Differentiating interfacial water structures via alkali metal cation promotor for H_2O_2 electrosynthesis in acid

Domain: neuroscience
Published: 2026-04-08
Editorial outcome: minor_revision

Abstract
Electrocatalytic oxygen reduction reaction (ORR) for H2O2 production represents a sustainable alternative route to the energy-intensive anthraquinone process. Nevertheless, under industrially-relevant acidic conditions, excessive protons at the reaction interface exacerbate low H2O2 selectivity and severe H2O2 reduction. Herein, we propose a universal alkali metal cation (AMC: Li+, Na+, K+, or Cs+) dosing strategy to markedly boost the acidic H2O2 electrosynthesis. Upon Cs+ addition, 2e− ORR selectivity increases from 20% to 80%, concurrently suppressing an H2O2 reduction current by 50% and achieving an H2O2 production rate of 9.2 mol g−1 h−1 at 500 mA cm−2. Microelectrode hydrogen evolution measurements witness impeded proton diffusion in AMC-dosed acidic electrolytes, directly restricting proton supply to catalytic active sites. In situ spectroscopic analysis combined with molecular dynamics simulation demonstrate AMCs help reconfigure interfacial water networks via cation hydration shells, thereby disrupting proton-hopping pathways. The efficacy trend (Li+ < Na+ < K+ < Cs+) originates from distinct cation-specific interfacial water restructure, delivering mechanistic insights into cation-promoted selective H2O2 electrosynthesis in acidic media.

Introduction
Hydrogen peroxide (H2O2) has emerged as a green and versatile oxidant with vast application potentials in wastewater treatment, organic synthesis and chemical technology1-7. Its on-site electrosynthesis via the two-electron oxygen reduction reaction (2e− ORR) with enriched choices of electrocatalysts8, offers an appealing alternative to the energy- and waste-intensive anthraquinone process9,10. Nevertheless, the practical implementation at industrially-relevant current densities is hindered by activity and stability limitations, primarily due to competing 4e− reduction pathways and H2O2 decomposition11-14.

Beyond deliberate electrocatalysis, the fundamental chemistry of aqueous interfaces has garnered intense interest for its role in spontaneous chemical transformations15-18. Notably, the purported spontaneous formation of H2O2 has sparked significant debate. While early reports emphasized remarkable accelerations at the air-water interface19-23, a growing body of evidence, including rigorous work from the Mishra group, suggests that such phenomena might be more pronounced at solid–water interfaces24-27. This ongoing controversy underscores that the intrinsic drivers of reactivity in confined aqueous environments remain a fundamental question in interfacial science. Resolving this question is critical not only for understanding natural processes but also for the rational manipulation of electrolyte–catalyst interfaces.

Electrolyte engineering, particularly the introduction of alkali metal cations (AMCs: Li+, Na+, K+, or Cs+)28-30, has emerged as a promising strategy to steer interfacial chemistry toward H2O2 electrosynthesis in acid31,32. AMCs are generally understood to mitigate the local proton activity, thereby stabilizing the key *OOH intermediate and inhibiting H2O2 decomposition. Particularly using metal-free carbon-based electrocatalysts, Zhang et al. observed that H2O2 production rates can be improved by an AMC-induced shielding effect, yet were relatively unaffected by the AMC types33. A recent study indicated a graphitic carbon electrode exhibited an increase in 2e− ORR selectivity following the trend of Cs+ > K+ > Na+ > Li+, where different cations would interfere in the *OOH binding and influence the local electric field34. It is interesting to note that the AMC roles in acidic 2e− ORR manifest inconsistent trends. In further contexts, the cation effect is effective across other catalyst types apart from carbonaceous materials. For instance, Cao et al. found the H2O2 production followed the trend of K+ > Na+ > Li+ ~ Cs+ with the employment of earth-abundant TiC catalyst, which was dictated by the hydrated ionic radius of AMCs35. These disparities imply that the dominant mechanism might be sensitive to specific conditions, such as catalyst identity, cation concentration and operating current density. A coherent picture of how AMCs reshape the interfacial environment in acidic 2e− ORR is still lacking, especially in the realm of prevailing transition metal-involved carbon electrocatalysts (e.g., metal-carbon nanotube, M-CNT).

Inspired by these considerations, we pay attention to a previously underappreciated role of AMCs that lies in their specific capability of reconfiguring the interfacial water to varied degrees. Such an interfacial water restructure governs the kinetics of proton transport to the active site. By employing Co-CNT as a model catalyst and AMC concentration series (0.1–0.5 M), we observe a clear cation-dependent selectivity trend (Li+ < Na+ < K+ < Cs+) that correlates directly with spectroscopic signatures of interfacial water reconfiguration, as captured by in situ attenuated total reflectance surface-enhanced infrared absorption spectroscopy (ATR-SEIRAS). Ab-initio molecular dynamics (AIMD) simulations provide strong evidence that AMCs promote the formation of less hydrogen-bonding (H-bonding) or free water molecules. This would markedly elevate the proton transfer energy barrier to impede H2O2 reduction, which is confirmed by microelectrode measurements. The varying influences of different AMCs stem from their distinct capabilities to perturb the H-bond network, associated with their hydration properties and coordination strengths. Such a perspective is essentially different from the general recognition of AMC roles in this field (Fig. 1). Leveraging this insight, we achieve a considerable improvement in acidic H2O2 selectivity and a production rate of 9.2 mol g−1 h−1 at 500 mA cm−2 with Cs+ additives, readily enabling efficient pollutant degradation and bacterial disinfection under practical application scenarios.

Results
Investigating AMC-dependent proton transport suppression

It is a general recognition that proton concentration at the electrode surface critically governs acidic 2e− ORR process33,36,37. We employed gold microelectrode measurement to probe the influence of AMCs upon the proton delivery toward the electrode. This could be achieved by evaluating the diffusion-limited current density of the hydrogen evolution reaction (HER)38,39. First, linear sweep voltammograms (LSVs) were recorded on a 5-mm diameter Au rotating disk electrode in a test solution of 0.5 M H2SO4 containing 0.3 M Cs+ at rotation rates of 400, 900, 1600, and 2500 rpm. We observed that the HER current was entirely mass-transport-limited by proton diffusion (Supplementary Fig. 1), confirming the validity of our measurement methodology. In bare acidic solution, a featured diffusion-limited current plateau of ~3.3 A cm−2 was attained on the Au microelectrode. Upon the AMC dosage, such a current markedly decreased to below 0.6 A cm−2 (Figs. 2a and 2b; Supplementary Figs. 2 and 3). Moreover, the extent of current suppression exhibited a strong dependence on the AMC types. Li+ exerted the weakest suppression effect, whereas Cs+ produced the most pronounced suppression, reaching a maximum factor of 13-fold. It is further noted that there existed a concentration-dependent variation of this suppression effect (Supplementary Figs. 4 and 5). The optimized AMC concentration refers to a key interfacial reaction engineering of the balance between suppressing parasitic pathways and maintaining sufficient reaction kinetics. Considering the trade-off between proton availability and reaction kinetics, a medium concentration of 0.3 M was selected to assess the AMC concentration effect on the H2O2 yield (Supplementary Fig. 6).

