favorable pathway that persists at low temperature (Supplementary Fig. 30).

Combining these TOF-SIMS results with microscopy and XPS data reveals a unique composite SEI structure in the MDCE system (Fig. 5g), consisting of two primary layers: an outer organic/inorganic composite layer and an inner layer dominated by inorganic components. The inorganic components are evenly distributed and more abundant than the organic components throughout the SEI. This strategic SEI composition directly influences the electrochemical performance of the MDCE system. This two-layer structure promotes a specific Zn2+ ion transport mechanism: the outer organic layer, acting as a water-blocking barrier, effectively suppresses side reactions such as HER and zinc corrosion. Benefiting from its porous structure and flexibility, this layer not only exhibits enhanced compatibility with the electrolyte (facilitating Zn2+ transport) but also adheres tightly to the inner inorganic SEI layer, preventing crack propagation at the interface caused by stress accumulation. Meanwhile, the inner inorganic layer supports efficient Zn2+ transport while inhibiting dendrite formation, directly contributing to the uniform deposition patterns and extended cycling stability observed in the MDCE system51.

Electrochemical performance of Zn||PANI cells
We evaluated the low-temperature performance of MDCE using full cells of Zn||PANI batteries6,52. The resulting cells exhibited improved performance over a temperature range from 25 to −50°C (Fig. 6a). The charge-discharge profiles of Zn||PANI batteries at a current density of 1 A g−1 are shown in Fig. 6a and Supplementary Fig. 31. The battery with HDCE exhibited a high capacity of 127.9 mAh g−1 at 25°C. However, the capacity rapidly decreased to 14.19 mAh g−1 at −50°C due to significantly hindered ion transport kinetics and reduced desolvation efficiency at low temperatures. In contrast, the MDCE system maintained 48.5% of its room-temperature discharge capacity at −50°C, showcasing effective modulation of the electrode-electrolyte interface across the temperature range. The MDCE effectively regulated the interface for Zn deposition, ensuring long-term cycling stability of the battery at both ambient and low temperatures. While PANI|HDCE|Zn cells failed within 2000 cycles due to Cl−-induced degradation of the Zn electrode (from ClO4− decomposition), the PANI|MDCE|Zn cell displayed improved durability. At a high current density of 3 A g−1, the PANI|MDCE|Zn cell achieved a near-unity CE (99.7%) and retained 98.1% capacity after 10,000 cycles (Fig. 6b).

The performance of the PANI|MDCE|Zn battery system at −50°C, maintaining 100% capacity retention and 99.9% CE over 10,000 cycles (approximately 73 days) at 1 A g−1 (Fig. 6c). To assess practical viability, we first evaluated standard coin cells configured with an 80 µm Zn electrode and a PANI positive electrode capacity of 0.13 mAh cm–2. These cells featured a high negative-to-positive capacity (N/P) ratio of 361:1 and an electrolyte-to-positive capacity (E/C) ratio of 9.8 mL Ah−1 with a 60 µL electrolyte volume (76 µL m−2) (Supplementary Fig. 32). Excessive electrolytes and low capacity in the button cell limit its large-scale application. In this context, Zn||PANI pouch cells were prepared with a carefully controlled N (10 μm, 5.86 mAh cm−2)/P (1.98 mAh cm−2) ratio of 2.96:1 and an E/C ratio of 9.8 mL Ah−1, constructed with assembled single-layer electrodes (Fig. 6d). The PANI|MDCE|Zn pouch cell delivered a high discharge capacity of 39.5 mAh (25°C) and retained 92.3% capacity over 100 cycles. At −50°C, the pouch cell achieved a high discharge capacity of 19.4 mAh and sustained 100% capacity retention over 400 cycles, demonstrating its low-temperature performance. The corresponding specific discharge capacities were 113.1 mAh g−1 (25°C) and 55.3 mAh g−1 (−50°C), based on the mass of the active materials. Furthermore, the pouch cell successfully powered the timer at approximately −50°C, showcasing the practicality of MDCE in low-temperature environments (Supplementary Movie 1).

Methods
Material preparation
The synthesis of the PANI positive electrode followed a well-established chemical oxidation method53. In a typical procedure, 3.65 mL of an aniline monomer (99%, Shanghai Macklin Biochemical Co.) was added to 150 mL of a 1 M HCl solution (37 wt%, Sinopharm Chemical Reagent Co.). Then, 150 mL of a 1 M HCl solution containing 2.28 g of ammonium persulfate (≥98.0%, Shanghai Macklin Biochemical Co.) was poured into the aniline solution and allowed to react for 1 h in an ice bath. The resulting dark green product was washed with deionized water and ethanol, and then dried in an oven at 60°C to obtain PANI powder. For the preparation of the electrolyte, Zn(ClO4)2·6H2O (>99.99%) was purchased from Sigma Aldrich, and EA (GR, 99.5%) was ordered from Shanghai Macklin Biochemical Co.

Electrode preparation and cell assembly
Zn foils (thickness: 80 μm), Cu foils (thickness: 10 μm), stainless steel (thickness: 10 μm), and stainless-steel mesh (400 mesh) were purchased from Qingyuan Metal Materials Co. All electrodes were polished with 200-grid Al2O3 powder before use to remove the oxide layer. PANI positive electrodes (diameter: 10 mm) were fabricated by casting a slurry with a weight ratio of 80% PANI powder, 10% acetylene black (Shenzhen Kejing Star Technology Co.), and 10% polyvinylidene fluoride (Shenzhen Kejing Star Technology Co.) in NMP onto the stainless steel mesh (400 mesh) using a doctor blade. Manual slicer (MSKT10, Shenzhen Kejing Star Technology Co.) used for cutting the electrode. Standard electrochemical testing used standard CR2025 coin cells (Shenzhen Kejing Star Technology Co.), evaluated at both room temperature (25°C) and low temperature (−40 and −50°C) in a controlled climate chamber (Shanghai Benting Equipment Co.). Zn||Cu cells were assembled using Zn foil (thickness: 80 μm, diameter: 10 mm) as one electrode and Cu foil (diameter: 16 mm) as the positive electrode. Symmetrical cells, including Zn||Zn and stainless steel||stainless steel cells, were assembled by using two Zn foils and stainless steel, respectively. Zn||PANI cells were assembled using Zn foil (thickness: 80 μm, diameter: 14 mm) as the electrode and PANI as the positive electrode (mass loading: 1.1 mg cm−2, diameter: 10 mm). The amount of electrolyte used was 80 µL.

Pouch cell: Zn||PANI cells were assembled using Zn foil (thickness: 10 μm, area: 4×5 cm2) as the negative electrode and PANI as the positive electrode (mass loading: 17.5 mg cm−2, area: 4×5 cm2) (Supplementary Fig. 32b). A glass fiber membrane (GF-A) served as the separator for all cells.

Electrochemical measurements
The charge/discharge behaviors of batteries were performed on a LAND battery test system. The Zn||PANI cells were charged and discharged between the voltage ranges of 0.5–1.6 V. A cut-off voltage of 0.5 V was applied in the symmetrical cell and Zn||Cu half-cell test. The CV, EIS, and CA measurements were performed on a CHI760E electrochemical workstation. EIS spectra were acquired in potentiostatic mode over a logarithmically spaced frequency range from 100,000 Hz to 0.1 mHz with 6 data points per decade and a sinusoidal amplitude of 10 mV. The Rct of the electrodes was obtained from the EIS data of the symmetric cells. CE is calculated as the ratio of discharge capacity divided by the charge capacity in the preceding charge cycle. Solvation energy measurements were conducted using a custom H-cell apparatus. A Zn metal electrode was paired with an asymmetric electrolyte, whereas a saturated ZnSO4 solution served as a salt bridge with a porous connection to separate the different electrolytes. The relative solvation energy was then assessed by measuring the OCP across the H-cell. This measurement employed two layers of hydrophilic polypropylene (PP) membrane as the porous junction previously developed in our group54. Arrhenius activation energy (Ea) was calculated by the following formula38:
1/Rct = A exp(−Ea/RT) (1)
where Rct is the charge transfer resistance, A is the frequency factor, R is the gas constant, and T is the absolute temperature.

In situ FTIR measurements
In situ electrochemical FTIR was performed on a Vertex 70 spectrometer (Bruker) coupled to a CHI760E workstation (Shanghai Chinstrument Co., Ltd.) and equipped with a liquid‑nitrogen‑cooled broadband MCT detector (Beijing bjscistar Co., Ltd.) at 25°C. Spectra were collected in absorbance mode [A=−log(I/I0)] at a fixed 60° incidence and 4 cm−1 resolution. A custom spectro‑electrochemical cell used a Cu‑mesh working electrode and Zn foil as both counter and reference; Zn was plated at 0.5 mA cm−2 for 900 s and stripped at 0.5 mA cm−2 to a 0.5 V cut‑off, the test was conducted at 25°C. Background spectra were first recorded in air, followed by A(t0) after electrolyte filling (open circuit) and time‑resolved spectra A(tx), A(ty) under current. Relative absorbance was defined as ΔA(tx)=A(tx)−A(t0) (abrupt change under bias) and ΔA(ty)=A(ty)−A(tx) (subsequent gradual evolution)55.

In situ pH measurements
Interface dynamics was investigated through fluorescence and confocal microscopy. Fluorescence measurements used an F-7100 spectrophotometer (Hitachi) operating at 240 nm/min scan speed with 5 nm excitation/emission slits and 550 V PMT voltage. Confocal imaging employed a TCSSP5II (Leica) system with 458 nm excitation and 490–550 nm detection range using a 20× dry objective at 20°C.

Characterization
A confocal inVia-Reflex (Renishaw) Raman microscope equipped with a green diode laser (λ = 532 nm) was used for the Raman analysis at 20°C. DSC (DSC 300 Caliris, Netzsch) tested the low-temperature phase transition process of the electrolytes. The XAS experiments were carried out at Shanghai Synchrotron Radiation Facility, China. The microstructure of Zn deposits was observed by SEM (SU8010, Hitachi). XPS spectra of Zn deposits were obtained by a Thermo Fisher ESCALAB 250Xi+ with an Al Kα X-ray source of 1486.6 eV. TOF-SIMS analysis of the Zn surface was measured on a PHI TOF-SIMS (nano TOF II) instrument; the ToF-SIMS depth profiling was carried out using a 30 keV Bi3++ primary ion beam (2 nA). A Gatan 698 cryo-transfer holder and a Talos F200X G2 TEM were employed to study the morphology of Zn deposits. Cryogenic transmission electron microscopy (cryo-TEM) was employed to characterize the components of Zn deposits (Thermo Fisher Talos L120C G2). The Zn||Cu cell was assembled by placing a Cu grid in contact with a Cu foil to serve as the working electrode. Zn was electrodeposited onto the working electrode with an areal capacity of 0.5 mAh cm−2 at a current density of 0.5 mA cm−2. During imaging, the specimens were maintained at a cryogenic temperature of −178°C. XRD measurements were performed on a Bruker D8 Advance diffractometer, in the range of 2θ=10.0–80.0° with a step size of 0.023° and 0.2 s per step. The conductivity of the electrolyte at 25, 0, −20, −40, and −50°C was measured using the Mettler Toledo Conductivity Meter (FE38-Standard).

Calculations
MD calculations. MD simulations were performed using GROMACS56 2024.4. The AMBER force field57, OPC water model58, and restrained electrostatic potential charge (obtained with Multiwfn59) were employed with the previously reported Zn2+ parameter of the Lennard-Jones 12-6 model (IOD set)60. All covalent bonds involving hydrogen atoms were constrained using the Linear Constraint Solver algorithm61. The initial configurations of atoms were constructed with the PACKMOL program62. The simulation cell was first defined as a cubic box with dimensions of 60×60×60 Å3, and periodic boundary conditions were imposed along all three spatial directions. Subsequently, the system was populated with 200 EA molecules (or without EA), 200 Zn(ClO4)2 units, and 3200 H2O molecules. Prior to data collection, the system was pre-equilibrated under the canonical (NVT) ensemble for 5.0 ns using a timestep of 1 fs, followed by an additional 5.0 ns equilibration under the isothermal–isobaric (NPT) ensemble with the same timestep. Production MD simulations were then carried out for 20 ns. During both the NVT production and NPT equilibration stages, the temperature was maintained at 298 K or 223 K using the Nosé–Hoover thermostat, while the pressure during NPT equilibration was regulated at 1 bar via the Parrinello–Rahman barostat. Long-range electrostatic interactions were evaluated using the Particle Mesh Ewald method63,64 with a cut-off distance of 1.0 nm. Visualization of the simulation trajectories was performed using the VESTA software65. The RDFs, g(r), describing the spatial correlations between Zn2+ cations and O atoms from ClO4−, H2O, and EA molecules were calculated according to the following equation:
g(r) = n(r) / (4πr2ρΔr) (2)
where r denotes the distance between a shell atom and the reference atom, ρ represents the average number density of a given species in the electrolyte, and n(r) corresponds to the mean number of atoms located within a spherical shell of thickness Δr, averaged over the entire simulation trajectory. By integrating g(r) over r, the CN of each species in the solvation shell surrounding the Zn2+ ion was determined. The subsequent hydrogen bonds analysis was carried out using the last 20 ns of the production run trajectories by the python code MDAnalysis66. Hydrogen bonds were identified using r_DA<3.5 Å and 150°≤∠D–H···A≤180°67. The calculation formula for the average value of hydrogen bonds is as follows:
Average HBs = (Total number of HBs) / (Total number of acceptor sites for all molecules) (3)
Dielectric constants of the solvent and electrolytes were calculated using the built-in dielectric analysis module in GROMACS, based on the final 20 ns of the production trajectory56. The OPC water model was employed for the simulations, known for its high accuracy in dielectric property simulations of aqueous systems68. The dipole moment of the system was output every 1000 steps during NVT simulations and used for the calculation of the dielectric constant. Calculation of dielectric constant using the following expression: the dielectric constant of liquid electrolytes is calculated according to Formula (4)27,28:
ε = 1 + (<M2> − <M>2) / (3ε0VkBT) (4)
where ε denotes the dielectric constant of the system, ε0 is the permittivity of vacuum, V represents the volume of the simulation box, k is the Boltzmann constant, and T is the absolute temperature. The total dipole moment of the system, M, is defined as M= Σqi ri, where qi and ri correspond to the atomic charges and their position vectors, respectively. Total dipole moment of the system (M in Formula 4) and square of it (M) were both averaged over 5000 frames, which means the dielectric constant is calculated from 5 ns to the end of the NVT production run.

DFT calculations. DFT calculations were performed using the CASTEP package. The Perdew–Burke–Ernzerhof formulation within the generalized gradient approximation was chosen as the calculation function for the above CASTEP package69,70. The cut-off energy was set to 571 eV. The convergence criteria of ionic relaxation and electronic minimization were set to 0.03 eV Å−1 and 10−6 eV, respectively. During the desolvation process, a solvated ion cluster will lose all water molecules of hydration one by one. The desolvation energy (Es) was defined as:
Es = Esolvent + EZn2+ − E[Zn2+(solvent)] (5)
Molecular‑orbital energies and EA‑reduction energetics were computed with Gaussian 1671 at the M05‑2X/6‑31++G(d,p) level, using a polarizable‑continuum (PCM) water model for solvation72.

Peer review rounds
[{"round": 1, "reviewer_comments": {"reviewer_1": "This manuscript has been revised as the comments, so it is suitable for publication without change.", "reviewer_2": "The authors provided a complete response to my questions. I suggest the consideration of this work.", "reviewer_3": "All my questions have been well addressed. It can be accepted in this journal.", "reviewer_4": "The authors addressed most of my comments. But critical details of the MD simulations are still missing. The authors said an optimized Zn2+ model was used in the study. The parameters, reference and/or other optimization details should be provided so that readers can reproduce the work.\nProper validation of the used force fields parameters for the studied systems are also needed. General force fields like OPLS can work on one system but fail badly on another. The choice of water model and other simulation protocols can change the results too. In spite of the references provided, direct comparison with experimental measurement on the same systems should be provided."}, "author_response": "Point-by-point response to the reviewers’ comments\nWe thank the reviewers for their careful assessment. Their comments led us to substantially\nstrengthen the manuscript: we clarified the scope and novelty of our dielectric-engineering strategy,\nadded mechanistic evidence and controls, corrected inaccuracies, and refined the presentation. The\nrevised manuscript and the point-by-point response address every comment with new analyses,\ndata, and citations. All textual changes are highlighted in yellow.\nClick to go to responses to a specific reviewer:\nReviewer #1 Reviewer #2 Reviewer #3 Reviewer #4\nP\nA\nG\nE\n\\\nRESPONSE TO REVIEWERS’ COMMENTS\nReviewer #1:\nIn this manuscript, the authors propose a solvation structure regulation mechanism dominated by\ndielectric constant, aiming to achieve stable cycling performance of aqueous zinc metal batteries\nat low temperatures. The assembled Zn||PANI full cell demonstrates an encouraging 100%\ncapacity retention after 10,000 cycles. Unfortunately, most of the conclusions presented are neither\nnovel nor original. Various aspects of the proposed mechanism have already been revealed by\nother researchers using similar strategies (Li, Xinpeng, et al. “Modulation of water reactivity by\nethyl acetate/water co-solvent for zinc-metal batteries.” Chemical Engineering Journal 487 (2024):\n150588. Chen, Yuqing, et al. “Breaking solvation dominance of ethylene carbonate via molecular\ncharge engineering enables lower temperature battery.” Nature Communications 14.1 (2023):\n8326.). In addition, the manuscript contains multiple apparent errors and lacks several important\ndatasets necessary to support the claims. Therefore, I don't recommend it to be published.\nResponse: We greatly appreciate the reviewer’s valuable comments. As mentioned by the\nreviewer, EA and similar ester solvents have been reported for zinc/lithium metal batteries.\nHowever, this study presents differences in mechanisms compared to previous research. In the\narticle “Breaking solvation dominance of ethylene carbonate via molecular charge engineering\nenables lower temperature battery.” Nature Communications 14, 1 (2023): 8326. Chen et al. aim\nto facilitate the coordination of low-ε solvents with Li+ to occupy its coordination sites, thereby\nconverting solvent-separated ion pairs (SSIPs) from high-ε solvent-dominated to low-ε-dominated.\nThis work targets zinc batteries with aqueous electrolytes, where the strong hydrogen-bond\nnetwork of water at low temperatures sharply reduces ionic conductivity. We propose a dielectric\nmatching strategy to construct an electrolyte with a moderate dielectric constant (ε ≈ 25.25),\nenabling synergistic regulation of Zn2+ solvation and accelerating Zn2+ transport and desolvation.\nCompared with CEJ (Li, Xinpeng, et al., Chemical Engineering Journal 487 (2024): 150588.), this\nwork employs a different electrolyte composition (EA/H O = 1:16 vs. 1:6), achieves higher ionic\n2\nconductivity and better performance at lower temperatures (−50 °C vs. −40 °C), and uses a\ndifferent zinc salt (Zn(ClO ) vs. Zn(CF SO ) ) for bilayer SEI. Li et al. and other researchers have\n4 2 3 3 2\nP\nA\nG\nE\n\\\nmainly focused on how co-solvents regulate solvation structures to improve low-temperature\nperformance. Still, they lack an in-depth investigation into how interfacial solvation structures\nevolve under an electric field and their role in promoting SEI formation.\nInstead, our contribution is to (i) define and target a medium‑ε window (ε≈25) that balances\ndissociation and ion pairing and (ii) quantify interfacial solvation dynamics under applied current\nthat link this window to fast, stable Zn plating at low temperature. The revised text now presents\ndielectric engineering as a practical screening criterion, including permittivity range and\nsolvation/interfacial checks. We have supplemented and corrected the data omissions and errors\nin the full text. We hope these clarifications and additions address the reviewer’s concerns about\nnovelty.\nOur work presents three fundamental advancements that distinguish it from the existing literature:\n1) Design strategy.\nWe proposed a dielectric matching strategy, wherein high-ε solvents promote the dissociation\nof zinc salts. In contrast, low-ε solvents optimize interface stability and ion transport kinetics at\nlow temperatures. In this work, we selected H O/EA (molar ratio 16:1) to construct an electrolyte\n2\nsystem with a moderate dielectric constant (ε≈25.25). This design has the following advantages:\n1: Greater flexibility in composition: the dielectric environment of the electrolyte can be precisely\ncontrolled by adjusting the ratio of high/low dielectric constant solvents. 2: practical screening\ncriterion: providing a quantifiable preliminary screening criterion for the design of low-\ntemperature-resistant aqueous electrolytes.\n2) Interfacial mechanism under bias (novelty).\nThe combination of in-situ infrared/confocal microscopy and theoretical calculations\nelucidates interfacial dynamics and decomposition mechanisms under an electric field. It further\ndemonstrates that incorporating organic solvents with low dielectric constants facilitates the\nformation of a bilayer SEI, providing insights into solid-liquid interfaces and solvation dynamics\nthat are crucial for cryogenic aqueous batteries.\n3) Advanced electrochemical performance.\nP\nA\nG\nE\n\\\nThe MDCE Zn||Zn cell cycled for 4,000 h at −50 °C and >10 months at 25 °C, and a negligible\ncapacity decay in low-temperature full cells over 12000 cycles, exceeding those of reported low-\ntemperature aqueous ZMBs (Supplementary Fig. 22).\nSupplementary Fig. 22 Comparison of the cycle life values of aqueous electrolyte in Zn||Zn cells\nat −50 °C-0 °C7-17.\nComment 1. The authors discuss the economic potential of EA in electrolyte formulation.\nHowever, we note that EA exhibits serious incompatibility with commonly used low-cost zinc\nsalts. As a compromise, the authors employed zinc perchlorate, which is relatively expensive and\nnot widely used in practical applications. Given that the zinc salt often constitutes a major\ncomponent of the electrolyte, I am concerned about the practical viability and cost-effectiveness\nof using EA in real-world aqueous Zn battery systems.\nResponse: Thank you for the suggestions. We will clarify this from the following two aspects:\n1. Compatibility and generality. We agree that solvent–salt compatibility is a constraint. We\ntherefore screened multiple zinc salts in EA:H O and report that both Zn(ClO ) and Zn(CF SO )\n2 4 2 3 3 2\nform single‑phase electrolytes at EA:H O=1:16 (molar). Zn(ClO ) was selected because it\n2 4 2\nprovides higher low‑temperature conductivity and better interfacial kinetics at the same\ncomposition; the Zn(CF SO ) control is now included to demonstrate that our approach is not\n3 3 2\nP\nA\nG\nE\n\\\nspecific to ClO − (conductivity at −50 °C: 24.92 vs 10.56 mS cm−1 at 1:16; corresponding CE and\n4\nnucleation metrics are provided).\n2. Electrolyte economy. This inventory of other salts, such as ZnCl , requires high concentrations\n2\n(7-30 mol L‒1) to achieve similar performance (Nat. Sustain. 6, 325–335 (2023), Nat. Commun.\n11, (2020) 4463), while relying on co-solvents like ethylene glycol (EG) may significantly reduce\nion transport efficiency due to a sharp increase in viscosity at low temperatures (Nat. Sustain. 5,\n205–213 (2022)). According to Table S3, i.e. 0.296 USD/g × 7 mol/L × 136 g/mol = 281 USD/L\nfor ZnCl electrolyte and 0.336 USD/g × 2 mol/L × 372.38 g/mol = 250 USD/L for Zn(ClO )\n2 4 2\nelectrolyte. Zn(ClO ) .6H O shows lower cost. Furthermore, Zn(ClO ) demonstrates superior\n4 2 2 4 2\nsolubility, allowing high cycling stability and rapid reaction kinetics to be achieved under low-\ntemperature conditions.\nCompared to other aqueous electrolytes and ester-based organic electrolytes, the electrolyte system\nproposed in this study shows significant advantages in cost-effectiveness and adaptability to\nextreme environments. In subsequent research, we will systematically investigate the\nincompatibility mechanisms of other low-cost zinc salts in the EA/H O system and pursue\n2\neffective regulation strategies to achieve electrolyte systems that combine cost advantages with\nhigh performance.\nComment 2. The authors have to cite relevant literatures to support your statement “ It is widely\naccepted that the form of ion association in high dielectric constant electrolytes consists of solvent-\nseparated ion pairs, whereas those with a low dielectric constant primarily feature contact ion\npairs”.\nResponse: Thank you for the suggestions. According to the reviewer’s recommendation, we\nadjusted the literature references after the statement in the manuscript on page 5, line 126:\n“It is widely accepted that the form of ion association in high dielectric constant electrolytes\nconsists of solvent-separated ion pairs, whereas those with a low dielectric constant primarily\nfeature contact ion pairs21,30, both scenarios are detrimental to ion transport.”\nP\nA\nG\nE\n\\\n[21] Yao, N. et al. An atomic insight into the chemical origin and variation of the dielectric constant\nin liquid electrolytes. Angew. Chem., Int. Ed. 60, 21473-21478 (2021).\n[30] Marcus, Y. & Hefter, G. Ion Pairing. Chem. Rev. 106, 4585-4621 (2006).\nThe corresponding description in the references is as follows: “A high dielectric constant is\nbeneficial for the dissociation of cations and anions of salts and achieving high ionic conductivity”.\n(Angew. Chem., Int. Ed. 60, 21473-21478 (2021) “…For solvents with high relative permittivity\nand electrolytes with low charges, the cutoff distance q is smaller than a. Therefore, no ion pairing\ntakes place in such cases, such as, e.g., for the larger alkali halides in water, where q = 0.357 nm\nat 298.15 K. Smaller ions and ions with higher charges, especially in solvents with lower\npermittivity, should associate to ion pairs to extents given by eqs 2 and 9…” Chem. Rev. 106,\n4585-4621 (2006).\nComment 3. The authors claim that solvent-separated ion pairs and contact ion pairs impede ion\ntransport, while electrolytes with medium dielectric constants form a balanced ion association\nstructure to enhance ion mobility. However, authors simultaneously assert that EA addition\ndisrupts the hydrogen-bond network to facilitate ion transport. Please provide experimental and\ntheoretical evidence to substantiate these claims. In addition, what is the difference between\nenhanced anion-cation pairing and contact ion pair?\nResponse: Thank you for the comment. At room temperature, solvent-separated ion pairs (SSIP)\ndetermine the number of charge carriers available for transport in electrolytes. High-dielectric-constant\nsolvents, such as water, effectively dissociate ion pairs, creating abundant charge carriers\nand enabling efficient ion transport. HDCE contains more SSIP structures, resulting in higher ionic\nconductivity at room temperature (Fig. R1a). However, at ultra-low temperatures (−50 °C), ionic\nconductivity drops dramatically due to strengthened hydrogen bonding and the freezing of water,\nwhich disrupts ion mobility.\nConversely, low-dielectric-constant solvents can significantly mitigate freezing issues but suffer\nfrom limited ion dissociation, resulting in fewer charge carriers and reduced conductivity. Our\nmethod addresses both problems simultaneously: incorporating ethyl acetate (EA) into water\nP\nA\nG\nE\n\\\ndisrupts hydrogen bonding among water molecules and prevents ice formation, while preserving\nwater’s excellent solvating properties. This maintains a high concentration of charge carriers and\nenables efficient ion movement even at ultra-low temperatures (−50 °C)\nThe introduction of EA significantly lowers the freezing point of the electrolyte, thereby\nimproving low-temperature performance. Meanwhile, the presence of contact ion pairs (CIP) can\npromote the formation of a stable solid electrolyte interface (SEI). However, as the EA content\nfurther increases, the concentration of CIP also rises (as shown by Raman spectroscopy and MD\nsimulations), which hinders ion migration and consequently reduces the ionic conductivity of the\nelectrolytes (Fig. R1a and Supplementary Table 4). This trend is clearly observable in conductivity\ntests. The key point is to regulate the concentration of CIP to synergistically optimize the freezing\npoint and ionic conductivity, thereby achieving optimal low-temperature electrochemical\nperformance.\nAdditionally, the introduction of EA disrupts the hydrogen-bond network, thereby promoting\nimproved ion transport. Molecular dynamics (MD) simulations were conducted to analyze the\nmean squared displacement (MSD) of zinc ions under both room temperature and low-temperature\n(−50°C) conditions. The results show that in MDCE (Zn(ClO ) H O/EA, molar ratios of EA: H O\n4 2 2 2\n= 1:16), zinc ion mobility is lower than in HDCE (Zn(ClO ) H O) at room temperature, whereas\n4 2 2\nat −50°C, MDCE demonstrates better zinc ion migration rates compared to HDCE (Fig. R1b, c).\nFig. R1 a, Ionic conductivities of Zn(ClO ) H O/EA solutions with different H O/EA molar ratios\n4 2 2 2\nfrom 1:0 to 6:1 at various temperatures, were obtained through measurements conductivity meter.\nb, c The MSD of the Zn2+ ion was calculated for both HDCE and MDCE at 25 °C (b) and −50 °C\n(c).\nP\nA\nG\nE\n\\\nLastly, the concepts of contact ion pairs and enhanced anion-cation pairs are not the same. The\nexpression of “enhanced anion-cation pairs” is intended to indicate that the introduction of EA\nenhances the cation–anion combination. We have revised this ambiguous expression in the\nmanuscript on page 4, line 107:\n“…low dielectric constant solvents enhance cation-anion pairing…”\nComment 4. The manuscript insufficiently addresses how low-ε solvents balance ion pair\nformation versus dissociation. The rationale for selecting EA as a co-solvent lacks systematic\nanalysis; comparative experiments with other low-ε solvents are required to validate EA’s\nuniqueness or the theory’s general applicability.