Gu et al. proposed that AMCs could hamper the electromigration of protons in solution via an electric field shielding fashion38,40. To examine the electric field modulation on proton transport in our system, we employed [Ru(NH3)6]3+ ion as a model redox probe and measured the dependence of its diffusion rate with AMCs39. Our findings indicated that AMC dosage enabled the suppression of [Ru(NH3)6]3+ electromigration, but marginal changes in the suppression effect were detected with respect to the dosed AMC types, all yielding an inhibition factor of ~1.38 (Figs. 2c and 2d). Additionally, this effect did not exhibit further enhancement with increasing Cs+ concentration (Supplementary Fig. 7), which cannot account for the ~13-fold suppression factor (Fig. 2e). To exclude potential interference from other variables, the pH value, ionic conductivity and dissolved oxygen in different electrolytes were probed. As shown in Supplementary Figs. 8 and 9, the AMC dosage did not significantly alter the values, indicating such parameters were unlikely to be the primary drivers behind the observed cation-dependent trend. These results motivated us to investigate the restructuring of interfacial H-bonding water networks in the presence of AMCs.

To elucidate the influence of AMCs on the interfacial water structures, we carried out in situ ATR-SEIRAS measurements (Supplementary Fig. 10). The SEIRAS signals were acquired at the catalyst-electrolyte interface in 0.5 M H2SO4 containing Li+, Na+, K+, or Cs+. The spectra clearly exhibit the O–H stretching vibration band of interfacial water (3000–3600 cm–1) (Fig. 2f; Supplementary Fig. 11)41. The intensity displays apparent potential dependence, indicating that the signals primarily originated from the first few layers of water molecules adjacent to the electrode surface. It is noted that the O–H stretching vibration band under different potentials in bare acid and Li+-containing system were predominantly located at ~3270 cm–1, affording a redshift of ~130 cm–1 as compared to that in the Cs+-containing system (~3400 cm–1). Such a shift implied that interfacial water molecules formed strong H-bond interactions at the electrode surface in the presence of small AMCs, which is analogous to the scenario in bulk water. In contrast, electrolytes containing large AMCs induced a reduction in the number of H-bonds formed by interfacial water molecules, resulting in a higher O–H stretching vibration frequency.

In detail, the O–H stretching band was deconvoluted into three components centered at 3550, 3400, and 3200 cm–1 (Fig. 2g)42-48. These signals could be assigned to isolated water molecules, 2-coordinated H-bonded water (2-HB), and 4-coordinated H-bonded water (4-HB), respectively, with their cation-dependent relative contributions plotted in Fig. 2h. As such, large AMC significantly enhanced the population of 2-HB and isolated water species, suggestive of pronounced disruption of the H-bond network within the interfacial water layer. The extent of this disruption also afforded the order of Li+ < Na+ < K+ < Cs+, which is consistent with the trend observed in proton transport suppression. Collectively, these findings underline the importance of AMC dosage in disrupting the H-bond network and altering the interfacial H-bonded water structures.

Theoretical insight into AMC-suppressed proton transport

Computational simulations were carried out to validate the experimental observations. Prior to this, the prepared catalyst underwent a full spectrum of characterizations including electron microscopy (Supplementary Figs. 12–14), X-ray photoelectron spectroscopy (Supplementary Figs. 15–17), Fourier transform infrared spectroscopy (Supplementary Fig. 18), X-ray diffraction (Supplementary Fig. 19), and synchrotron radiation-based analysis (Supplementary Figs. 20–23; Supplementary Table 1). The introduction of potassium thiocyanate (KSCN) into the electrolyte is known to selectively bind and poison transition metal sites. As demonstrated in Supplementary Fig. 24, the addition of 10 mM KSCN leads to a pronounced decrease in disk current density, suggesting cobalt acts as the active site for the 2e– ORR. These unambiguously revealed the morphological, chemical and electronic information of catalysts.

We constructed an orthorhombic supercell model incorporating 1080 explicit water molecules. Protons (H+) and AMCs (Li+, Na+, K+, Cs+) were then introduced into such an aqueous environment to replicate the experimental electrolyte conditions (Supplementary Fig. 25). The system underwent 200 ps of pre-equilibration by NVT simulations to generate physically realistic initial configurations, followed by extensive equilibration by 2 ns NPT and 2 ns NVT. After equilibration, quantitative analysis of H-bond populations across the different systems revealed the sequence of Li+ < Na+ < K+ < Cs+, where the Cs+-containing system exhibited the lowest H-bond density per unit volume (Supplementary Fig. 26; Supplementary Table 2). Subsequently, path integral molecular dynamics (PIMD) simulations were performed to probe the proton transport in the presence of AMCs (Fig. 3a). Quantification of proton conduction pathways reveals that AMC suppresses long pathways while promoting short trajectories, with the trend remaining Li+ < Na+ < K+ < Cs+ (Figs. 3b and 3c; Supplementary Figs. 27–29). Moreover, larger AMCs coordinating more water molecules markedly reduce the number of H-bonds, as reflected in the radial distribution functions (RDFs) (Supplementary Fig. 30). These results collectively indicate that large-sized AMCs impede proton transport.

Gomez et al. elucidated a three-step mechanism for aqueous proton transport, governed by sequential H-bond exchange events: (i) the reduction in proton-acceptor coordination to facilitate proton transfer; (ii) the rate-limiting step involving enhanced proton-donor coordination to prevent back-transfer; and (iii) the dynamic equilibrium between Eigen and Zundel configurations49-51. Based upon the theoretical framework and in situ ATR-SEIRAS analysis, three distinct proton-donating water species were introduced as reference configurations, encompassing 4-HB water, 2-HB water, and isolated (free) water, with 4-HB water serving as the universal proton acceptor52. Consequently, donor-acceptor pairs were categorized into three types: Type I: 4-HB water (donor) → 4-HB water (acceptor); Type II: 2-HB water (donor) → 4-HB water (acceptor); Type III: Free water (donor) → 4-HB water (acceptor).

Density functional theory (DFT) calculations were performed to elucidate how different H-bond configurations affect proton transfer. In Type I system, the proton transfer from donor to acceptor proceeded in a stable manner. At the IM1 intermediate state, the proton-donor water distance ratio reached 1%, indicating the proton remains associated with the donor. At the IM6 state, this ratio reversed with the proton-acceptor water distance attaining 9%, demonstrative of successful proton transfer to the acceptor (Fig. 3d; Supplementary Fig. 31). In contrast, Type II systems exhibited comparatively slower proton transfer kinetics. The proton-donor water distance ratio merely shifted from 4% at IM1 to 82% at IM6 (Fig. 3d; Supplementary Fig. 32). During the transfer from IM1 to IM3, the proton reached the midpoint between the donor and acceptor (50%) more rapidly than in the Type I system. However, its mobility decreases significantly along the pathway from IM4 to IM6. After advancing to 82% of the donor–acceptor distance at IM5, the proton remained at this position in IM6. This behavior indicates that in the Type II system, the proton exhibited a stronger tendency to form and maintain a Zundel-like configuration ([H2O-H-OH2]+)51.

Paradoxically, while the Zundel state is a key intermediate, efficient proton transfer relies on the dynamic transition between this state and the intrinsic Eigen-type configuration. Stabilization of the Zundel state in the Type II system introduces a kinetic bottleneck, which likely suppresses overall proton transfer. Furthermore, the complete proton transfer in the Type II system shows an energy barrier of 0.011 eV, about twice that in the Type I system (0.005 eV), indicating hindered proton mobility considering the intrinsic uncertainty of DFT calculations (Fig. 3e). The kinetic impediment likely originated from reduced H-bond coordination of the donor water molecule in Type II systems, which stabilized the symmetric Zundel intermediate during proton transfer. Additionally, H-bond network deformation was observed during the 1-HB to 4-HB transition, manifested through the formation of 2-HB donors and 3-HB acceptors, suggesting that direct 1-HB→4-HB transfer faced elevated energy barriers (Supplementary Fig. 33). Consequently, the experimentally observed suppression of proton transfer by AMCs can be attributed to the cation-dependent restructuring of interfacial water molecules.

AMC-modulated *OOH binding

The kinetics of *OOH intermediate formation, desorption, and dissociation are the key to governing the 2e− ORR selectivity toward high-rate H2O2 electrosynthesis53-55. To elucidate the origin of high 2e− pathway selectivity over Co-CNT electrocatalysts under varying AMC dosages, different types of reaction models with respect to pure H+, H+ + Li+, H+ + Na+, H+ + K+, and H+ + Cs+ media were constructed. Upon evaluating various adsorption sites for key ORR intermediates (*OOH, *O, and *OH), the most stable configurations were identified (Fig. 4a; Supplementary Figs. 34–39). The reaction initiated with proton-coupled electron transfer (PCET) to form *OOH adsorbed on Co sites. The negative free energy change implied the thermodynamic favorability of *OOH formation. Subsequently, *OOH could either undergo further hydrogenation via the 4e− pathway to form *O and H2O, or desorb as H2O2 via the 2e− pathway.

As depicted in Fig. 4b, at 0.7 V vs. RHE (the equilibrium potential of H2O2/H2O), the energy barrier for the conversion of O2 to H2O2 in the Cs+-dosed system remained the lowest (–1.05 eV), reaching nearly half of that observed in the bare acid system (–2.07 eV). A comparison of the energy barriers across different AMC systems revealed a progressive decrease (Li+: –1.33 eV, Na+: –1.22 eV, K+: –1.16 eV). This trend indicates that the H2O2 generation required overcoming a higher energy barrier in the pure acid system, while the incorporation of AMCs mitigated this issue by reducing the transport of protons to the surface. Furthermore, H2O2 production is progressively facilitated in the order Li+ < Na+ < K+ < Cs+, in close agreement with experimental observations.

The tendency of the *OOH intermediate to desorb/dissociate is equally crucial to influence H2O2 production1,56,57. In bare acidic systems, excessive protons promoted dissociation of adsorbed *OOH intermediates. The energy barrier for *O formation (1.58 eV) was lower than that for H2O2 desorption (2.07 eV), favoring 4e− ORR pathway. In contrast, AMC-dosed systems exhibited substantially higher energy barrier for the *OOH → *O step (1.05 eV with Cs+ compared to 2.49 eV in bare acid), demonstrating the presence of AMCs suppressed the dissociation of *OOH by limiting the proton availability at the surface. This was a common effect across all AMC-incorporated systems, rendering 2e− pathway and promoting H2O2 formation. Notably, the Cs+-dosed system showed the pronounced reduction in the H2O2 formation barrier. The computational trend in AMC efficacy (Li+ < Na+ < K+ < Cs+) was in good agreement with our experimental measurements in ORR selectivity (Figs. 4c and 4d).

Electrochemical analysis of AMC impact

The acidic ORR performance of our catalyst was evaluated using a rotating ring-disk electrode (RRDE) system in O2-saturated 0.5 M H2SO4 containing 0.3 M AMC additive (Fig. 5a). In this sense, A2SO4 (A: AMC) was deliberately employed. Fig. 5b displays LSV curves collected at a rotation speed of 1600 rpm. The platinum ring electrode was polarized at 1.2 VRHE to oxidize H2O2 generated on the disk electrode, enabling quantitative analysis of the ORR products. The addition of AMCs significantly delayed the onset potential and lowered the disk current density as compared to the bare electrolyte scenario, which was likely to originate from hindered proton transfer in impeding the HER or 4e− ORR pathways occurring at the catalyst surface10,58. The diffusion-limited ring current density, strongly correlated with the 2e− ORR pathway, was obviously enhanced upon AMC addition. The introduction of AMCs managed to restructure the H-bond network and interfacial water configuration, thereby reducing the proton availability at the reaction interface. This effectively suppressed the dissociation of oxygen reduction intermediates, favoring their desorption and augmenting H2O2 selectivity. The electron transfer numbers (n) derived from LSV analysis confirmed that the Cs+-dosed system exhibited the highest H2O2 selectivity (reaching 80%) and lowest n value (2.5) across a broad potential range of 0-0.6 VRHE (Fig. 5c; Supplementary Fig. 40). We observed that the 2e− ORR selectivity in acid followed a monotonically increasing trend from Li+ to Cs+. Our catalyst showcased a competitive turnover frequency (TOF) value of 0.51 s−1 at 0.5 VRHE, which compared favorably with reported electrocatalysts59,60. In addition, Tafel slope plots derived from Koutecky-Levich diffusion equation manifested that the Cs+ system exhibited the fastest ORR kinetics (Fig. 5d; Supplementary Fig. 41)61,62.

In addition, electrochemical double-layer capacitance measurements demonstrate that the catalyst exhibited the highest double-layer capacitance in Cs+-dosed electrolyte, suggesting an optimized interfacial environment that promotes reaction activity and electron transfer (Supplementary Fig. 42)63,64. RRDE tests in N2-saturated 0.5 M H2SO4 containing 10-50 mM H2O2 showed significantly suppressed H2O2 reduction currents in AMC-dosed electrolytes65. The Cs+ system yielded the weakest current response for H2O2RR, particularly at high H2O2 concentrations (Fig. 5e; Supplementary Fig. 43). This effectively mitigates the thorny issue of excessive H2O2 reduction caused by high proton concentrations in acidic media. The suppression efficacy of AMCs on H2O2RR followed the trend of Li+ < Na+ < K+ < Cs+ (Fig. 5f). Moreover, in situ Raman spectroscopy was employed to monitor interactions between oxygen intermediates and the catalyst in both bare and Cs+-modified acidic electrolyte. Within the applied potential range of 0 to 0.9 VRHE, characteristic Raman peaks emerged at 983 and 1051 cm−1 in both systems (Fig. 5g), corresponding to *OOH and *O2 species, respectively14. Comparative analysis revealed the *OOH signal intensity was markedly enhanced in the Cs+-containing electrolyte relative to the bare acid counterpart, suggesting the preferential stabilization of the *OOH intermediate in the presence of Cs+. Taken together, these findings corroborate that AMC dosage could impede the proton transfer and accordingly promote the 2e− ORR selectivity for boosted H2O2 production in acid.