\nResponse: Thank you for the suggestions. In our work, the high ε solvent is responsible for\npromoting ion pair dissociation in the electrolyte. Based on this, we formulated electrolytes with\nvarying contents of low dielectric constant solvents to explore an optimal ratio that ensures\nsufficient ion dissociation while promoting moderate cation–anion association to support SEI\nformation at low temperature while retaining adequate ionic transport, resulting in improved\nbattery performance.\nIn accordance with the reviewer's suggestions, we conducted experiments utilizing low\ndielectric constant co-solvents, specifically diethyl carbonate (DEC) (ε = 2.8), tetrahydrofuran\n(THF) (ε = 7.4), and methyl formate (MF) (ε = 8.5). Our findings reveal that the electrolyte with\nDEC led to macroscopic phase separation (Fig. R2). Notably, THF and MF did not exhibit any\nlayering. Zn||Cu batteries were assembled to investigate zinc ion plating and stripping at both room\nand low temperatures. The electrolyte containing DEC resulted in an average Coulombic\nefficiency of 91.2% at room temperature and 96.66% at low temperature (−50 °C), which is\nmarkedly lower than that of other electrolytes. Furthermore, the nucleation overpotential at low\ntemperature (η = 670 mV) was significantly higher than that observed with different electrolytes\n(Fig. R3).\nThese results suggest that electrolytes with moderate dielectric constants (for example, Zn(ClO )\n4 2\nEA/H O molar ratio EA: H O = 1:16, ε = 25.25) universally regulate solvation structures. Despite\n2 2\nP\nA\nG\nE\n\\\nhaving similar dielectric constants, THF and MF demonstrated poor electrochemical stability at\nroom temperature. While MF and THF can serve as additives in organic low-temperature\nelectrolytes, MF undergoes hydrolysis in aqueous electrolytes. EA exhibits superior stability\ncompared to MF for the following reasons: (1) the carbonyl carbon in MF is more positively\ncharged due to the weaker electron-donating ability of -H compared to -CH , rendering it more\n3\nsusceptible to nucleophilic attack. (2) In EA, the methyl group disperses the positive charge of the\ncarbonyl carbon through hyperconjugation, thereby enhancing its stability. (March, J. (1992).\nAdvanced Organic Chemistry (4th ed.). Wiley.). Similarly, THF belongs to the class of ether\nsolvents. The oxygen in the ether group contains two lone pairs of electrons, which gives it basicity\nand nucleophilicity. Influenced by electronic effects, when there is a hydrogen on the α-carbon of\nthe ether, it is prone to oxidation. (John M. Fundamentals of organic chemistry, 5th ed.), which\ncan form peroxides or decompose under anodic conditions, leading to side reactions that degrade\nelectrolyte performance and form unwanted byproducts. In contrast, EA, which belongs to\ncarbonates, has better antioxidant properties.\nBased on the research results mentioned above, we found that the introduction of co-solvents\nwith low dielectric constants can effectively enhance the low-temperature performance of batteries.\nThis phenomenon fully validates the universality of the electrolyte strategy with moderate\ndielectric constants. Notably, EA exhibits particularly outstanding low-temperature\nelectrochemical performance due to its unique physicochemical properties, such as lower viscosity\nand suitable solvation ability. This provides an important design concept and research direction\nfor the subsequent development of high-performance low-temperature resistant electrolyte\nsystems.\nWe have added the corresponding content to the main text (page 8, line 225):\n“Combining the results of ionic conductivity at various temperatures, incorporating low dielectric\nconstant solvents (EA) facilitates ion transport at low temperatures, while maintaining a balance\nbetween sufficient ion pair dissociation and moderate cation-anion association. This facilitates\nstable SEI formation, ultimately enhancing battery performance.”\nWe have also added the test results for other solvents to the supplementary information\nP\nA\nG\nE\n\\\n(Supplementary Fig. 2 and Supplementary Note 2)：\n“To systematically evaluate the impact of low dielectric constant co-solvents on the low-\ntemperature performance of batteries, we selected a series of solvents with similar dielectric\nconstants, specifically diethyl carbonate (DEC) (ε = 2.8), tetrahydrofuran (THF) (ε = 7.4), and\nmethyl formate (MF) (ε = 8.5). Our findings reveal that the electrolyte with DEC led to\nmacroscopic phase separation (Supplementary Fig. 2a). Notably, THF and MF did not exhibit any\nlayering. Zn||Cu batteries were assembled to investigate zinc ion plating and stripping at both room\nand low temperatures. The electrolyte containing DEC resulted in an average Coulombic\nefficiency of 91.2% at room temperature and 96.66% at low temperature (−50 °C), which is\nmarkedly lower than that of other electrolytes. Furthermore, the nucleation overpotential at low\ntemperature (η = 670 mV) was significantly higher than that observed with different electrolytes\n(Supplementary Fig. 2b, c). As a result, EA exhibits particularly outstanding low-temperature\nelectrochemical performance. This provides an important design concept and research direction\nfor the subsequent development of high-performance low-temperature resistant electrolyte\nsystems.”\nPhase separation\nFig. R2/Supplementary Fig. 2 Digital photograph of an electrolyte containing low dielectric\nconstant solvents such as THF, DEC, and MF.\nP\nA\nG\nE\n\\\nFig. R3/ Supplementary Fig. 2 a, b The CEs of Zn||Cu batteries at a current density of 1 mA cm−2\nand 0.5 mAh cm−2 at 25 °C (a) and −50 °C (b). c, The first discharge curve of Zn||Cu batteries at a\ncurrent density of 1 mA cm−2 and an areal capacity of 0.5 mAh cm−2.\nComment 5. In the sentence “Electrostatic portion of the metallic lithium is zero for both sides, as\nit is electrically neutral,” the mention of metallic lithium appears to be inconsistent with the rest\nof the manuscript, which focuses on zinc-based systems. This seems to be an oversight or a\ntypographical error. Please revise it accordingly to maintain accuracy and consistency in\nterminology throughout the manuscript.\nResponse: Thank you for reminding us of this. We have revised this section in the Supplementary\nDiscussion on page 1 in Supplementary Information:\n“Electrostatic portion of the metallic zinc is zero for both sides, as it is electrically neutral.”\nComment 6. There appears to be a clear error in Fig. 1d regarding the DN of water. According to\nGutmann’s widely cited work (Gutmann, V. “Empirical parameters for donor and acceptor\nP\nA\nG\nE\n\\\nproperties of solvents.” Electrochimica Acta 21.9 (1976): 661-670.), the donor number of liquid\nwater is approximately 33, not 18, which corresponds to the gas phase. Please revise the value\naccordingly to ensure consistency with established empirical data.\nResponse: Thank you for catching this; we fixed the donor number of liquid water to 33. (Figure\n1d and Supplementary Table 1)\nFig. R4/Fig. 1 Comparison of solvent properties, including donor number, dielectric constant,\nfreezing point, and viscosity.\nSupplementary Table 1. Comparison of solvent properties, including donor number, dielectric\nconstant, freezing point, and viscosity.\nDonor Freezing point\nSolvent η (mPa·s) ɛ\nnumber (°C)\nEA 17.2 −83.6 1.9 6.4\nPC 15.0 −49.2 2.5 64.9\nDMSO 29.8 −19.0 2.0 48.9\nEG 20.0 −12.9 25.66 37.0\nDMF 26.6 −62.0 1.0 36.7\nAN 14.0 −48.0 0.34 35.7\nP\nA\nG\nE\n\\\nMeOH 19.0 −97.5 0.59 32.7\nTHF 20.0 −108.5 0.55 7.4\nTMP 23.0 −46.0 1.76 21.3\nDMC 17.2 4.0 0.59 3.1\nEC 16.0 35.0 1.90 90.1\nH O ~33.0 0 0.85 78.46\nderived from EIS measurements. However, no corresponding chronoamperometry data and EIS\nfitting results are provided. According to previous studies (Yang, Jin-Lin, et al. “Designing single-\nion conductive electrolytes for aqueous zinc batteries.” Matter 7.6 (2024): 1928-1949.), the\nimpedance of aqueous Zn-ion batteries can be unstable. Therefore, for ionic conductivity\nmeasurements at different temperatures, we recommend the use of a conductivity meter to obtain\nmore reliable and reproducible results. Moreover, based on the data presented, the ionic\nconductivity of the 20(H O):1(EA) electrolyte is comparable to that of the 16:1 formulation across\n2\nvarious temperatures. Therefore, selecting 16:1 as the optimal concentration requires further\njustification, such as by comparing Coulombic efficiency, symmetric cell lifespan, or other\nelectrochemical performance metrics.\nResponse: Thank you for the comment. The suggestion has helped us enhance our manuscript.\nAccordingly, we measured the ionic conductivity at different temperatures using a conductivity\nmeter (Mettler Toledo Conductivity Meter, FE38-Standard) and included this information in\nSupplementary Fig. 6 and Supplementary Table 4. In the revised manuscript, the corresponding\nconductivity values have been changed on page 5, line 142:\n“Among the tested EA:H O ratios, 1:16 delivered the highest ionic conductivity at −50 °C\n2\n(24.92 mS cm⁻1).”\nP\nA\nG\nE\n\\\nIn addition, we compared the Coulombic efficiency of various electrolytes and the lifespan of\nsymmetric batteries. We have added this data as a new Supplementary Fig. 7. The corresponding\nexplanations were also added in the Supplementary Note 5 (Page 10 in Supplementary\ninformation):\n“Zn||Cu asymmetric cells were assembled to assess Zn plating/stripping reversibility across\nEA:H O ratios from 0:1 to 1:6. As shown in Supplementary Fig. 7a,b, the MDCE (EA:H O = 1:16)\n2 2\ndelivered average Coulombic efficiencies of 98.2% at 25 °C and 97.14% at -50 °C, outperforming\nthe other formulations. In contrast, the HDCE (0:1) exhibited an operational life of only ~100\ncycles with an average Coulombic efficiency of 43.34% and failed to operate at -50 °C.\nSymmetric Zn||Zn cells were then used to probe cycling stability (Supplementary Fig. 7c-e). Most\nEA-containing electrolytes showed markedly enhanced stability relative to 0:1 and 1:20 ratios.\nNotably, in the MDCE cell the overpotential decreased from 92.6 mV (0-10 h) to 21.1 mV (90-\n100 h), consistent with the formation of a stable, ion-conductive SEI during initial cycles. These\nresults identify EA:H O = 1:16 as the optimal composition for MDCE.”\n“To assess zinc plating/stripping reversibility across compositions, we assembled Zn||Cu and\nZn||Zn cells at 25 °C and −50 °C; both experiments identified EA:H₂O = 1:16 as the optimal ratio\nfor low‑temperature operation (Supplementary Fig. 7; Supplementary Note 5).”\nP\nA\nG\nE\n\\\nFig. R5/Supplementary Fig. 6 The ionic conductivities of Zn(ClO ) in H O/EA solutions with\n4 2 2\ndifferent H O/EA molar ratios from 1:0 to 6:1 at different temperatures test by conductivity meter\n2\n(The conductivity of the electrolyte at 25, 0, −20, −40, and −50°C was measured using the Mettler\nToledo Conductivity Meter (FE38-Standard)).\nSupplementary Table 4. Temperature-dependent ionic conductivity of 2 M Zn(ClO ) in H O/EA\n4 2 2\nacross EA:H O ratios from 0:1 to 1:6, measured with a conductivity meter.\nEA:H O molar\n2\n25 °C 0 °C -20 °C -40 °C -50 °C\nratios\n0:1 105.01 94.78 84.28 -- --\n1:20 69.37 66.5 52.22 27.37 24.5\n1:16 64.47 54.11 45.71 25.69 24.92\n1:12 61.11 57.26 45.36 24.29 19.88\n1:8 48.37 44.52 34.23 16.45 14.49\n1:6 45.71 42.56 34.09 14.7 12.88\nP\nA\nG\nE\n\\\nFig. R6/Supplementary Fig. 7 a, b The CEs of Zn||Cu batteries at a current and capacity density\nof 1 mA cm−2 and 0.5 mAh cm−2 at 25 °C (a) and −50 °C (b). c-e, The cycling performance of\nZn||Zn batteries at a current and capacity density of 1 mA cm−2 and 1 mAh cm−2 (c), and local time\nvoltage curve within 0-10 hours (d) and 90-100 hours (e).\nComment 8. The authors provided WT-EXAFS data in Fig. 2c, effectively visualizing the EXAFS\nsignals. However, for rigorous validation and clarity of EXAFS fitting results, the authors are\nrequested to also provide the original Zn K-edge EXAFS oscillation data (in k-space) and the\nP\nA\nG\nE\n\\\ncorresponding fitting curves in R-space. Specifically, please include plots clearly showing the\nexperimental and fitted EXAFS data overlaid in both k-space and R-space, along with a detailed\nsummary table of structural parameters (coordination number, bond length, etc.). This information\nis essential for ensuring transparency, reproducibility, and robust evaluation of the EXAFS\nanalysis.\nResponse: Thank you for the comments. According to the reviewer’s suggestion, we added plots\nshowing the experimental and fitted EXAFS data overlaid in both k-space and R-space, along with\na detailed summary table of structural parameters (coordination number, bond length, etc.) of all\nsamples in the supplementary information. As shown in Supplementary Fig. 10 and Supplementary\nTable 5. Therefore, we have corrected the writing to make it more concise and accurate on page 7,\nline 185:\n“To resolve the Zn2+ solvation structure, we compared experimental XAFS with spectra simulated\nfrom DFT‑optimized models of [Zn(ClO )(H O) )]+ and [Zn(ClO )EA(H O) ]+ (Supplementary\n4 2 5 4 2 4\nFig.10a-f). In the EXAFS region, a pronounced peak at 1.65 Å corresponds to the Zn–O distance\nin [Zn(ClO )(H O) )]+ complex32. By contrast, MDCE exhibits enhanced multi‑shell features,\n4 2 5\nindicating a more heterogeneous coordination environment. Wavelet‑transform EXAFS further\nseparates these environments and reveals additional high‑k components in MDCE attributable to\nEA/ClO ⁻ coordination (Fig. 2c; Supplementary Fig. 10g).”\nA\nG\nE\n\\\nFig. R7/Supplementary Fig. 10 a, b Fitting of EXAFS spectra in R space (a) and k space (b) for\nZn foil. c, d, e, f Fitting of EXAFS spectra in R space and k space for HDCE based on DFT\noptimized molecular geometries of [Zn(ClO )(H O) ]+ (c, d) and MDCE based on DFT optimized\n4 2 5\nmolecular geometries of [Zn(ClO )EA(H O) ]+ (e, f). (g) Wavelet transforms for the Zn foil.\n4 2 4\nP\nA\nG\nE\n\\\nSupplementary Table 5. Curve-fitting parameters for Zn K-edge EXAFS of Zn(H O) (ClO )+ and\n2 5 4\nZn(H O) EA(ClO )+standards. Reported parameters: coordination number (C.N.), bond distance\n2 4 4\n(R), Debye–Waller factor (σ2), inner potential correction (ΔE ), and R-factor (goodness of fit). S 2\n0 0\nwas fixed at 0.85.