H2O2 generation and utilization enabled by AMCs

We first deposited the Co-CNT catalyst at a loading of 0.1 mg cm−2 onto a gas diffusion electrode (Supplementary Figs. 44 and 45). With the AMC promotor in hand, we then employed 0.5 M H2SO4 + 0.3 M A2SO4 as the flowing electrolyte and Pt as the counter electrode. We managed to assemble a three-phase flow cell reactor to evaluate the impact of AMCs in practical H2O2 production (Supplementary Figs. 46 and 47). Under a constant-current operation at 50 mA cm−2, Cs+ dosage demonstrated enhanced promotion effects relative to other cations. Both H2O2 production rate and Faradaic efficiency followed the trend of Li+ < Na+ < K+ < Cs+ (Fig. 6a; Supplementary Figs. 48–50). Notably, the H2O2 production rate increased substantially when elevating the current densities, particularly at industrially-relevant levels (>300 mA cm−2). As shown in Fig. 6b, at 500 mA cm−2, the production rate reached 9.2 mol g−1 h−1. We further assessed the production efficiency and operational stability. Under a constant-current electrolysis at 300 mA cm−2 in 0.5 M H2SO4/0.3 M Cs+ electrolyte (50 mL), H2O2 rapidly accumulated to 5100 mg L−1 within 6 h, outperforming 2000 mg L−1 attained in bare H2SO4 electrolyte (Fig. 6c). Continuously stable H2O2 generation was maintained for 140 h at 50 mA cm−2 (electrolyte flow rate: 6 mL min−1), yielding a consistent output concentration of ~5200 mg L−1 (Fig. 6d). These results underscore the optimized performance of the Co-CNT catalyst for electrochemical H2O2 synthesis in AMC-dosed acidic media, competing the state-of-the-art systems (Supplementary Table 3). Note that analogous effect was also observed under identical reaction conditions when using Fe-CNT or Ni-CNT catalysts (Supplementary Fig. 51).

The on-site generation of H2O2 in our system is highly attractive for targeted applications such as wastewater treatment. Our foregoing evaluation confirmed that dosed AMC concentrations ranging from 0.1–0.5 M would enhance H2O2 electrosynthesis, with core investigation at 0.3 M that could be encountered in industrial effluents such as reverse osmosis concentrates. Under such conditions, endogenous ions serve as inherent promoters, eliminating the external additives. Specifically, the electro-synthesized H2O2 reacts with supplemented Fe2+ via an electro-Fenton process, producing hydroxyl radicals (•OH)66-68. We selected model contaminant species spanning (i) broad-spectrum dyes (rhodamine B, methyl orange, methylene blue), (ii) aromatic pollutants (phenol, 4-chlorophenol, 4-nitrophenol) and (iii) sulfonamide contaminants (sulfamethoxazole, sulfadiazine, sulfa pyridine) (Supplementary Figs. 52–54). Our protocol achieved high removal efficiencies (>92% within 30 min) across all contaminant types (Fig. 6e). As for methylene blue exhibiting rapid degradation (>95% removal in 5 min), time-resolved UV-vis spectra confirmed progressive decay of featured absorption peaks at 664 nm (Fig. 6f; Supplementary Fig. 55), validating efficient ·OH-mediated mineralization. Furthermore, the electro-synthesized H2O2 exhibited effective bactericidal properties. Both E. coli and S. aureus cultures achieved a 100% sterilization rate within 90 min of exposure (Figs. 6g–i).

Methods
Materials. Cobalt(II) nitrate hexahydrate (Co(NO3)2·6H2O, 99 wt%), melamine (C3H6N6, 98 wt%), cyanuric acid (C3H3N3O3, 98 wt%), and potassium titanyl oxalate (K2TiO(C2O4)2 or C4H2K2TiO, 99 wt%) were purchased from Shanghai Macklin Biochemical Co., Ltd. Potassium hexacyanoferrate(III) (K3[Fe(CN)6], 99 wt%) was obtained from Aladdin Reagent Co., Ltd. A 5 wt% Nafion 117 solution was supplied by DuPont Co., Ltd. Concentrated sulfuric acid (H2SO4, 98 wt%), nitric acid (HNO3, 68 wt%), and ethanol (C2H5OH) were purchased from Sinopharm Chemical Reagent Co., Ltd. Lithium sulfate (Li2SO4, 99.5 wt%), sodium sulfate (Na2SO4, 99.5 wt%), potassium sulfate (K2SO4, 99.5 wt%), and cesium sulfate (Cs2SO4, 99.5 wt%) were acquired from Xi'an Yizhichen Biotechnology Co., Ltd. All chemicals were used as received without further purification. Carboxyl-functionalized carbon nanotubes (OCNTs) were obtained from Beijing Boyu Gaoke New Material Technology Co., Ltd. These CNTs were synthesized via chemical vapor deposition. Ultrapure water (resistivity: 18.25 MΩ cm−1) used throughout the experiments was purified using a Milli-Q system. The Au microdisk electrode was obtained from CH Instruments, Inc.

Synthesis of Co-CNT. Oxygen-functionalized carbon nanotubes anchored with cobalt clusters (denoted as Co-CNT) were synthesized via a sequential chemical impregnation, pyrolysis, and acid etching process. Initially, OCNTs were immersed in a mixed acid solution of concentrated H2SO4 (75 mL) and HNO3 (25 mL) and ultrasonicated at room temperature (20 ˚C) for 30 min to enhance dispersion. The suspension was subsequently refluxed at 80 °C in an oil bath for 1 h. After cooling to room temperature (20 ˚C), the sample was isolated by centrifugation at 6000 rpm, thoroughly washed with ethanol, and finally dried overnight in a vacuum oven at 60 °C. For cobalt precursor impregnation, cobalt (II) nitrate hexahydrate (Co(NO3)2·6H2O, 2 mM), cyanuric acid (1.02 g), and melamine (1 g) were dissolved in a 100 mL 1:1 (v/v) ethanol/ultrapure water mixture. This solution was ultrasonicated for 30 min followed by continuous stirring for 24 h. The resulting solid was collected by centrifugation at 6000 rpm, washed three times with ethanol, and vacuum-dried. The dried precursor powder was fine ground and then subjected to pyrolysis in a tube furnace under a continuous N2 flow at 900 °C for 3 h, employing a heating rate of 5 °C min−1.

Characterizations. Powder X-ray diffraction (XRD) patterns were acquired on a Bruker D8 Advance diffractometer using Cu Kα radiation (λ = 1.5418 Å) operated at 40 kV and 40 mA. High-angle annular dark-field scanning transmission electron microscopy (HAADF-STEM) images were obtained using a JEOL JEM-2100F field-emission electron microscope operating at an accelerating voltage of 200 kV. Elemental composition analysis was performed by inductively coupled plasma atomic emission spectroscopy (ICP-AES) on a Thermo Fisher Scientific iCAP 6300 spectrometer. X-ray photoelectron spectroscopy (XPS) measurements were conducted on a Thermo Scientific ESCALAB 250Xi spectrometer with monochromatic Al Kα X-ray radiation (1486.6 eV). Fourier transform infrared (FT-IR) spectra of the catalysts were acquired using a Nicolet Nexus 670 spectrometer employing the KBr pellet technique.

In situ ATR-SEIRAS. Electrochemical ATR-SEIRAS measurements were performed using a Bruker INVENIO-R FTIR spectrometer equipped with a liquid-nitrogen-cooled mercury cadmium telluride (MCT) detector. A conventional three-electrode configuration was employed for electrochemical control1. In situ SEIRAS measurements were conducted in electrolytes with pH 0, comprising 0.5 M H2SO4 (Sigma-Aldrich, 70 wt%, 99.999% trace metals basis) and 0.1–0.5 M sulfate solutions containing Li+, Na+, K+, or Cs+ (Li2SO4, Sigma-Aldrich, 99.99%; Na2SO4, Sigma-Aldrich, 99.9%; K2SO4, Sigma-Aldrich, 99.99%; Cs2SO4, Sigma-Aldrich, 99.9%). A platinum foil and an Ag/AgCl electrode served as the counter electrode and reference electrode, respectively. Unpolarized infrared radiation was focused onto the interface at an incident angle of 60° using an external reflection accessory (Elema) and the reflected radiation was detected. Prior to ATR-SEIRAS experiments, the gold-coated Si prism surface was electrochemically cleaned by cycling the potential between 0.05 V and 1.1 V vs. the reversible hydrogen electrode (RHE). Time-resolved spectra were recorded within the working electrode potential window at a resolution of 8 cm−1, co-adding 44 scans per spectrum. The reference spectrum was collected at 1.1 V vs. RHE69.

Electrochemical Measurements. Prior to each measurement campaign, the Ag/AgCl reference electrode was calibrated against a freshly prepared saturated calomel electrode (SCE, 0.241 V vs. SHE at 25 °C) to ensure potential accuracy and reproducibility. Both electrodes were immersed in a thermostatted 3.0 M KCl solution maintained at 25.0 ± 0.1 °C using a water circulator. The open-circuit potential difference was measured using a high-impedance voltmeter after stabilization for 10 min, with data recorded at 1 s intervals over a 5 min period. The potential of the Ag/AgCl electrode versus the standard hydrogen electrode (SHE) was calculated as EAg/AgCl vs. SHE = ESCE vs. SHE + ΔE, where ΔE is the measured potential of the Ag/AgCl electrode relative to the SCE. Only electrodes exhibiting a potential within ±2 mV of the theoretical value (e.g., +0.197 V vs. SHE for saturated KCl internal filling) and drift less than 0.5 mV over the measurement period were deemed acceptable for subsequent electrochemical experiments. All potentials reported herein have been corrected to the SHE scale based on this calibration.

The apparent radius of the microelectrode was determined using the steady-state limiting current (Ilim) according to the following equation:
Ilim = 4nFDCr (1)
where n represents the number of electrons transferred per molecule, F is the Faraday constant (96485 C mol−1), D denotes the diffusion coefficient of the reactant (cm2 s−1), C is the bulk concentration of the reactant (mol cm−3), and r signifies the true geometric radius of the microdisk electrode. Specifically, Ilim was measured as the steady-state limiting current for the reduction of 10 mM [Ru(NH3)6]Cl3 in a 0.1 M K2SO4 solution. For this well-established redox couple, the parameters are defined as follows: n=1 (one-electron transfer), F=96,485 C mol−1, D=5.71×10−6 cm2 s−1 (diffusion coefficient of [Ru(NH3)6]3+), and C=10 mM (bulk concentration of [Ru(NH3)6]Cl3). All electrochemical measurements were conducted at ambient temperature (25 ± 1 °C) using a CHI 760E electrochemical workstation.

In a custom three-electrode system, linear sweep voltammograms of the Au microelectrode were obtained by performing hydrogen evolution experiments in 0.5 M H2SO4 solutions containing x M A+ (where x = 0.1–0.5; A+ = Li+, Na+, K+, Cs+) at a scan rate of 0.02 V s−1.

Rotating disk electrode (RDE) tests were conducted using a Pine Research Instrumentation MSR rotator coupled to a CHI-760D bipotentiostat. Polarization curves were recorded on a polished Au RDE (5 mm diameter) in hydrogen-saturated 0.1 M H2SO4 + 0.3 M K+ solution at 0.02 V s−1.

Cyclic voltammograms for [Ru(NH3)6]Cl3 reduction were acquired in a custom three-electrode system using aqueous solutions containing 10 mM [Ru(NH3)6]Cl3 and x M K+ (x = 0.1-0.3). Scans were performed from 0.7 V to -0.2 VSHE at 0.02 V s−1 to determine the limiting current.

Electrocatalytic ORR performance was evaluated using a CHI 760e electrochemical workstation (CH Instruments, Inc., Shanghai, China) with a standard three-electrode configuration and an RRDE (Pine Instruments). The system comprised an Ag/AgCl reference electrode, Pt foil counter electrode, and an RRDE working electrode with a glassy carbon disk (diameter: 5.61 mm, area: 0.247 cm2) and Pt ring (inner diameter: 6.25 mm, outer diameter: 7.92 mm, area: 0.186 cm2). All electrodes were mechanically polished and rinsed with ultrapure water prior to measurements.

Catalyst ink was prepared by dispersing 5 mg synthesized catalyst in a mixture of 480 μL isopropanol, 480 μL ultrapure water, and 20 μL (5.0 wt%) Nafion solution, followed by 2 h sonication. A 10 μL aliquot of the homogeneous ink was drop-cast onto the disk electrode to achieve a catalyst loading of 0.2 mg cm−2. All electrochemical potentials were converted to the reversible hydrogen electrode (RHE) scale using:
E (RHE) = E (Ag/AgCl) + 0.0592 × pH + 0.197 (2)
All electrochemical tests were conducted at room temperature (20 ˚C) in 0.5 M H2SO4 electrolyte thoroughly saturated with N2 or O2 prior to measurements.

Before electrochemical testing, all catalysts were electrochemically activated via cyclic voltammetry (CV) at 50 mV s−1 until stable voltammograms were obtained. Linear sweep voltammetry (LSV) was performed under O2-saturated conditions at a scan rate of 10 mV s−1 with an RRDE rotation speed of 1600 rpm, while the ring electrode potential was held constant at 1.2 VRHE to monitor H2O2 production. H2O2 selectivity and electron transfer number (n) for the RRDE were calculated using the following equations:
Selectivity (%) = 200 × (Ir / N) / (Id + Ir / N) (3)
Electron transfer number (n) = 4 × Id / (Id + Ir / N) (4)
where Ir and Id denote the ring and disk currents, respectively, and N represents the collection efficiency of the ring electrode (0.37). Catalyst stability was evaluated via chronoamperometry at a constant potential of 1.2 VRHE.

To probe active sites, poisoning experiments were conducted by adding 10 mM KSCN to the 0.5 M H2SO4 electrolyte. Linear sweep voltammetry (LSV) was performed under identical configurations as described above. For hydrogen peroxide reduction reaction (H2O2RR) tests, electrolytes comprised N2-saturated 0.5 M H2SO4 containing 10-50 mM H2O2 (10, 20, 30, 40, 50 mM).