\nσ2×103\nSample Path R (Å) C. N. ΔE (eV) R factor\n0\n(Å2)\nZn Zn-Zn 2.64 6 12.9 −0.038 0.009\nZn-O (H O) 2.13 5.1 6.4 −1.8 0.007\nZn(H 2 O) 5 (ClO 4 )+) Zn-O (ClO 4 ) 2.09 0.9\nZn-O (H O) 2.13 3.4 9.6 −2.6 0.013\nZn-O (ClO ) 2.09 1.5\n2 4 4\nZn-O (EA) 2.15 1.1\nComment 9. The description of the in situ FTIR difference spectra is quite confusing. First, it is\nunclear whether the spectra shown in Fig. 3c and Fig. S14b correspond to HDCE or MDCE, the\nmanuscript does not clarify this, leaving the reader uncertain. Second, the definitions of A(t ),\nx\nA(t ), and A(t ) are not provided in the main text, and to the best of my knowledge, these are not\n0 z\nstandard or widely recognized terms. In Fig. S14c and d, there is also a clear inconsistency between\nthe y-axis label and the figure captions. The axis label reads “A(t) - A(t )”, while the caption refers\n0\nto “A(t ) - A(t )”, which are not equivalent concepts based on the definitions implied in the\nz 0\nmanuscript. This inconsistency is confusing and should be corrected to ensure clarity. Moreover,\naccording to the authors, A(t ) represents the first spectrum collected after applying current and\nz\nshould therefore be time-independent. However, the spectral changes shown in the figure appear\nto vary over time, which contradicts that definition. Additionally, the observed enhancement of\nW-H relative absorbance for MDCE under biased potential deviates from previous reports in the\nliterature (Yu, Xiaoyu, et al. “Unlocking dynamic solvation chemistry and hydrogen evolution\nmechanism in aqueous zinc batteries.” Journal of the American Chemical Society 146.25 (2024):\n17103-17113.). The authors should provide a clear explanation for this discrepancy. Lastly,\nimportant experimental details are missing, such as the duration of the charge/discharge process,\nP\nA\nG\nE\n\\\nspectral acquisition resolution, and the number of scans. Providing this information is necessary\nfor readers to fully understand and evaluate the experimental setup and results.\nResponse: Thank you for your comments. We will address them below one by one.\nFirst, we reorganized the content and logic to make the connection between images and samples\nclearer. Specifically, we added the following text to the manuscript on page 10, line 269:\n“In MDCE, significant spectral changes occurred during the initial discharge/charge cycle,\nincluding bands related to decomposition and species formation for interfacial H O (νOH), EA\n2\n(νCH , νOCH ), and ClO ⁻ (νClO) (Fig. 3e). These changes illustrate the dynamics of weak\n3 2 4\nhydrogen bonds (W-H) and strong hydrogen bonds (S-H) of interfacial water, that is water at the\ninterface between the electrode surface and the electrolyte solution (Fig. 3f). In HDCE, the increase\nin W-H signal during plating indicates higher water activity and stronger HER (Supplementary\nFig. 16b).”\nSecond, we redefined the acquired spectral information and supplemented the experimental details\nin the method part on page 22, line 585:\n“Background spectra were first recorded in air, followed by A(t ) after electrolyte filling (open\n0\ncircuit) and time‑resolved spectra A(t ), A(t ) under current. Relative absorbance was defined as\nx y\nΔA(t )=A(t )−A(t ) (abrupt change under bias) and ΔA(t )=A(t )−A(t ) (subsequent gradual\nx x 0 y y x\nevolution) (Fig. R8).”\nP\nA\nG\nE\n\\\nFig. R8 The dynamic interface evolution during Zn plating/stripping. The process can be\ndecoupled into two stages, abrupt change of interface reconstruction (expressed as ΔA(t ) = A(t )\nx x\n− A(t )) and gradual change of (de)solvation evolution (expressed as ΔA(t ) = A(t ) − A(t ))[56].\n0 y y x\n[56] Yu, X. et al. Deciphering multi-dimensional interfacial mechanisms via organic cosolvent\nengineering for sustainable zinc metal batteries. Nat. Commun. 16, 3820 (2025).\nDuring infrared spectroscopy testing, the high humidity in the laboratory significantly influenced\nthe infrared absorption peaks of water molecules, causing a continuous rise in water peak intensity.\nTo reduce environmental humidity's effects on the results, controlled humidity chamber was\nemployed to lower humidity below 10%. Under these controlled conditions, it was observed that\nduring charging, the infrared absorption intensity of the weak hydrogen bond (W-H) increased,\nwhereas during discharging, this peak's intensity noticeably decreased (Fig. R9).\nFig. R9 The A(t )-A(t ) relative absorbance evolution during charging (left, red) and discharging\ny x\n(right, blue), which depicts the evolution of W-H stretching vibration of interfacial H O in MDCE.\nmanuscript accordingly on page 22, line 579:\n“In situ electrochemical FTIR was performed on a Vertex 70 spectrometer (Bruker) coupled to\na CHI760E workstation (Shanghai Chinstrument Co., Ltd.) and equipped with a\nliquid‑nitrogen‑cooled broadband MCT detector (Beijing bjscistar Co., Ltd.) at room temperature.\nSpectra were collected in absorbance mode [A = −log(I/I₀)] at a fixed 60° incidence and 4 cm⁻1\nresolution. A custom spectro‑electrochemical cell used a Cu‑mesh working electrode and Zn foil\nP\nA\nG\nE\n\\\nas both counter and reference; Zn was plated at 0.5 mA for 900 s and stripped at 0.5 mA to a 0.5 V\ncut‑off. Background spectra were first recorded in air, followed by A(t ) after electrolyte filling\n0\n(open circuit) and time‑resolved spectra A(t ), A(t ) under current. Relative absorbance was\nx y\ndefined as ΔA(t )=A(t )−A(t ) (abrupt change under bias) and ΔA(t )=A(t )−A(t ) (subsequent\nx x 0 y y x\ngradual evolution).”\nComment 10. Why does MDCE exhibit a significantly higher nucleation overpotential difference\nat −40°C compared to −50°C in Figure 4d?\nResponse: Thank you for flagging this discrepancy. We repeated the nucleation measurements at\n−40 °C and −50 °C using a stricter protocol (identical cell history, extended thermal equilibration,\nand standardized pre-conditioning) across independent cells. The corrected data show that the\nnucleation overpotential in MDCE is smaller in magnitude at −40 °C than at −50 °C, as expected\nfrom transport-limited nucleation kinetics at lower temperature. We have replaced Fig. 4d with the\ncorrected results (and included the corresponding response figure), added representative voltage–\ntime traces with error bars, and updated the Methods to detail the revised protocol so that the\noutcome is reproducible.\nFig. R10 The nucleation overpotential (η) and the plateau overpotential (Δη) at 0.5 mA cm−2 in\nMDCE at −40°C.\nP\nA\nG\nE\n\\\nFig. R11/Figure 4d The nucleation overpotential (η) and the plateau overpotential (Δη) at\n0.5 mA cm−2 in MDCE at 25°C, −40°C, and −50°C.\nP\nA\nG\nE\n\\\nReviewer #2\nComments (major strengths):\nThe paper reports a dielectric constant modulation strategy to improve the cryogenic performance\nof aqueous Zn metal batteries. By introducing a low-ε co-solvent (ethyl acetate, EA) into a\nZn(ClO ) electrolyte, the authors successfully disrupt the hydrogen-bonding network of water,\n4 2\nregulate Zn²⁺ solvation structure, and enable stable Zn plating/stripping under extreme cold\nconditions (down to –50 °C).\nThe study is thorough, well-executed, and presented. The approach is novel, and the conclusions\nare convincingly supported by complementary experimental and simulation data. The\nelectrochemical results are outstanding, especially the symmetric Zn||Zn cycling for over 4,000 h\nat −50 °C and 10,000 cycles in PANI||Zn full cells.\nThe manuscript is generally well-organized and experimentally comprehensive. However, the\nfollowing concerns should be addressed before publication.\nResponse: Thank you very much for your positive comments and words of praise.\nComment 1. Mechanistic ambiguity in “altering dielectric response” The statement\n“Zn(ClO₄)₂·6H₂O enhances the miscibility of EA and water by altering the solution’s overall\ndielectric response” (page 5, paragraph 1) lacks sufficient mechanistic explanation. While it is\nsuggested that the salt improves solvent compatibility, the phrase “altering dielectric response”\nremains vague. It would be helpful to clarify how Zn(ClO₄)₂·6H₂O influences the dielectric\nenvironment — for instance, by coordinating with water to reduce its effective polarity, by\nscreening unfavorable interactions between EA and H₂O, or by modifying hydrogen bonding\nnetworks. A brief discussion or reference to a specific interaction model (e.g., salt-mediated dipole\nalignment or suppression of phase separation) would strengthen this point and aid the reader's\nunderstanding.\nP\nA\nG\nE\n\\\nResponse: Thank you for the helpful comment. We have replaced the phrase “altering dielectric\nresponse” with a mechanism-based explanation. In brief, Zn(ClO ) (H O) modulates the local\n4 2 2 6\npermittivity and improves EA–H O miscibility through two coupled effects. First, Zn2+\n2\ncoordinates water to form [Zn(H O) ]2+, which decreases the population of strongly hydrogen-\n2 6\nbonded, highly polar free water and its orientational polarization. Second, the weakly coordinating,\nchaotropic ClO – anion disrupts the water hydrogen-bond network and screens unfavorable EA–\n2\ndipole correlations, and a reduced free-energy penalty for mixing. We now state this mechanism\nin the expanded Supplementary Discussion 2 with a concise interaction-model rationale (ion–ion,\nion–dipole, dipole–dipole) framed in the Kirkwood–Frohlich picture of dielectric response,\ntogether with supporting literature (Hou et al., Adv. Energy Mater. 2023; Chen and Zhang, Acc.\nChem. Res. 2020).\nThe article has been revised in Supplementary Information on page 2:\n“To rationalize the structure–activity relationship of the electrolyte, we introduce classical models\nfor the principal intermolecular interactions (ion–ion, ion–dipole, and dipole–dipole). Their\norientation-averaged potential energies are given by Eqs. (7)–(9) (SI units suppressed for clarity)4,5.\n1 z z e2 (7)\nU =- \n1 2\nion-ion 4πε r\n1 zeμcosθ (8)\nU =- \nion-dipole 4πε r2\n1 2μ2μ2 (9)\nU =- 1 2\ndipole-dipole (4πε)23k\nB\nTr6\nThe dielectric constant ε quantifies the response to an applied field and arises microscopically from\ndipole orientation. In a simple mean-field picture, it scales with the effective molecular dipole\nmoment and the number density of dipoles:\nμ2μ2 (10)\nε∝ 1 2\nk T\nB\nHere, U denotes the interaction energy, ε is the dielectric constant, ze is the charge of the ion, µ is\nthe dipole moment of the dipole, r is the distance between ions or the centers of dipoles, θ is the\ndipole angle relative to the line joining the ion and the center of the dipole, k is the Boltzmann\nB\nP\nA\nG\nE\n\\\nconstant, and T is the absolute temperature.\nElectrolytes modify dielectric behavior through coordination and network effects. For example,\nZn2+ forms [Zn(H O) ]2+ complexes, binding water and reducing the number density of free water\n2 6\ndipoles, thereby lowering their dielectric contribution. At 25 °C, pure water has ε = 78.466, whereas\na Zn(ClO ) aqueous solution exhibits epsilon ~ 45. In mixed EA/H O electrolytes, EA engages in\n4 2 2\ndipole–dipole interactions with water, and the chaotropic anion ClO - disrupts the water hydrogen-\n4\nbond network. Both effects increase orientational disorder, reduce the effective vector sum of\nwater dipoles, and diminish the bulk dielectric response. The resulting local dielectric shielding,\nintegrated over the liquid, decreases the overall ε. (Page 2, Supplementary information)\n[4] R. Hou, S. Guo, H. Zhou, Atomic insights into advances and issues in low-temperature\nelectrolytes. Adv. Energy Mater. 2023, 13, 2300053.\n[5] Chen, X. & Zhang, Q. Atomic insights into the fundamental interactions in lithium battery\nelectrolytes. Acc. Chem. Res. 2020, 53, 1992-2002.\nComment 2. Spectral interpretation lacks quantification or deconvolution\nThe FTIR and Raman shifts (e.g., C–O–C from 1046 to 1040 cm⁻¹, O–Cl–O from 1079 to 1073\ncm⁻¹) are invoked as evidence of solvation changes, but the authors present only qualitative\ndiscussion. There is no spectral deconvolution or integration data (e.g., peak area ratios) to quantify\nsolvation shell reorganization or binding strength. Especially for the ClO₄⁻ shift, which is subtle,\na clearer attribution is needed—could the observed change be due to simple dilution, not\ncoordination? A deeper spectroscopic analysis would strengthen the claims.\nResponse: Thank you very much for your valuable suggestions. We conducted quantitative\ndeconvolution of infrared band to investigate and quantify the solvation shell reorganization\nfocusing on the C–O–C from 1046 to 1040 cm⁻1 and O–Cl–O from 1079 to 1073 cm⁻1. CIP band\narea increases from 16% to 21.9% on adding EA, SSIP decreases from 49.5% to 43.3%\n(Supplementary Fig. 7d). Spectra were baseline-corrected and fitted with Voigt profiles using\nshared-width constraints; peak positions, FWHM, and integrated areas were extracted.\nFurthermore, the following was added to the manuscript in order to explain this analysis in main\ntext on page 6, line 162:\nP\nA\nG\nE\n\\\n“To elucidate ion pairing in the MDCE, we probed Zn2+–ClO –coordination by FTIR\n4\n(Supplementary Fig. 8c,d). The results showed peaks assigned to solvent-separated ion pairs (SSIP,\nnoncoordinated ClO –, 1076 cm–1), contact ion pairs, (CIPs, Zn-(ClO –) , 1047 cm–1), indicating\n4 4 n≥1\nstrong adsorption for ClO –, free anion (1114 cm–1) and free EA (1037 cm–1). Upon EA addition,\n4\nthe SSIP and free‑anion bands diminish while the CIP band grows, indicating strengthened cation–\nanion association with a fraction of EA remaining uncoordinated (Supplementary Fig. 8c,d).”\nFig. R12/Supplementary Fig. 8d Deconvolution result of FTIR spectra for 2 M Zn(ClO ) H O\n4 2 2\n(EA: H O = 0:1) and 2 M Zn(ClO ) EA:H O (EA:H O = 1:16) electrolytes at 1,200–1,000 cm−1.\n2 4 2 2 2\nComment 3. The authors state that FTIR/Raman results (e.g., shifts in O–H vibrations)\n“promote the conversion of weak and medium-strength hydrogen bond networks to strong hydrogen\nbond networks,” which contradicts the usual interpretation that the addition of organic co-solvents\nweakens water’s H-bond network. The use of “strong H-bond network” here needs clarification:\ndoes this imply localized but fewer, stronger H-bonds or just a spectral artifact? Otherwise, the\ninterpretation appears inconsistent to disrupt water structure to improve low-temperature\nperformance.\nResponse: Thank you for bringing the problem to our attention. As you mentioned, we have\ncorrected our inaccurate descriptions based on the results of Raman and infrared analysis. Both\nFTIR and Raman indicate progressive weakening of the water H-bond network with EA:\nP\nA\nG\nE\n\\\ncluster/bulk-water O–H bands shift to higher wavenumber while the non–H-bonded component\ngrows (to ~60% at EA:H O=1:6), consistent with reduced ordering at low temperature (Main Fig.\n2\n2a; Supplementary Fig. 9)\nWe have corrected this statement in the main text on page 7, line 176:\n“These results suggest an increased disruption of the hydrogen-bonding network of water and\npromote the conversion of weak and strong-strength hydrogen bond networks to non-hydrogen\nbond networks in water, which potentially inhibits the ordered arrangement of water molecules\ninduced by a decrease in temperature.”\nComment 4. The potentiometric measurement of solvation energy (OCP between asymmetric\nelectrolytes) is elegant, but its accuracy depends on minimizing liquid junction potential and\nassuming identical Zn activity at both electrodes. While Supplementary Discussion 1 outlines these\nassumptions, no control (e.g., symmetric electrolyte test, or multiple Zn salts) is shown to validate\nthe result.