Assuming metallic Co serves as the active site and all metal sites participate in the reaction, the TOF is calculated as:
TOF = I Mmetal / n F mcatalyst wmetal (5)
The current density I (unit: A) for H2O2 generation can be obtained by multiplying the disk current with the Faraday efficiency of H2O2. Mmetal (unit: g/mol) represents the molar mass of the transition metal. The parameter n = 2 denotes the number of electrons transferred during the H2O2 production process. F stands for the Faraday constant, while mcatalyst (unit: g) indicates the mass of the catalyst loaded on the electrode. The mass fraction of metal in the catalyst, wmetal is determined through ICP analysis. This approach, which presumed all Co atoms in the clusters are active, yielded an overestimated count of apparent active site. This, in turn, resulted in a calculated TOF that systematically underestimated the actual catalytic activity.

The Tafel slope is calculated by the Tafel equation:
η = b log (j / j0) (6)
η represents overpotential, b is Tafel slope, j is current density, and j0 is exchange current density.

H2O2 Concentration Quantification. H2O2 content was determined via potassium titanyl oxalate (K2TiO(C2O4)2) titration. In this method, titanium ions react with H2O2 under acidic conditions to form a stable orange peroxotitanium complex, with color intensity proportional to H2O2 concentration. Experimentally, H2O2 solutions of known concentration were mixed with a solution containing 3 M H2SO4 and 0.1 M potassium titanyl oxalate. The absorbance of the mixture was measured at 400 nm using UV-Vis spectroscopy (20 ˚C). A standard calibration curve correlating H2O2 concentration (C) and absorbance (a) was established:
a = 0.2028C+0.0101 (R2=0.9991) (7)
The Faradaic efficiency (FE) of H2O2 generation was calculated via the following equation:
FE = 2FCV/It (8)
where F represents the Faraday constant (96485 C mol−1), C refers to the H2O2 concentration (mol L−1) in the electrolyte, V indicates the volume of electrolyte (L) in the cell, and Q is the total charge amount (C) consumed. I represents the current, and t refers to the running time. Q was determined by the integration area of I over t in this chronoamperometry experiment.

Flow Cell H2O2 Production. Practical H2O2 electrosynthesis performance was evaluated in a two-compartment flow cell with a three-electrode configuration. An Ag/AgCl electrode served as the reference electrode and a Pt foil as the counter electrode. Catalyst ink was uniformly drop-cast onto a 2 × 2 cm2 carbon paper to form a catalyst layer, creating a gas diffusion electrode (GDE) with a loading of 0.1 mg·cm−2. The GDE was assembled as the oxygen reduction reaction (ORR) cathode in a three-phase flow cell: the catalyst layer faced the cathode chamber while the opposite side was exposed to O2 flow. The cell comprised proton exchange membrane (Nafion 117, size: 2.5 cm × 2.5 cm, thickness: 183 μm)-separated cathodic and anodic compartments (2 × 4 mL), with a Pt plate anode and an Ag/AgCl reference electrode fixed in the cathode chamber. The proton exchange membrane was not subjected to any pretreatment.

The entire experimental process was carried out at room temperature (20 ˚C). The corresponding electrolyte was prepared by adding a certain amount of solid salt to the sulfuric acid solution. Note that a fresh electrolyte was prepared 30 min prior to each test.

Fenton Reaction for Pollutant Degradation. In situ-generated H2O2 was coupled with the classical Fenton reaction to achieve efficient pollutant degradation. The reaction scheme is expressed as:
2Fe2+ + H2O2 → 2Fe3+ + OH-+ ·OH (9)
This reaction was implemented in the flow cell by introducing both the pollutant solution and Fe2+ into the catholyte compartment. Through this process, hydroxyl radicals (·OH) generated from the reaction between in situ-produced H2O2 and Fe2+ continuously degrade pollutants in the solution, enabling efficient contaminant removal.

MD Simulations. The MD simulations were carried out using GROMACS 2020.3 software package with the OPLS-AA force fields.70-72 To simulate the experiment solutions, 20 hydronium ions, 10 sulfate anions, and around 1080 water molecules were inserted randomly into cubic boxes with side length of 32 Å for 0.5 M H2SO4 without AMCs. For 0.5 M H2SO4 with 0.15 M A2SO4 (A=Li, Na, K, or Cs), 6 AMCs and 3 more sulfate anions were added into the boxes. The steepest descent method was then used to perform 1000 steps of minimization. The models were pre-equilibrated by NVT ensemble simulations for 200 ps. This was followed by extensive equilibration by 2 ns NPT and 2 ns NVT. The timestep was 1 fs in all simulations. The Berendsen method was used to control both temperatures and pressures at 298 K and 1.0 atm.73 The Columbic integrations were treated with the particle mesh Ewald (PME) method.74 These models were used to analysis the number of HBs and the RDFs of AMCs verse water39,75-79.

To study the connectivity of HB networks, we also built cubic boxes (7 Å in length) filled with 50 water molecules, 2 hydronium ions and a sulfate anion. A hydronium ion was replaced by an AMC for solutions with AMCs. Same relaxations as described above were applied to these systems. PIMD simulations were then conducted using built-in PINT module in CP2K software package.80 8 beads are adopted to sample the quantum nuclei. The PIMD trajectories were used for HB networks analyses by ChemNetworks software.81 An O···H bond distance cutoff of 2.5 Å and a bond angle cutoff of 145° were employed.

DFT Calculations. The free energy profiles of the ORR pathway were investigated by DFT methods using the Vienna ab initio simulation package (VASP) version 5.4.4.82-85 The core−valence electron interaction was described by the projector augmented wave (PAW) approach with a plane-wave cutoff energy of 400 eV.86 The electronic exchange−correlation effect was represented by the Perdew−Burke−Ernzerhof (PBE) functional within the generalized gradient approximation (GGA).87,88 The Grimme D3 correction was added to describe the London dispersion integrations.89 All calculations are spin-polarized. The entropy correction and the zero-point energy correction at 298.15 K were taken into consideration for the Gibbs free energy. The 2e- reaction pathway for H2O2 production is calculated as following:
O2 (g) + 2H+ + 2e- + * → *OOH+ H++ e- (10)
*OOH + H+ + e- → H2O2 (g) + * (11)
where * represents the catalyst, and *OOH donates the intermediate species adsorbed on the catalyst. For the Gibbs free energy, both entropy and the zero-point energy correction were considered at 298.15 K. The Gibbs free energy differences were calculated:
ΔG (1) = ΔGOOH – 4.92 (eV) (12)
ΔG (2) = 3.56 – ΔGOOH (eV) (13)
where ΔGOOH represents the adsorption energy of molecular OOH.

We applied a constant potential (+0.15 V vs. SHE, consistent with experiments) to the ORR simulations by CP-VASP.90,91 The targeted Fermi level referenced to the electrolyte was set to –4.75 V with a convergence criterion of 0.01 V for the electrolyte-referenced Fermi level. The solvation model during in calculations was considered by using the VASPsol++ framework implemented in VASP.92

Each AMC was put on the surface of Co-CNT and the adsorption energies were calculated by DFT:
Ead = Etotal – ECo-CNT – EAMC (14)
where Etotal is the total energy of an AMC adsorbed on the catalyst, while ECo-CNT and EAMC represent the energies of Co-CNT and AMC, separately. Positive-charged calculations were performed by manually setting the number of valence electrons to consider the positive charge state of AMCs.

To study the proton transfer process, we constructed 3 clusters of 8, 8 and 10 water molecules with an extra proton, for 1-HB to 4-HB, 2-HB to 4-HB and 4-HB to 4-HB, respectively. Flexible scans were performed for the O···H···O bond between the proton donor and the proton acceptor using Jaguar program at the BLYP/6-31G* level.93 The proton transfer energy barrier was calculated as the energy difference between the intermediate state with highest energy and the initial structure. All the files of optimized computational models are provided in Supplementary Data 1.

Due to the inherent limitations of simulations in capturing the full complexity of real catalytic environments, these results serve as theoretical references requiring experimental validation.

Peer review rounds
[{"round": 1, "reviewer_comments": {"reviewer_1": "The revision has addressed most of my previous concerns and the manuscript is now close to acceptance. I believe the work can be accepted after the authors clarify the following minor points to improve transparency and reproducibility.\n1. In the rebuttal, the authors state that adsorption energies of alkali metal cations (AMCs) on the catalyst surface were calculated to exclude direct surface binding. It would be helpful to explicitly clarify whether the positive charge state of the cations (e.g., Li⁺/Na⁺/K⁺) was considered in these calculations, and how charge compensation was handled in the computational model (e.g., background charge, counter ions, solvation treatment, etc.), since this may influence adsorption energetics and interfacial electrostatics.\n2. The authors performed constant-potential DFT calculations at 0.15 V vs. SHE to account for electrochemical conditions. While this value is described as being consistent with experimental conditions, it would strengthen the manuscript to provide a clearer electrochemical/theoretical rationale for selecting this potential, or to briefly discuss how representative it is within the relevant ORR operating window and whether the conclusions are expected to remain robust over a reasonable potential range.\n3. In the Methods section, particularly in the “DFT calculations” subsection near the end of the manuscript, there are currently no references cited for the computational methods and models employed. Please add appropriate methodological references (e.g., DFT setup, electrochemical/solvation model, constant-potential method) to improve completeness and reproducibility.", "reviewer_2": "Dear Authors, Dear Editors,\nI feel like all my remarks have been answered appropriately.\nBest regards.", "reviewer_3": "I am satisfied by the additional work the co-authors have done to elevate their contribute to Nature Comms. level. I recommend it's acceptance."}, "author_response": "Response to Reviewers’ Comments (NCOMMS-25-75578)\nReviewer #1 (Remarks to the Author):\nThe study investigates how alkali metal cations (AMCs: Li⁺, Na⁺, K⁺, Cs⁺) enhance H₂O₂ production\nduring the oxygen reduction reaction (ORR). The authors combine a wide range of experimental\nand computational techniques, including physicochemical analyses (XRD, TEM, ICP, XPS, FT-IR),\nelectrochemical characterization, in situ ATR-SEIRAS, MD and DFT simulations, as well as a flow-\ncell demonstration coupled with a Fenton reaction for pollutant degradation.\nWhile the conventional interpretation would be that larger cations simply block or hinder the 4-\nelectron ORR pathway at the active sites or the outer Helmholtz plane, the authors propose a more\nsophisticated mechanism; water network, Eigen-Zundel interconversion, and proton hopping. The\nresults provide a coherent and convincing narrative that the altered interfacial water structure\ngoverns the ion-specific enhancement of H2O2 selectivity. The mechanistic consistency across\nexperiments and simulations is a notable strength of the work, but there are several comments to\nfurther develop the research.\nResponse: We are grateful for your positive feedback. All requested revisions based on your\ninstructive comments have been carried out.\nSpecific comment:\n1. It would be beneficial to include evidence (e.g., from FT-IR, XPS, or DFT adsorption-energy\ncalculations) clarifying whether the introduced AMCs interact with or adsorb onto the catalyst\nsurface, or remain fully hydrated near the interface.\nReply 1: Thank you for your valuable advice. To directly address the state of the introduced AMCs,\nwe performed supplementary DFT calculations to evaluate the adsorption energies of anhydrous\nAMCs on the Co-CNT catalyst surface. The computed adsorption energies for all AMCs (Li⁺, Na⁺,\nK⁺, Cs⁺) are positive, indicating that these cations are not thermodynamically favorable to adsorb\nonto the catalyst surface. Instead, they remain fully hydrated and reside near the electrode–\nelectrolyte interface, as illustrated in Figure 4a. This finding confirms that the promotional effect of\nAMCs originates from their hydrated state at the interface, rather than through direct adsorption\nonto the catalytic sites.\nRevisions to Manuscript:\nSupplementary Fig. 34 | Optimized structures of anhydrous AMCs on Co-CNT and its adsorption\nenergy values.\n2. The MD and DFT models appear to describe neutral interfaces. In reality, the electrochemical\nreactions occur under an applied potential, especially in the presence of charged species like Cs⁺.\nA discussion of possible approaches—or the limitations of the current setup—to model potential-\ncontrolled interfaces would strengthen the computational section.\nReply 2: Thank you for the constructive comment in improving the quality of our manuscript. We\nfully agree that the calculations of electrochemical reactions under an applied potential could\ndescribe the nature of electrochemical reaction mechanism in a more accurate manner.\nIn response to your suggestion, we have carried out additional constant-potential DFT\ncalculations using the CP-VASP method (+0.15 V vs. SHE, consistent with key experimental\nconditions) to explicitly simulate the electrode potential during the ORR1,2. These calculations\nconfirm that the high selectivity toward the 2e⁻ pathway could be maintained in the presence of\nAMCs under a representative applied potential (Fig. R1). The results further support our central\nconclusion that the cation-induced restructuring of the interfacial water network, rather than a\nchange in the fundamental reaction mechanism at the active site, is the dominant promotional effect.\nThe reactions remain a high 2e⁻ ORR selectivity with AMCs.\nWhile our present MD simulations employed a charge-neutral setup to efficiently explore the cation-\nspecific perturbation of the hydrogen-bond network, we acknowledge that explicitly modeling the\nfull electrochemical double layer under bias remains a significant computational challenge. Future\nwork will benefit from more advanced simulation techniques to fully couple the interfacial water\nstructure with the electrostatic field. Nevertheless, the strong consistency between our potential-\ninclusive DFT results, the potential-dependent in situ spectroscopic data, and the observed\nelectrochemical trends provides a coherent and convincing mechanistic picture.\nFig. R1. Free energy profiles of the ORR pathways on Co-CNT catalysts under 0.15 V vs. SHE.\n3. The flow-cell experiment successfully illustrates the practical applicability of H₂O₂ generation,\nbut it contributes little to the mechanistic understanding, which is the core focus of this paper. This\nsection could be presented more clearly as an application-oriented addendum rather than\nmechanistic evidence.