\nResponse: Thank you for the thorough review. We supplemented the experiments and conducted\nfurther analysis based on your suggestions, which are included in Supplementary Fig. 13 and\nSupplementary Table 6. We supplemented with symmetric electrolyte testing to verify that the Zn\nactivity of the two electrodes is the same under minimized liquid junction potential. When the\nreference electrolyte and the test electrolyte are the same, the potential remains around 0 V. The\nmeasured potential ensures that it only reflects the electrode reactions themselves, eliminating the\nimpact of additional potentials at the electrolyte interface. In addition, we supplemented with\ndifferent reference electrolytes to verify the more negative free energy of the MDCE.\nP\nA\nG\nE\n\\\nFig. R13/Supplementary Fig. 13 Potentiometric measurement of relative Zn2+ solvation energy.\n(a) H-cell setup: two Zn electrodes, each immersed in a different electrolyte, a reference electrolyte\n(EL ) and a test electrolyte (EL ). The open-circuit cell voltage E_cell reports the difference in\nref test\nZn2+ solvation free energy between EL and EL (ΔG = -nFE _cell; see Supplementary\ntest ref\nDiscussion S1). (b) Representative voltage trace for HDCE vs an aqueous reference electrolyte.\n(c–h) HDCE (c,e,g) and MDCE (d,f,h) measured against 2 M ZnSO , 2 M Zn(CF SO ) , and\n4(aq) 3 3 2(aq)\n2(aq)\nmixing-induced drift; over 60 min, drift is approximately 1–2 mV.\nSupplementary Table 6. Summary of solvation energy for each electrolyte.\nP\nA\nG\nE\n\\\nSolvation Energy Solvation Energy\nSolvent Reference electrolyte Voltage (V)\n(eV) (kJ mol-1)\nHDCE HDCE ~0 ~0 ~0\nHDCE 2 M ZnSO H O 0.0529 ± 0.00017 -0.1058 ± 0.00034 -10.1568 ± 0.0326\n4 2\nMDCE 2 M ZnSO H O 0.0860 ± 0.00110 -0.1720 ± 0.00220 -16.5954 ± 0.2123\n4 2\nHDCE 2 M Zn(CF SO ) H O 0.0401 ± 0.00007 -0.0774 ± 0.00014 -7.7381 ± 0.0135\n3 3 2 2\nMDCE 2 M Zn(CF SO ) H O 0.0718 ± 0.00012 -0.1386 ± 0.00232 -13.8552 ± 0.0232\n3 3 2 2\nHDCE 1 M Zn(OAc) H O 0.0960 ± 0.00017 -0.1853 ± 0.00033 -18.5251 ± 0.0328\n2 2\nMDCE 1 M Zn(OAc) H O 0.1301 ± 0.00009 -0.2511 ± 0.00174 -25.1054 ± 0.0174\n2 2\nComment 5. The authors propose that stronger Zn2⁺–ClO ⁻ pairing (enabled by lower ε) leads to\n4 4\nstandard potentials. No electrochemical or chemical evidence (e.g., cyclic voltammetry, ion\nchromatography) is provided to confirm ClO ⁻ reduction or its conversion to SEI-forming species.\nResponse: Thank you for this valuable suggestion. We measured the onset of anion reduction by\nlinear sweep voltammetry (LSV) and incorporated the detailed description below verbatim in the\nmanuscript on page 10, line 280:\n“Consistent with the FTIR transients pointing to early interphase formation, linear sweep\nvoltammetry differentiates the two electrolytes (Supplementary Fig. 17). The full-scale LSV\ncurves in Supplementary Fig. 17a show the behavior of both electrolytes. As can be seen from the\nenlarged curves in Supplementary Fig. 17b, the first reduction scan in the HDCE reveals HER\nstarting at −0.86 V, where zinc deposition initiates sharply. In contrast, the first reduction scan in\nthe MDCE includes a distinctive reduction peak starting at −0.71 V, which reaches its maximum\nat −0.76 V (Supplementary Fig. 17c). The reduction peak can only be observed in the first cycle\nand disappeared in the second cycle. This phenomenon is a typical signature of SEI formation44.”\nP\nA\nG\nE\n\\\nFig. R14/Supplementary Fig. 17 (a) Full-scale LSV and magnified views for (b) HDCE and (c)\nMDCE, recorded with Cu foil (working), Zn metal (counter), and Ag/AgCl (reference).\n[44] Chen, T., Wang, Y., Li, X., Yuan, T., Pang, Y. & Zheng, S. Enabling solid-electrolyte interphase\nformation prior to water reduction in aqueous zinc batteries by mild protic chemistry. Angew.\nChem. Inter. Edit. 64, e202424642 (2025).\nFurthermore, ex situ mass spectrometry of the SEI shows a strong Cl signal with negligible ClO −,\n4\nindicating consumption of perchlorate to form Cl-containing SEI species (Fig. R15).\nP\nA\nG\nE\n\\\nFig. R15/Supplementary Fig. 29 a,b, Mass spectrometry of cycled Zn anode in HDCE (a) and\nMDCE (b).\nComment 6. While the manuscript convincingly demonstrates the formation of a robust SEI layer\nin the MDCE system through cryo-TEM, XPS, and TOF-SIMS analysis, the underlying\nmechanism by which EA promotes SEI formation remains insufficiently explained. Given the\nknown challenges in forming SEI in aqueous systems due to competitive HER, it is unclear how\nEA reduction products contribute to SEI nucleation and growth. Additionally, the possible\nsynergistic decomposition of EA and ClO ⁻ is not well discussed.\nany evidence of EA-derived species (e.g., acetate) directly contributing to SEI structure?\nIncluding reaction pathway analysis or DFT simulations would significantly strengthen the\nmechanistic claims.\nResponse: Thank you very much for your valuable suggestions. Density functional theory\ncalculations indicate that the downstream reactions resulting from the decomposition of EA are\nthe cause of the formation of organic components in the SEI (Fig. R16/Supplementary Fig. 30).\nWe calculated the reduction pathway for the decomposition of EA using DFT. And we have added\nthe corresponding content in the main text on page 17, line 458:\n“DFT indicates that EA undergoes single‑electron reduction; the resulting radical/carbanion then\ndeprotonates to an alkene, while acetate (CH COO⁻) coordinates Zn2⁺ and contributes to forming\n3\nan organic‑rich SEI. The computed reaction free energy for overall decomposition reaction EA\n→CH COO⁻ + CH CH is −1.12 eV at 298 K and −1.06 eV at 233 K, indicating a\n3 2 3\nthermodynamically favorable pathway that persists at low temperature (Supplementary Fig. 30)”\nP\nA\nG\nE\n\\\nFig. R16/Supplementary Fig. 30 Calculated reaction energetics (Delta G) for the electron-\ninduced reduction and decomposition of EA.\nComputational details are provided in the Methods on page 24, line 656:\nMolecular‑orbital energies and EA‑reduction energetics were computed with Gaussian 16[69] at the\nM05‑2X/6‑31++G(d,p) level, using a polarizable‑continuum (PCM) water model for solvation[70].\n[69] Gaussian 16 rev. C.01 (Wallingford, CT, 2016).\n[70] Scalmani, G. & Frisch, M. J. Continuous surface charge polarizable continuum models of\nsolvation. I. General formalism. J. Chem. Phys. 132 (2010).\nThe distribution results of Cl⁻ and CH COO⁻ in TOF-SIMS indicate that both ClO ⁻ and EA have\n3 4\nbeen decomposed. The presence of the C-Cl peak in XPS suggests a possible synergistic\ndecomposition of EA and ClO ⁻. According to your suggestion, we further discussed the promoting\n4\neffect of EA addition on anion decomposition. Following this, we employed density functional\ntheory (DFT) calculations to examine the highest occupied molecular orbital (HOMO) and lowest\nunoccupied molecular orbital (LUMO) energy levels of the primary Zn²⁺ solvation structures and\nprovide the corresponding explanation in the main text on page 8, line 221:\n“DFT frontier‑orbital analysis shows that EA‑containing CIP complexes have a smaller HOMO–\nLUMO gap than CIPs without EA, implying greater electron‑transfer propensity and facilitating\nanion‑derived SEI formation (Supplementary Fig. 11e) [36].”\nP\nA\nG\nE\n\\\n[36] Hei, P., Sai, Y., Yu, L., Lin, Y., Li, B., Hu, G., Wu, W., Wang, J., Sun, X., Liu, X.-X. & Song,\nY. Cosolvent electrolyte design for high-voltage aqueous zinc–sulfur batteries. J. Am. Chem. Soc.\n147, 27802-27811 (2025).\nTogether, these results link EA reduction to acetate formation and to an increased propensity for\nanion-derived SEI, providing a coherent mechanism for EA-promoted SEI nucleation and growth\nin MDCE.\nFig. R17/Supplementary Fig. 11e HOMO/LUMO energy levels and molecular orbitals for\nsolvation structures in HDCE and MDCE; structures are defined by the local fractions of\nEA/H O/ClO ⁻.\n2 4\nComment 7. The authors attribute Zn corrosion in the HDCE system to Cl⁻ derived from ClO ⁻\n4\ndecomposition. However, ClO ⁻ is generally considered electrochemically inert under standard\n4\naqueous conditions, and the manuscript lacks direct evidence for its decomposition. This assertion\nis critical, as it underpins the comparison between HDCE and MDCE degradation behavior.\nResponse: We agree that additional support research was needed. We now provide converging\nevidence that perchlorate can be reduced under the strongly cathodic conditions at a Zn surface in\nP\nA\nG\nE\n\\\nHDCE. First, TOF-SIMS of cycled HDCE anodes shows a pronounced Cl− signal within the SEI\nwith negligible ClO −, whereas MDCE shows much weaker Cl− (Fig. R18/Supplementary Fig. 29).\ncontaining interphases. These observations indicate that the chlorine found in the SEI does not\noriginate from trapped undecomposed ClO −.\nour assignment that ClO − undergoes reduction to Cl− in HDCE and contributes to Zn corrosion.\nperchlorate becomes reactive\nSecond, we added a brief literature context on perchlorate reactivity under reducing conditions.\nLi et al. demonstrated that ClO − participates in the primary solvation sheath of Zn2+, facilitating\n4\nthe in-situ formation of a Zn (OH) Cl ·H O-rich solid electrolyte interphase (SEI) to suppress the\n5 8 2 2\ncorrosive effect of ClO −. (Li, YS., Geng, LS., Zhang, BM. et al. Concentrated perchlorate-based\n4\nelectrolyte facilitates Zn anode-compatible in-situ solid electrolyte interphase. Rare Met. 44, 950–\n960 (2025).) Yang et al. also demonstrated the decomposition of ClO −: “Then, these cells were\n4\ndisassembled and the over production of undesired insulating products (i.e., Zn ClO (OH) and\n4 4 7\nZn (OH) Cl ) were detected from X-ray diffraction (XRD). (Yang et al., Joule 4, 1557–1574\n5 8 2\n(2020)).\nFig. R18/Supplementary Fig. 29 a,b, Mass spectrometry of cycled Zn anode in HDCE (a) and\nMDCE (b).\nP\nA\nG\nE\n\\\nReviewer #3:\nIn this work, the author proposed a strategy for low-temperature ZMBs based on tailoring the Zn2+\nsolvation environment by engineering the dielectric constant. Specifically, according to\nintroduction of EA into Zn(ClO ) electrolyte to weaken water’s hydrogen-bond network and\n4 2\nincrease cation-anion pairing. This strategy accelerates Zn2⁺ transport and desolvation, promotes\nthe formation of a protective solid electrolyte interphase rich in organic and inorganic components,\nand inhibits parasitic hydrogen evolution. Although it achieves excellent electrochemical\nperformance for the cryogenic aqueous zinc batteries, many issues should be addressed before\npublication.\nResponse: Thank you very much for your positive comments.\nComment 1. It is mentioned that Zn(ClO ) ·6H O enhances the miscibility of EA and water by\n4 2 2\naltering the solution's overall dielectric response. In the method section, the calculation method\nand equation of dielectric constants of hybrid salt and solvent (such as Zn(ClO ) /EA, Zn(ClO )\n4 2 4 2\nEA:H O) are not clear. Please supplemented the section.\ncorresponding content to the Methods section of the manuscript on page 24, line 636:\n“Dielectric constants of the solvent and electrolytes were calculated using the built-in dielectric\nanalysis module in GROMACS, based on the final 20 ns of the production trajectory[57]. The OPC\nwater model was employed for the simulations known for its high accuracy in dielectric property\nsimulations of aqueous systems[66]. The dipole moment of the system was output every 1000 steps\nduring NVT simulations and used for calculation of dielectric constant.” The dielectric constant of\nliquid electrolytes is calculated according to formula 5[28,29]:\n< 𝑀2 > −< 𝑀 >2 (5)\nε = 1+ \n3𝜀 𝑉𝑘 𝑇\n0 𝐵\nP\nA\nG\nE\n\\\nwhere 𝜀 is the dielectric constant of specific system, 𝜀 the dielectric constant of vacuum, 𝑉 the\nvolume of the simulation box, 𝑘 the Boltzmann constant, 𝑇 the temperature. 𝑀=∑𝜇⃗ =∑𝑞𝑟⃗ is the\ntotal dipole moment of the system where 𝑞 and 𝑟⃗ are the charge and position of atoms, respectively.\nTotal dipole moment of the system (𝑀 in formula 5) and square of it (𝑀) were both averaged over\n5000 frames, which means the dielectric constant is calculated from 5 ns to the end of NVT\nproduction run.\n[28] Jorgensen, W. L., Maxwell, D. S. & Tirado-Rives, J. Development and testing of the opls all-\natom force field on conformational energetics and properties of organic liquids. J. Am. Chem. Soc.\n118, 11225-11236 (1996).\n[29] Neumann, M. & Steinhauser, O. Computer simulation and the dielectric constant of\npolarizable polar systems. Chem. Phys. Lett. 106, 563-569 (1984).\n[57] Abraham, M. J., Murtola, T., Schulz, R., Páll, S., Smith, J. C., Hess, B. & Lindahl, E.\nGromacs: High performance molecular simulations through multi-level parallelism from laptops\nto supercomputers. SoftwareX 1-2, 19-25 (2015).\n[66] Kadaoluwa Pathirannahalage, S. P., Meftahi, N., Elbourne, A., Weiss, A. C. G., McConville,\nC. F., Padua, A., Winkler, D. A., Costa Gomes, M., Greaves, T. L., Le, T. C., Besford, Q. A. &\nChristofferson, A. J. Systematic comparison of the structural and dynamic properties of commonly\nused water models for molecular dynamics simulations. Journal of Chemical Information and\nModeling 61, 4521-4536 (2021).\nAnd we explained how Zn(ClO ) ·6H O affects the dielectric environment to enhance solvent\n4 2 2\ncompatibility and supplemented this in Supplementary Discussion 2:\n“To rationalize the structure–activity relationship of the electrolyte, we introduce classical models\nfor the principal intermolecular interactions (ion–ion, ion–dipole, and dipole–dipole). Their\norientation-averaged potential energies are given by Eqs. (7)–(9) (SI units suppressed for\nclarity)4,5.\n1 z z e2 (7)\n1 2\nU =- \nion-ion 4πε r\nP\nA\nG\nE\n\\\n1 zeμcosθ (8)\nU =- \nion-dipole 4πε r2\n1 2μ2μ2 (9)\nU =- 1 2\ndipole-dipole (4πε)23k Tr6\nB\nThe dielectric constant epsilon quantifies the response to an applied field and arises\nmicroscopically from dipole orientation. In a simple mean-field picture, it scales with the effective\nmolecular dipole moment and the number density of dipoles:\nμ2μ2 (10)\nε∝ 1 2\nk T\nB\nHere, U denotes the interaction energy, ε is the dielectric constant, ze is the charge of the ion, µ is\nthe dipole moment of the dipole, r is the distance between ions or the centers of dipoles, θ is the\ndipole angle relative to the line joining the ion and the center of the dipole, k is the Boltzmann\nB\nconstant, and T is the absolute temperature.\nElectrolytes modify dielectric behavior through coordination and network effects. For example,\nZn2+ forms [Zn(H O) ]2+ complexes, binding water and reducing the number density of free water\n2 6\ndipoles, thereby lowering their dielectric contribution. At 25 °C, pure water has ε = 78.466, whereas\na Zn(ClO ) aqueous solution exhibits epsilon ~ 45. In mixed EA/H O electrolytes, EA engages in\n4 2 2\ndipole–dipole interactions with water, and the chaotropic anion ClO - disrupts the water hydrogen-\n4\nbond network. Both effects increase orientational disorder, reduce the effective vector sum of water\ndipoles, and diminish the bulk dielectric response. The resulting local dielectric shielding,\nintegrated over the liquid, decreases the overall ε. (Page 2, Supplementary information)\n[4] R. Hou, S. Guo, H. Zhou, Atomic insights into advances and issues in low-temperature\nelectrolytes. Adv. Energy Mater. 2023, 13, 2300053.\n[5] Chen, X. & Zhang, Q. Atomic insights into the fundamental interactions in lithium battery\nelectrolytes. Acc. Chem. Res. 2020, 53, 1992-2002.\nP\nA\nG\nE\n\\\nComment 2. The author proposed that EA is an ideal co-solvent because it has a low dielectric\nconstant, low viscosity, and low freezing point. But as can be seen from Table S1, MeOH and THF\nalso have similar characteristics. Besides, other carboxylic esters, such as methyl acetate and\nmethyl formate also have similar characteristics. Why EA was chosen as the co-solvent, the\nexplanation is not sufficient.\nResponse: Thank you for the comment. We investigated the electrochemical performance of\nMeOH, THF, and methyl formate (MF) as co-solvents at different temperatures (Fig. R18). Our\nexperimental results (Table S1) demonstrate that the EA-containing electrolyte exhibits\nsignificantly higher Coulombic Efficiency (CE) at both room and low temperatures compared to\nMeOH (ɛ = 32.7), THF (ɛ = 7.4) and MF (ɛ = 8.5), despite their similar physical properties. During\nthe 200 cycles of zinc plating/stripping process, the CE of EA at 25 °C and -50 °C is 97.6% and\n97.3% respectively. And the overpotential at -50 °C being 373 mV. This suggests that EA's\nadvantage is not solely determined by dielectric constant or viscosity, but rather by its unique\ninterfacial compatibility with electrodes and Zn salts.\nMeOH has a relatively high dielectric constant, and when mixed with water at a molar ratio of\n1:16, the resulting electrolyte still maintains a high dielectric constant. And MF demonstrated poor\nelectrochemical stability at room temperature, attributable to the intrinsic properties of the solvents.\nWhile MF and THF can serve as additives in organic low-temperature electrolytes, MF undergoes\nhydrolysis in aqueous electrolytes. EA exhibits superior stability compared to MF. The stability of\nester groups is influenced by inductive and conjugation effects: (1) The carbonyl carbon in MF is\nmore positively charged due to the weaker electron-withdrawing ability of -H compared to -CH ,\n3\nrendering it more susceptible to nucleophilic attack. (2) In EA, the methyl group disperses the\npositive charge of the carbonyl carbon through hyperconjugation, thereby enhancing its stability.\n(March, J. (1992). Advanced Organic Chemistry (4th ed.). Wiley.)\nP\nA\nG\nE\n\\\nFig. R18 a, b The CEs of Zn||Cu batteries at a current density of 1 mA cm−2 and 0.5 mAh cm−2 at\n25 °C (a) and −50 °C (b). c, The first discharge curve of Zn||Cu batteries at a current density of 1\nmA cm−2 and 0.5 mAh cm−2.\nComment 3. The results of Figure S1 show that ZnOTF also enables EA to form a homogeneous\nsolution with H O, so why not choose ZnOTF instead of ZnClO ?\n2 4\nResponse: Thank you for this important question. Figure S1 indeed shows that Zn(OTf)2\n(Zn(CF SO ) ) also yields a single-phase EA:H O solution; however, our salt choice is driven by\n3 3 2 2\nlow-temperature transport and interfacial behavior rather than miscibility alone. At identical\ncomposition (molar ratio of EA:H O =1:16), Zn(ClO ) delivers substantially higher ionic\n2 4 2\nconductivity: 64.47 mS cm−1 at 25 °C and 24.92 mS cm−1 at −50 °C, versus 41.32 and 10.56 mS\ncm−1 for Zn(OTf) , respectively (Table R1). This transport advantage translates to cell\n2\nperformance: in Zn||Cu tests (1 mA cm−2, 0.5 mAh cm−2), the Zn(ClO ) EA:H O electrolyte\n4 2 2\nexhibits higher Coulombic efficiency, more stable cycling at both temperatures, and a lower\nnucleation overpotential than the Zn(CF SO ) analogue (Fig. R19).\n3 3 2\nP\nA\nG\nE\n\\\nTable R1 Conductivity of Zn(CF SO ) EA:H O (molar ratio = 1:16) and Zn(ClO ) EA:H O\n3 3 2 2 4 2 2\n(molar ratio = 1:16) measured by the conductivity meter.\nZinc salt σ (mS cm−1)\n25 °C −50 °C\nZn(CF SO ) 41.32 10.56\n3 3 2\nZn(ClO ) 64.47 24.92\n4 2\nWe have added the corresponding content in Supplementary Note 1:\n“Additionally, Zn(CF SO ) enables miscibility of EA and water at a molar ratio of 1:16.\n3 3 2\nConductivity measurements indicate that Zn(CF SO ) in an EA:H O mixture (1:16 molar ratio)\n3 3 2 2\nachieves a conductivity of 41.32 mS cm⁻1 at 25 °C and 10.56 mS cm⁻1 at −50 °C. In contrast,\nunder the same conditions, Zn(ClO ) in EA:H O (1:16 molar ratio) exhibits higher\n4 2 2\nconductivities of 64.47 mS cm⁻1 at 25 °C and 24.92 mS cm⁻1 at −50 °C. Thus, the Zn(ClO ) -\n4 2\nbased electrolyte demonstrates superior conductivity at low temperatures, making it a promising\ncandidate for cryogenic applications.”\nP\nA\nG\nE\n\\\nFig. R19 a, b The CEs of Zn||Cu batteries in Zn(CF SO ) EA:H O (molar ratio = 1:16) and\n3 3 2 2\nZn(ClO ) EA:H O (molar ratio = 1:16) at a current density of 1 mA cm−2 and 0.5 mAh cm−2 at 25\n4 2 2\n°C (a) and −50 °C (b). c, The first discharge curve of Zn||Cu batteries at a current density of 1 mA\ncm−2 and 0.5 mAh cm−2.\nComment 4. In Line 165, it is stated that weak and medium hydrogen bonding networks evolve to\nstrong hydrogen bonding networks, but the statement is contrary to the cited reference.\nResponse: Thank you for your comment. As you mentioned, we have corrected our inaccurate\ndescriptions based on the results of Raman and infrared analysis. The Raman spectroscopy with\nhydrogen-bonded peaks in the 3,100 cm−1 to 3,700 cm−1 range (Supplementary Fig. 9a). And as\nshown in Supplementary Fig. 9b-h, the area of peaks representing non-hydrogen-bonded water\nincreases progressively to 60.3% with the rise in EA concentration, indicating the disruption of the\nhydrogen-bonding network of water. The FTIR result in Fig. 2a, we observed a shift to higher\nwavenumber in the peaks for cluster water (3335 cm⁻1) and bulk water (3229 cm−1), along with a\nshift to lower wavenumber in the peaks for isolated water (3552 cm−1). These results all confirm\nthat the addition of organic cosolvents weakens the hydrogen bond network of water.\nP\nA\nG\nE\n\\\nWe have revised the text on page 7, line 176:\n“These results suggest an increased disruption of the hydrogen-bonding network of water and\npromote the conversion of weak and strong-strength hydrogen bond networks to non-hydrogen\nbond networks in water, which potentially inhibits the ordered arrangement of water molecules\ninduced by a decrease in temperature.”\nComment 5. The results of Fig. 3a, b seems to contradict each other with Fig. 3a showing that\nMDCE has a greater solvation energy, while the opposite is true for Fig. 3b.\nResponse: Thank you for the comment. Figure 3a reports the apparent solvation free energy in the\nbulk electrolyte, inferred from the potentiometric transfer between compositions; this quantity\nreflects the thermodynamics of the solvated ion, including ion–ion correlations and activity\ncoefficients. In EA-rich MDCE, the lower permittivity promotes contact ion pairs/aggregates\naround Zn2+, which stabilizes Zn2+ in the bulk and renders the apparent solvation free energy more\nnegative.\nFigure 3b represents the desolvation energy, indicating the energy barrier for ion desolvation at\nthe interface under an electric field. Based on previous literature reports and our experimental\nresults, it can be concluded that solvents with weak salt dissociation ability can promote ion\ndesolvation. First, the weak ion–solvent interaction reduces the desolvation barrier of the solvated\ncharge-carrier ions and realizes a faster desolvation process. Second, weak ion–solvent interaction\nmeans that solvents have limited ability to dissociate salt crystals, which makes anions in the\nelectrolyte more likely to participate in the solvation structure of charge-carrier ions. The resulting\nstrong cation–anion interaction promotes the preferential reduction of anions to form an inorganic-\nrich SEI film, which accelerates the ion diffusion in the SEI film. (R. Hou, S. Guo, H. Zhou,\nAtomic Insights into Advances and Issues in Low-Temperature Electrolytes. Adv. Energy Mater.\n2023, 13, 2300053).”\nComment 6. Figure S13 only shows that it is easier to remove water molecules in the solvated\nconfiguration containing EA formed in the electrolyte after the addition of EA, while the ability of\nP\nA\nG\nE\n\\\nEA to replace some of the water molecules suggests that it has a stronger bonding force with Zn2+,\nwhich may cause a greater obstruction in the removal of EA molecules and ions. Is the statement\nin the abstract that the modified solvated structure accelerates Zn2+ desolvation not appropriate?\nResponse: Thank you for raising this point. Our previous results demonstrated that the\nintroduction of EA facilitates faster desolvation of water molecules in the Zn2⁺ solvation structure.\nFurthermore, we performed additional computational analyses to thoroughly assess the stability of\nthe modified solvation structure, offering deeper insights into how EA influences Zn²⁺ desolvation\nkinetics.\nFrom a dielectric perspective, introducing EA (ε ≈ 6.4) markedly diminishes the electrolyte's\ndielectric shielding capability. This reduction destabilizes the Zn²⁺ solvation shell, driving a\ntransition from a tightly bound, water-dominated configuration to a more loosely coordinated, EA-\nwater mixed structure. Although this electrolyte environment has the single-point binding strength\nof EA-Zn2⁺, it overall reduces the rigidity of the solvation shell, making Zn2⁺ more prone to partial\ndesorption. The energy levels (HOMO and LUMO) and molecular orbitals of solvation structures\nin HDCE and MDCE, compared to the solvation structures without EA, the structure with EA\nexhibits a narrower HOMO-LUMO energy gap, indicating that the structure containing EA is more\nsusceptible to facilitate the charge transfer reaction under an electric field (J. Am. Chem. Soc. 147,\n27802-27811 (2025)), facilitating the Zn²⁺ plating/stripping reactions (Supplementary Fig. 11e).\nFig. R20/Supplementary Fig. 11e HOMO/LUMO energy levels and molecular orbitals for\nsolvation structures in HDCE and MDCE; structures are defined by the local fractions of\nEA/H O/ClO -.\n2 4\nP\nA\nG\nE\n\\\nReviewer #4:\nThe work by Zhu and co-workers presented the performance and characterization of ZnClO\n4\naqueous electrolytes with or without EA additive. The electrolyte with EA, MDCE, showed\nsuperior performance compared to the aqueous counterpart, HDCE, especially at low temperature.\nThe study combined various experimental and simulation techniques and insightful details were\nprovided for the studied systems.\n1) The work emphasized “dielectric engineering” in the title. However, the meaning of which is\nnot very clear to me. The study included two electrolytes, one with relative high dielectric and the\nother one lower. Is the dielectric the decisive reason for the superior performance? If one uses a\ndifferent additive to prepare an electrolyte with the same dielectric, or even a completely different\nelectrolyte with the same dielectric constant, would the same performance be expected? Or is the\nperformance reported in the work specific for EA at the studied composition?\nResponse: Thank you for your careful review. The overall dielectric constant (ε) of the electrolyte\nas a quantitative preliminary screening indicator is a key design criterion for low-temperature\naqueous electrolytes.\nIn our work, “dielectric engineering” means deliberately tuning the bulk permittivity of the\nsolvent matrix by mixing a high-permittivity and a low-permittivity component to place the\nelectrolyte in a medium-permittivity regime. In Zn electrolytes, this regime promotes moderate\ncation–anion association, lowers effective water activity, and increases the likelihood of anion\nparticipation in the primary solvation shell. The combined effect is to support SEI formation at\nlow temperature while retaining adequate ionic transport, resulting in improved battery\nperformance.\nWhile the dielectric is a primary design lever and a quantitative screening metric, it is not\nsufficient on its own. Different solvents adjusted to the same bulk permittivity can still yield\ndistinct solvation motifs, kosmotropic/chaotropic effects, and reduction pathways; these specific\ninteractions co-determine interfacial kinetics and stability. The trend we emphasize is nevertheless\ngeneral. For example, Tao et al. used different electrolytes with varying degrees of cation-anion\nassociation, including a medium dielectric constant electrolyte 2 M Zn(BF ) -AN: H O (7:3, v: v)\n4 2 2\nP\nA\nG\nE\n\\\n(ε=46.85), a low dielectric constant electrolyte 2 M Zn(BF ) -THF: H O (7:3, v: v) (ε=15.28), and\n4 2 2\na high dielectric constant (ε =78.46) electrolyte 2 M Zn(BF ) -H O. They demonstrated that\nwater 4 2 2\nmoderate ion association electrolytes using medium-ε solvents like AN can keep high ion transport\nin the bulk electrolytes and stable anion-derived SEI formation ability on the Zn anode. (J. Am.\nChem. Soc. 2024, 146, 46, 31612–31623). This supports targeting a medium-permittivity window\nrather than any single solvent.\nIn this work, we selected EA/H O (molar ratio 1:16) to construct an electrolyte system with a\n2\nmoderate dielectric constant (ε≈25.25). This design has the following advantages: 1: Greater\nflexibility in composition: the dielectric environment of the electrolyte can be precisely controlled\nby adjusting the ratio of high/low dielectric constant solvents. 2: practical screening criterion:\nproviding a quantifiable preliminary screening criterion for the design of low-temperature-resistant\naqueous electrolytes. However, the permittivity is a screening variable, not a sole determinant\n(solvent identity still matters).\nWe have provided corresponding supplementary descriptions in the main text (page 3, line 69):\n“we proposed a dielectric matching strategy, placing the aqueous electrolyte in a\nmedium‑permittivity window by mixing a high‑ε and a low‑ε component. Wherein high-ε solvents\npromote the dissociation of zinc salts and low-ε solvents optimize interface stability and ion\ntransport kinetics at low temperatures.”\n2) Other specific comments:\na) Line 130, “pure electrolyte” likely means “pure aqueous electrolyte”, but it could also mean\n“pure EA electrolyte”. Please clarify.\nResponse: Thank you for pointing out the problem. Thank you for catching these ambiguities. We\nnow write “The electrolyte without EA” instead of “pure electrolyte” in main text on page 5, line\n135:\n“The electrolyte without EA has a crystallization temperature (T ) of −38.5 °C, and adding EA at\nc\na 1:20 ratio reduced the crystallization temperature to −51.3 °C.”\nb) Line 138: “nEA:nH O” should be “EA:H O”.\n2 2\nP\nA\nG\nE\n\\\nResponse: Thank you for catching these ambiguities. We have corrected “nEA:nH O” to EA:H O\n2 2\nin the main text on page 5, line 142:\n“Among the tested EA:H O ratios, 1:16 delivered the highest ionic conductivity at −50 °C\n2\n(24.92 mS cm⁻1)”\nc) Starting from Line 155, the authors discussed the FTIR results. The authors mentioned “red\nshift” and “blue shift”, but it is hard to see from the context the shifts are relative to what. On lines\n162-163, the terms “left shift” and “right shift” were used whereas Figures 2a and 2b show FTIR\nin opposite frequency orders, which make it hard to follow. These terms are not professional either.\nOpposite x-axis is also seen in Figure 6 in SI. Please make it consistent throughout the paper.\nResponse: Thank you for flagging the ambiguity and axis inconsistencies. We standardized all\nFTIR plots to the same axis orientation—decreasing wavenumber from high to low—across the\nmain text and SI (updated Figs. 2a, b and Supplementary Figs. 8), and we revised captions to state\nthe reference and axis direction. To quantify changes, spectra were baseline-corrected and\ndeconvolved. The C–O–C mode shifts from 1046 to 1040 cm−1 and the O–Cl–O mode from 1079\nto 1073 cm−1 with EA addition. Peak assignments used in the analysis are: SSIP (non-coordinated\nClO −, 1076 cm−1), CIP (Zn–(ClO −)n≥1, 1047 cm−1), free anion (1114 cm−1), and free EA (1037\n4 4\ncm−1). Area evolution shows a decrease of SSIP and free-anion contributions and a corresponding\nincrease of CIP, consistent with enhanced anion participation in the primary solvation shell.\nSpectra were baseline-corrected and fitted with Voigt profiles using shared-width constraints; peak\npositions, FWHM, and integrated areas were extracted. These revisions remove ambiguity and\nmake the comparison reproducible.\nWe have added the corresponding content in the manuscript on page 6, line 162:\n“To comprehensively understand the molecular mechanisms of different segments in MDCE, the\ncoordination of Zn2+ with ClO – anions were investigated by FTIR spectroscopy (Supplementary\n4\n1076 cm–1), contact ion pairs (CIPs, Zn-(ClO –) , 1047 cm–1), indicating strong adsorption for\n4 n≥1\nClO –, free anion (1114 cm–1) and free EA (1037 cm–1). The results indicate that upon the\nA\nG\nE\n\\\nintroduction of EA, the content of SSIP and free anions in the electrolyte decreases, while the\nconcentration of CIP increases. Additionally, a fraction of EA exists as free solvent.”\nFig. R22/Supplementary Fig. 8 a, FTIR spectra of 2 M Zn(ClO ) H O/EA electrolytes with\n4 2 2\ndifferent EA:H O ratios from 0:1 to 1:6 at 4,000–400 cm−1. b, c FTIR spectra of H O, EA, 2 M\n2 2\nZn(ClO ) H O and 2 M Zn(ClO ) EA:H O (EA:H O = 1:16) electrolytes at 3,800–800 cm−1 (b)\n4 2 2 4 2 2 2\nand 1,200–1,000 cm−1 (c). d, Deconvolution result of FTIR spectra for 2 M Zn(ClO ) H O and 2\n4 2 2\nM Zn(ClO ) EA:H O (EA:H O = 1:16) electrolytes at and 1,200–1,000 cm−1.\n4 2 2 2\nd) Line 185 mentioned average number of HB but it is not clear average over what? Are the\nnumbers per water molecule or per acceptor site or any other definition?\nResponse: Thank you for pointing out the problem. The average number of HB is per acceptor\nsite (water or EA), including water and EA. The calculation formula is as follows:\nTotal number of HBs\nAverage HBs = \nTotal number of acceptor sites for all molecules\nWe have added the corresponding calculation methods in the manuscript on page 24, line 633:\n“The calculation formula for the average value of hydrogen bonds is as follows:\nP\nA\nG\nE\n\\\nTotal number of HBs (4)\nAverage HBs = \nTotal number of acceptor sites for all molecules\n”\ne) For the FTIR results shown in Figure 2, what is the temperature? Does FTIR show the same\ntemperature dependent trend as the simulation? Temperature is not provided for most of the results\nreported in the manuscript.\nResponse: Thank you for highlighting this. The FTIR spectra in Figure 2 were collected at 25 °C;\nwe now state the temperature explicitly in the figure caption and in Methods. We also standardized\ntemperature reporting across the manuscript (electrochemistry, spectroscopy, and conductivity),\nand corrected axis labels where needed for consistency:\n“In situ electrochemical FTIR was performed on a Vertex 70 spectrometer (Bruker) coupled to a\nCHI760E workstation and equipped with a liquid‑nitrogen‑cooled broadband MCT detector at\nroom temperature. (Page 22, line 579);\nConfocal imaging employed a Leica TCS SP5II system with 458 nm excitation and 490-550 nm\ndetection range using a 20× dry objective at room temperature. (Page 22, line 592);\nA confocal Raman microscope (Renishaw, inVia-Reflex) equipped with a green diode laser\n(λ = 532 nm) was used for the Raman analysis at room temperature. (Page 23, line 597);\nThe conductivity of the electrolyte at 25°C, 0°C, −20°C, −40°C, and −50°C was measured using\nthe Mettler Toledo Conductivity Meter (FE38-Standard). (Page 22, line 607)”\nf) The sentences on lines 229-233 are hard to follow. In the beginning it mentioned\n“immobilization of the ClO − anion”, what does “immobilization” mean here? Then it talks about\n4\nsolvents: “solvents with stronger binding properties tend to produce more negative solvation\nenergies”. In the next sentence, it goes back to ion: “corresponding to a greater number of ion\npairs.” Then the final part, what does “reduced ion dissociation” mean? Is this about structure or\nkinetics?\nP\nA\nG\nE\n\\\nResponse: The phrase “immobilization of the ClO ⁻ anion” indicates that ClO ⁻ serves as the\n4 4\nfixed anionic component in both HDCE and MDCE systems. Different anion salts may influence\nthe test results, all electrolytes used in this study are zinc salts containing ClO ⁻ anions. This\n4\neliminates the need for extensive discussion of alternative anion systems, thereby avoiding\npotential reader confusion or misinterpretation regarding other salt types, so we removed this\nsentence from the main text.\nSince solvation energy measurements can be influenced by both the solvent and the salt\ncomposition, we specifically employed ClO ⁻ as the sole anion throughout this study to isolate and\n4\ncompare the effects of different solvents on solvation energy. “Reduced ion dissociation” refers to\nthe ability of a solvent to dissociate cation-anion pairs. Solvents with low dielectric constants have\npoor ion dissociation capabilities, adding a low dielectric constant solvent to a high dielectric\nconstant solvent can reduce the degree of ion dissociation in the high dielectric constant electrolyte,\nthereby increasing the number of ion pairs in the high dielectric constant electrolyte. To enhance\nclarity and avoid redundancy, we have revised the manuscript language in main text on page 9,\nline 245:\n“Because stronger binding yields more negative solvation free energies, the more negative ΔGₛₒₗᵥ\nmeasured for MDCE (Fig. 3a; Supplementary Table 6) indicates tighter Zn²⁺ binding and,\ncorrespondingly, a higher fraction of ion pairs in the liquid[37].”\n[37] Wang, H., Kim, S. C., Rojas, T., Zhu, Y., Li, Y., Ma, L., Xu, K., Ngo, A. T. & Cui, Y.\nCorrelating Li-ion solvation structures and electrode potential temperature coefficients. J. Am.\nChem. Soc. 143, 2264-2271 (2021).\ng) Lines 240-241: E were reported for HDCE and MDCE, and commented that “the latter\na\n(MDCE) being notably lower than …” The E of HDCE and MDCE are only 1.2 kJ/mol different\na\nfrom each other, which is lower than kT. These E are similar to me.\na\nResponse: Thank you for the insightful comment regarding the apparent similarity in activation\nenergies (E ) between HDCE and MDCE. First, although the desolvation energy difference\na\nP\nA\nG\nE\n\\\nbetween HDCE and MDCE is relatively small, based on the Arrhenius equation (𝑘 = 𝐴𝑒− 𝐸 𝑅 𝑎 𝑇)，it\ncan be inferred that the zinc ion desolvation rate in MDCE is 1.623 times that of HDCE:\n𝑘𝑀𝐷𝐶𝐸\n= 𝑒\n−(𝐸𝑎𝑀𝐷𝐶\n𝑅\n𝐸\n𝑇\n−𝐸𝑎𝐻𝐷𝐶𝐸) \n=1.623\n𝑘𝐻𝐷𝐶𝐸\nA relatively small difference in desolvation energy can also significantly impact the reaction rate.\nMoreover, the resulting E values were 22.1 kJ mol⁻1 for HDCE and 20.9 kJ mol⁻1 for MDCE,\na\nwith the latter being notably lower than values reported for most comparable electrolyte systems.\nOur use of “notably lower” referred to comparison with many reported electrolytes rather than\nto HDCE specifically. For context, the table below compiles representative values (Table R2).\nTable R2 Compared with the desolvation energy reported in previous literature.\nElectrolyte E (kJ mol−1) Ref\na\n2 2\n54.8 Nat Commun 14, 5443 (2023)\nvolume ratios of 2:3\n0.5 M Zn(SO ) ∙7H Obutanone/H O\n4 2 2 2\n30.6 Nat Commun 15, 302 (2024)\nvolume ratios of 1:7\n2 2\n41.65 Nat Commun 14, 3067 (2023)\nweight ratios of 0:1\n2 2\n65.15 Nat Commun 14, 3067 (2023)\nweight ratios of 3:7\nAngew. Chem. Int. Ed. 2024,\nZn(Ac) /Fad molar ratios of 1:9 22.53\n2\n63, e202406906.\n4 2 2\n22.1 This work\nmolar ratios of 0:1 (HDCE)\n4 2 2\n20.9 This work\nmolar ratios of 1:16 (MDCE)\nh) Line 578 says the OPLS-AA force field was used in the study. Does OPLS-AA include\nparameters for Zn2+ ion? If not, please specify what model was used for Zn2+.\nP\nA\nG\nE\n\\\nResponse: Thank you for pointing this out. The built-in OPLS-AA force field does not include\nparameters for Zn2+ ions. In our study, we employed a Zn2+ model based on parameters reported\nin a previous publication (J. Chem. Theory Comput. 2020, 16, 4429−4442), which we further\noptimized for aqueous zinc battery simulations. The employed Zn2+ model has been validated in\nseveral previous studies (e.g., Adv. Mater. 2025, 2502649; Adv. Funct. Mater. 2022, 32, 2109322),\nand has demonstrated good accuracy in reproducing the solvation structure and transport behavior\nof Zn2+ in aqueous electrolytes.\nThe fit results of EXAFS also confirmed the validity of the OPLS-AA force field. The Zn-O\n(H2O)\nZn-O and Zn-O bond length in Zn(H O) (ClO )+ and Zn(H O) EA(ClO )+ acquired by MD\n(ClO4) (EA) 2 5 4 2 4 4\nsimulation, EXAFS fitting, and DFT calculation are consistent within the error of 5% and\nsummarized in Table R3.\nTable R3. The bond length of Zn–O in HDCE and MDCE by MD simulation, EXAFS fitting, and\nDFT calculation.\nMethod Path MD R (Å) DFT R (Å) RDF R (Å)\nZn-O (H O) 2.13 2.14 2.04\nHDCE (path in 2\nZn(H 2 O) 5 (ClO 4 )+) Zn-O (ClO 4 ) 2.09 2.09 2.01\nZn-O (H O) 2.13 2.16 2.04\nZn-O (ClO ) 2.09 2.14 2.01\n2 4 4\nZn-O (EA) 2.15 2.07 2.06\ni) For the simulation study, validation of the force field needs to be provided.\nResponse: Thank you for raising this point. We have noted the previously reported examples:\nOPLS-AA combined with the Zn2+ model has been successfully applied to aqueous Zn electrolytes,\nreproducing solvation structure and transport trends (e.g., Proc. Natl. Acad. Sci. 120, e2221980120\n(2023), Nat. Commun. 16, 3820 (2025), Nat. Commun. 16, 5740 (2025), Nat. Commun. 15, 6471\n(2024), Nat. Commun. 16, 1273 (2025), Nano-Micro Lett. 18, 3 (2025), J. Am. Chem. Soc. 144,\nP\nA\nG\nE\n\\\n7160-7170 (2022), Energy Environ. Sci. 13, 3527-3535 (2020), Adv. Mater., 2502649.\nhttps://doi.org/10.1002/adma.202502649, ACS Energy Lett. 8, 608-618 (2023)). These works\ncollectively support the validity and applicability of the OPLS-AA force field for modeling the\nstructural and transport properties of aqueous zinc battery systems.\nOverall, the work presented a systematic study of an electrolyte system and insightful results were\nreported. An emphasis on dielectric engineering was made. However, it is not clear what the\nauthors mean by “dielectric engineering”. On the other hand, the effect of dielectric constant has\nbeen discussed in the literature. Works on low temperature electrolyte are available in the literature\ntoo. Therefore, I do not think the work carry enough novelty to be published in Nature\nCommunication. I would suggest the authors address above comments and resubmit it to a different\njournal.\nResponse: Thank you for the suggestions. We agree that the manuscript needed a clearer definition\nof “dielectric engineering” and a sharper articulation of what is new. We will respond to the\nreviewers’ questions as follows:\nFirst, we clearly define the specific meaning of “dielectric engineering” in the context of this\nstudy. This study proposes a dielectric matching strategy based on the physical property of the\nsolvent’s dielectric constant (ε), whereby high-ε solvents facilitate zinc salt dissociation and low-\nε solvents optimize interface film formation.\nSecond, the reviewers noted that previous studies have discussed dielectric constants and low-\ntemperature electrolytes. Our work fundamentally differs from these previously reported studies\nfor the following reasons:\n1) Mechanism difference.\nThis study presents differences in mechanisms compared to previous research. Chen et al. aim\nto facilitate coordination of the low-ε solvents with Li+ to occupy its coordination sites, thus\nconverting the solvent-separated ion pairs (SSIPs) from high-ε solvent-dominated to low-ε-\ndominated (Nat. Commun. 14, (2023): 8326). Yang et al. (Nat. Sustain. 6, 325–335 (2023)) and\nWang et al. (Nat Commun. 14, 2023: 5443) other researchers have predominantly focused on how\nthe regulation of solvation structures by co-solvents enhances low-temperature performance,\nP\nA\nG\nE\n\\\nwithout in-depth investigation into the evolution of interfacial solvation structures under an electric\nfield and their role in promoting SEI formation.\nWe proposed a dielectric matching strategy, wherein high-ε solvents promote the dissociation\nof zinc salts, while low-ε solvents optimize interface stability and ion transport kinetics at\nlow temperatures. And we focus on exploring how the dielectric properties of co-solvents affect their\nsolvation structure, and clarify their relationship with low-temperature ion transport dynamics and\ncycling stability.\n2) Deeper consideration of the challenges posed by low-temperature conditions.\nSome studies have reported regulating the solvation structure to lower the crystallization\ntemperature of electrolytes, while suppressing HER (hydrogen evolution reaction) and Zn dendrite\ngrowth. However, the following issues remain to be addressed: 1. Although the freezing point of\nelectrolytes can be effectively reduced, this is often achieved at the expense of ion transport\nkinetics. Minimizing the loss of ionic conductivity under ultra-low temperature conditions while\nmaintaining the stability of the solvation structure is key to achieving efficient zinc ion transport.\nTo address this, we selected chaotropic ClO ⁻ anions, which exhibit strong hydrogen bond\n4\ndisruption capabilities, along with low-viscosity ethyl acetate, to minimize the ionic transport\nimpedance at low temperatures. 2. The cycle life of existing low-temperature zinc metal anodes is\ngenerally limited to a few hundred hours, which is insufficient for practical applications. To\novercome the limitations of cycle stability, we investigated how low dielectric constant influences\nthe regulation of interfacial solvation structure, promoting SEI formation and the potential\nmechanism for reversible deposition/dissolution at low temperatures. Our zinc metal anode\nsymmetrical cell provides an ultra-long time of 7,200 hours of stable cycling at 25 °C, and 4,000\nhours at the ultra-low temperature of −50 °C with an ultra-thin zinc foil of 30 μm.\n3) Deeper investigation into interfacial electrochemical reactions.\nMost work has focused on how the regulation of solvation structures by co-solvents enhances\nlow-temperature performance, lacking in-depth investigation into the evolution of interfacial\nsolvation structures under an electric field and their role in promoting SEI formation. We addressed\nthis issue by employing advanced in-situ characterization techniques. For the first time, in situ\ntwo-dimensional pH characterization was conducted using confocal laser scanning microscopy\n(CLSM) to monitor the HER of the electrolyte. Operando infrared elucidates interfacial dynamics\nP\nA\nG\nE\n\\\nand decomposition mechanisms under an electric field, providing insights into solid/liquid\ninterfaces and solvation dynamics crucial for cryogenic aqueous batteries.\nWe emphasize generality over solvent specificity; EA is a useful example of the medium-ε\nstrategy. The revised text now presents dielectric engineering as a practical screening criterion,\nincluding permittivity range and solvation/interfacial checks, clarifies how our mechanism goes\nbeyond a single cosolvent, and places the results within the context of existing literature. We hope\nthese clarifications and additions address the reviewer’s concerns about definition and novelty.\nP\nA\nG\nE\n\\\nResponse to the reviewers’ comments\nWe sincerely appreciate the reviewer’s endorsement of our work, and their insightful\nand constructive input. The manuscript has now been carefully revised, and the\nrevisions are highlighted by yellow color. Our point-to-point responses to the comments\nare given in the following.\nComment: This manuscript has been revised as the comments, so it is suitable for\npublication without change.\nResponse: Thank you for your endorsement of our work.\nReviewer #2:\nComment: The authors provided a complete response to my questions. I suggest the\nconsideration of this work.\nResponse: Thank you for your endorsement of our work.\nReviewer #3:\nComment: All my questions have been well addressed. It can be accepted in this\njournal.\nResponse: Thank you for your endorsement of our work.\nReviewer #4:\nComment: The authors addressed most of my comments. But critical details of the\nMD simulations are still missing. The authors said an optimized Zn2+ model was used\nin the study. The parameters, reference and/or other optimization details should be\nprovided so that readers can reproduce the work.\nProper validation of the used force fields parameters for the studied systems are also\nneeded. General force fields like OPLS can work on one system but fail badly on\nanother. The choice of water model and other simulation protocols can change the\nresults too. In spite of the references provided, direct comparison with experimental\nmeasurement on the same systems should be provided.\nResponse: We thank the reviewer for the insightful and thoughtful comment.\nWe agree with the reviewer that the OPLS-AA force field may present limitations in\naccurately describing Zn2+ systems. To comprehensively address this concern, we\nrevalidated our previous MD simulations using the AMBER force field. In the new\nAMBER-based simulations, we employed the Zn2+ parameters developed by Merz and\nco-workers (Lennard–Jones 12–6, IOD parameter set, compatible with the OPC water\nmodel) [J. Chem. Theory Comput. 2020, 16, 4429−4442]. The results indicate that the\nsolvation structure, dielectric constant, and hydrogen-bonding characteristics obtained\nfrom the AMBER force field are consistent with those derived from OPLS-AA, with\n2\nonly minor deviations that do not alter our main conclusions. To fully address the\nreviewer’s concerns, all MD simulation results in the manuscript have been\nreplaced with those obtained from the AMBER force field, and the corresponding\nsections have been revised accordingly. The specific changes are detailed below.\nWe have added the relevant MD simulation details in Calculation part from the main\ntext (page 23, line 614) as:\n“Molecular dynamics (MD) calculations: MD simulations were carried out using\nGROMACS57 2024.4. The AMBER force field58, OPC water model59 and restrained\nelectrostatic potential (RESP) charge (obtained with Multiwfn60) were employed with\npreviously reported Zn2+ parameter of Lennard-Jones 12-6 model (IOD set)61.\npotential (resp) model perform in calculating conformational energies of organic\nand biological molecules? J. Comput. Chem. 21, 1049-1074 (2000).\napproach. J. Phys. Chem. Lett. 5, 3863-3871 (2014).\nChem 33, 580-592 (2012).\nmetal ions for the opc3, opc, tip3p-fb, and tip4p-fb water models. J. Chem. Theory\nComput. 16, 4429-4442 (2020).”\nTo fully address the reviewer's concerns regarding the use of OPLS-AA for\ncalculating Zn2+, we have replaced all MD calculations with the AMBER force field-\nbased simulations as below:\n1. Solvation structure analysis\nUnder identical computational conditions, the radial distribution functions (RDF) and\ncoordination number values generated by both methods exhibited similar results. The\nMD simulations presented in Figures R1 and R2 employed AMBER force fields, while\nFigure R3 utilized the OPLS-AA force field. We have updated the corresponding results\naccordingly on page 8, line 211:\n“For MDCE, as the temperature decreased from 25 °C to −50 °C, the average CN of\nZn2+ with oxygen in EA and ClO ⁻ increased slightly, from 0.19 to 0.21 and from 1.35\n4\nto 2.12, respectively (Fig. 2f and Supplementary Fig. 11d).”\nHDCE at −50°C.\nFig. R2/Supplementary Fig. 11. (a,b) Snapshots of representative Zn2+ solvation\nmotifs in HDCE and MDCE at 25 °C (a) and −50 °C (b); anion distributions are\nhighlighted. (c,d) Distributions of Zn2+ solvation species and RDFs/CNs in MDCE (c)\nand HDCE (d) at 25 °C.\nMDCE (b, d) at 25 °C and −50 °C; anion distributions are highlighted. (e-h)\nDistributions of Zn2+ solvation species and RDFs in HDCE (e, f) and MDCE (g, h) at\n25 °C. (i-l) Distributions of Zn2+ solvation species and RDFs in HDCE (i, j) and MDCE\n(k, l) at −50 °C.\nWe performed a comprehensive statistical analysis of the solvation structures. The\nstatistical results reveal similar distributions of solvation structures across different\nforce field implementations. Figure R4 presents results obtained using AMBER force\nfield, while Figure R5 displays results obtained using OPLS-AA force fields.\nFig. R4/Fig. 2g Distribution of primary solvation structures obtained from MD\nsimulations.\n2. Statistics of hydrogen bonds\nWe conducted a statistical analysis of the solvation structures, which revealed similar\naverage numbers of hydrogen bonds across different force fields. The results presented\nin Figures R6 and R7 were obtained using AMBER force fields, while Figures R8 and\nR9 show results from OPLS-AA force fields. The corresponding statistical data have\nbeen updated accordingly on page 8, line 200:\n“At 25 °C, the nHBs for HDCE and MDCE were calculated to be 2.32 and 2.03,\nrespectively (Supplementary Fig. 12). Upon cooling to −50 °C, HDCE exhibited a\nsubstantial increase in hydrogen bonding, with nHBs rising to 2.69, as visualized by the\ncolor change in Fig. 2e. In contrast, MDCE showed a more moderate increase, with\nnHBs reaching only 2.28 (Fig. 2d).”\nFig. R6/Fig. 2d, e Radial angular distribution functions, g(r, θ), for MDCE and HDCE\nat 223 K.\nMDCE and HDCE at 298 K.\nFig. R8 Radial angular distribution functions, g(r, θ), for MDCE (d) and HDCE (e) at\nFig. R9 Radial angular distribution functions, g(r, θ), for MDCE and HDCE at 298 K.\n3. Calculation of dielectric constant\nobtained using AMBER force fields, while Figures R11 show results from OPLS-AA\nforce fields. The corresponding statistical data have been updated accordingly on page\n5, line 122:\n“The dielectric constant of the Zn(ClO ) /H O electrolyte is 42.48, categorizing it as a\n4 2 2\nhigh-dielectric-constant electrolyte (HDCE). In contrast, the Zn(ClO4) /EA electrolyte\n2\nexhibits a much lower dielectric constant of 5.47. When Zn(ClO ) is dissolved in a\n4 2\nwater-EA mixture, the resulting electrolyte has an intermediate dielectric constant of\n24.13 (Fig. 1e), and is thus designated as a medium-dielectric-constant electrolyte\n(MDCE).”\nFig. R10/Fig. 1e ε values of 2 M Zn(ClO ) in H O, 2 M Zn(ClO ) in H O/EA\n4 2 2 4 2 2\n(EA/H O = 1:16, molar ratio) and 2 M Zn(ClO ) in EA electrolytes.\n2 4 2\nFig. R11 ε values of 2 M Zn(ClO ) in H O, 2 M Zn(ClO ) in H O/EA (EA/H O =\n4 2 2 4 2 2 2\n1:16, molar ratio) and 2 M Zn(ClO ) in EA electrolytes.\n4 2\nimproved the quality and clarity of our manuscript. We hope that the revisions made in\nresponse will fully address the reviewer’s concerns.\nWe sincerely appreciate the reviewer’s endorsement of our work, and their insightful\nand constructive input. The manuscript has now been carefully revised, and the\nrevisions are highlighted by yellow color. Our point-to-point responses to the comments\nare given in the following.\nP\nA\nG\nE\n\\\nReviewer #4:\nComment: The authors provided more simulation details including information about\nthe force field parameters in the revised manuscript. Considering the fact that\nsimulation is only part of this work, I would support the publication of the work if the\nauthors can provide input files for the simulation in the SI, including enough\ninformation about simulation details and all the force field parameters that others can\nreproduce the work.\nMeanwhile, I want to make it clear to the authors that consistent results using two sets\nof force fields do not necessarily validate the simulation results. As I commented\npreviously, a direct comparison with experimental measurements on the same system\nunder the same condition is needed for this purpose.\nResponse: We sincerely thank the reviewer for the supportive comments. Following\nthis suggestion, we have now included a set of simulation details and the force field\nparameters in the Supporting Information (SI) as following:\nMore details are provided for the MD simulations: The parameters of Zn2+ were\nobtained from previously reported research7, (R /2 = 1.373Å, ε = 0.01179373\nmin,M M\nkcal mol-1). The parameters of ClO - and EA were generated with Sobtop8. The\n4\nsimulations were conducted using AMBER force field with the OPC water model. The\nproduction MD simulation was performed for 20 ns with a time step of 1 fs using the\nleap-frog integrator. Bond constraints were handled using the LINCS algorithm with\nan order of 4 and one iteration. A Verlet cutoff scheme was employed for neighbor\nsearching with a frequency of 10 steps. Both the short-range electrostatic and van der\nWaals cutoffs were set to 1.0 nm. Long-range electrostatic interactions were treated\nusing the Particle Mesh Ewald (PME) method with a fourth-order interpolation and a\nFourier grid spacing of 0.16 nm. The temperature was maintained at 298 K or 223K\nusing the Nosé-Hoover thermostat with a time constant of 0.5 ps applied to the system.\nThe pressure was controlled at 1.0 bar using the Parrinello-Rahman barostat with\nisotropic coupling, a time constant of 2.0 ps, and a compressibility of 4.5×10-5 bar-1.\nPeriodic boundary conditions were applied in all three directions.\nP\nA\nG\nE\n\\\nmetal ions for the OPC3, OPC, TIP3P-FB, and TIP4P-FB water models. J. Chem.\nTheory Comput. 16, 4429-4442 (2020).\nWe sincerely appreciate the reviewer's insightful comment and fully concur that\nexperimental comparison is the gold standard for validation. We will be more mindful\nof this in our future work. Thank you.\nP\nA\nG\nE\n\\"}, {"round": 2, "reviewer_comments": {"reviewer_4": "The authors provided more simulation details including information about the force field parameters in the revised manuscript. Considering the fact that simulation is only part of this work, I would support the publication of the work if the authors can provide input files for the simulation in the SI, including enough information about simulation details and all the force field parameters that others can reproduce the work.\nMeanwhile, I want to make it clear to the authors that consistent results using two sets of force fields do not necessarily validate the simulation results. As I commented previously, a direct comparison with experimental measurements on the same system under the same condition is needed for this purpose."}, "author_response": "Response to the reviewers’ comments\nWe sincerely appreciate the reviewer’s endorsement of our work, and their insightful\nand constructive input. The manuscript has now been carefully revised, and the\nrevisions are highlighted by yellow color. Our point-to-point responses to the comments\nare given in the following.\n\nReviewer #4:\nComment: The authors provided more simulation details including information about\nthe force field parameters in the revised manuscript. Considering the fact that\nsimulation is only part of this work, I would support the publication of the work if the\nauthors can provide input files for the simulation in the SI, including enough\ninformation about simulation details and all the force field parameters that others can\nreproduce the work.\nMeanwhile, I want to make it clear to the authors that consistent results using two sets\nof force fields do not necessarily validate the simulation results. As I commented\npreviously, a direct comparison with experimental measurements on the same system\nunder the same condition is needed for this purpose.\nResponse: We sincerely thank the reviewer for the supportive comments. Following\nthis suggestion, we have now included a set of simulation details and the force field\nparameters in the Supporting Information (SI) as following:\nMore details are provided for the MD simulations: The parameters of Zn2+ were\nobtained from previously reported research7, (R /2 = 1.373Å, ε = 0.01179373\nmin,M M\nkcal mol-1). The parameters of ClO - and EA were generated with Sobtop8. The\n4\nsimulations were conducted using AMBER force field with the OPC water model. The\nproduction MD simulation was performed for 20 ns with a time step of 1 fs using the\nleap-frog integrator. Bond constraints were handled using the LINCS algorithm with\nan order of 4 and one iteration. A Verlet cutoff scheme was employed for neighbor\nsearching with a frequency of 10 steps. Both the short-range electrostatic and van der\nWaals cutoffs were set to 1.0 nm. Long-range electrostatic interactions were treated\nusing the Particle Mesh Ewald (PME) method with a fourth-order interpolation and a\nFourier grid spacing of 0.16 nm. The temperature was maintained at 298 K or 223K\nusing the Nosé-Hoover thermostat with a time constant of 0.5 ps applied to the system.\nThe pressure was controlled at 1.0 bar using the Parrinello-Rahman barostat with\nisotropic coupling, a time constant of 2.0 ps, and a compressibility of 4.5×10-5 bar-1.\nPeriodic boundary conditions were applied in all three directions.\nmetal ions for the OPC3, OPC, TIP3P-FB, and TIP4P-FB water models. J. Chem.\nTheory Comput. 16, 4429-4442 (2020).\nWe sincerely appreciate the reviewer's insightful comment and fully concur that\nexperimental comparison is the gold standard for validation. We will be more mindful\nof this in our future work. Thank you."}]

Decision letter