\nReply 3: We sincerely thank you for raising this important point regarding the role of flow-cell\nexperiments in the context of our manuscript. We fully agree that the flow-cell section primarily\nillustrates the practical applicability of the H O generation strategy. In revising the manuscript,\n2 2\nwe have refined the presentation of this part to more clearly distinguish between mechanistic\ninvestigation and application demonstration, as suggested.\nAt the same time, the flow-cell performance data (Figure 6) demonstrates that the Cs+-promoted\nsystem exhibits not only the highest H O production rate (9.2 mol g⁻1 h⁻1 at 500 mA cm⁻2) but also\n2 2\nexceptional operational stability over 140 h, which directly reflects the robustness of the interfacial\nwater restructuring mechanism proposed in our work.\nThe observed cation-dependent performance trend (Li+ < Na+ < K+ < Cs+) in the flow-cell mirrors\nthe trend obtained from fundamental microelectrode, spectroscopic, and computational analyses,\nthereby corroborating the consistency of the mechanism under industrially relevant conditions.\nTo better frame this section, we have revised the sub-heading to “H O generation and utilization\n2 2\nenabled by AMCs” which more accurately presents the flow-cell experiments as an\napplication-oriented extension that concurrently verifies the practical implications of the proposed\nmechanism. This format aligns with similar high-impact studies that include flow-cell\ndemonstrations as a logical endpoint of mechanistic discovery3,4.\nRevisions to Manuscript:\n“H O generation and utilization enabled by AMCs” in Page 12.\n2 2\nReviewer #2 (Remarks to the Author):\nI have read the manuscript “Differentiating interfacial water structures via alkali metal cation 1\npromotor for H2O2 electrosynthesis in acid”. Overall, the manuscript presents interesting findings\nand is addressing the relevant topic. However, there are some issues I feel need to be addressed\nbefore publication.\nResponse: We are grateful for your positive and constructive feedback. All requested revisions\nbased on your professional comments have been carried out.\nSpecific comment:\n1. “Employing Co-CNT as a model catalyst, we observe that the 2e. ORR selectivity in acid follows\na monotonically increasing trend from to which is not revealed by prior studies (Figure 1).”\nAnd similar claims authors make throughout the manuscript, while strictly speaking true for this\nspecific catalyst, are not generally applicable. The exact same cation trend for the reaction has been\nrecently observed, e.g., for facets of graphitic carbon https://doi.org/10.1021/acscatal.4c04734\nAdditionally, I think the authors should comment on this and similar literature findings.\nIn fact, I find that one of the biggest weaknesses of the manuscript currently is the very limited\nreferral to literature data for oxygen reduction to hydrogen peroxide. For instance, Metal-CNT\n(https://doi.org/10.1021/cr900136g,https://doi.org/10.1002/anie.201509241, https://doi.org/10.10\n16/j.jcat.2005.11.024), discussions about 2-electron ORR\n(https://doi.org/10.1021/acscatal.8b00217 ), and cation (https://doi.org/10.1021/acscatal.4c04734)\nand electrolyte effects in general (\nhttps://doi.org/10.1002/cssc.200800257, https://www.sciencedirect.com/science/article/abs/pii/S00\n21951704001952 –33,127,131https://doi.org/10.1016/S0926-3373(02)00232-1). This is by no\nmeans an exhaustive list\nReply 1: Thank you for this critical comment. We fully agree that a comprehensive discussion of\nthe existing literature is essential to enhance the quality of our work. Based on your suggestions and\nfurther analysis, we have expanded the discussion to acknowledge and refer to the important prior\nstudies, including cation effect on acidic 2e⁻ ORR, the employment of different catalysts, as well as\ncation and electrolyte effects in general. The relating revision has been made in the rephrased\nIntroduction section, with the supplementary incorporation of advised references (New Ref. [3–\n8][30][34]).\nRevisions to Manuscript:\n“Hydrogen peroxide (H O ) has emerged as a green and versatile oxidant with vast application\n2 2\npotentials in wastewater treatment, organic synthesis and chemical technology1-7. Its on-site\nelectrosynthesis via the two-electron oxygen reduction reaction (2e⁻ ORR) with enriched choices of\nelectrocatalysts8.\nElectrolyte engineering, particularly the introduction of alkali metal cations (AMCs: Li+, Na+,\nK+, or Cs+)28-30, has emerged as a promising strategy to steer interfacial chemistry toward H2O2\nelectrosynthesis in acid31,32. A recent study indicated a graphitic carbon electrode exhibited an\nincrease in 2e⁻ ORR selectivity following the trend of Cs+ > K+ > Na+ > Li+, where different cations\nwould interfere in the *OOH binding and influence the local electric field34” in Page 3-4.\nAdded references:\non Fenton’s Reaction Chemistry. Chem. Rev. 109, 6570-6631 (2009).\n2 2 2 2 2\nby palladium in an aqueous medium. J. Catal. 225, 249-255 (2004).\nHydrogen Peroxide using Supported Gold–Palladium Catalysts. ChemSusChem 2, 575-580\n(2009).\nselective oxidation of H by O to H O over supported Pd catalysts in an aqueous medium.\n2 2 2 2\nJ. Catal. 238, 28-38 (2006).\nof hydrogen peroxide from hydrogen and oxygen. Appl. Catal., B 42, 203-211 (2003).\nSelective Electrochemical Reactions. Angew. Chem. Int. Ed. 55, 2058-2062 (2016).\nSurface-Specific Spectroscopic Survey. J. Chem. Phys. eprint arXiv:2508.15422 (2025).\nSurfaces. ACS Energy Lett. 9, 1331-1338 (2024).\n2 2\nthe Catalysis. ACS Catal. 8, 4064-4081 (2018).\nCarbon Electrodes: Influence of Defects, Oxygen Functional Groups, and Alkali Metals in\nthe Electrolyte. ACS Catal. 14, 17675-17689 (2024).\n2. “The varying influences of different AMCs stem from their distinct capabilities to perturb the H-\nbond network, associated with their hydration properties and coordination strengths. Impressively,\nthe elucidated mechanism is proven to be universal, as Fe-CNT and Ni-CNT catalysts exhibit\nanalogous AMC-dependent behavior”\nOnce again, this would be an opportunity to refer to and comment on existing literature data on\ncation effects on other catalysts.\nReply 2: Thanks for your kind suggestion. In the revised Introduction section, we have expanded\nour discussion on cation effects using other catalysts based on existing literatures. Our claim of\n“universality” is carefully confined in the catalyst type of our employed metal-CNT, since it has not\nbeen reported in the cation-sensitive acidic 2e⁻ ORR study. By integrating these discussions, we aim\nto present our work not as an isolated finding, but as a conceptual bridge that connects electrolyte\nengineering, interfacial water science and catalyst design for sustainable H O production.\n2 2\nRevisions to Manuscript:\n“Particularly using metal-free carbon-based electrocatalysts, Zhang et al. observed that H O\n2 2\nproduction rates can be improved by an AMC-induced shielding effect, yet were relatively\nunaffected by the AMC types33. A recent study indicated a graphitic carbon electrode exhibited an\nincrease in 2e⁻ ORR selectivity following the trend of Cs+ > K+ > Na+ > Li+, where different cations\nwould interfere in the *OOH binding and influence the local electric field34. It is interesting to note\nthat the AMC roles in acidic 2e⁻ ORR manifest inconsistent trends. In further contexts, the cation\neffect is effective across other catalyst types apart from carbonaceous materials. For instance, Cao\net al. found the H O production followed the trend of K⁺ > Na⁺ > Li⁺ ~ Cs⁺ with the employment\n2 2\nof earth-abundant TiC catalyst, which was dictated by the hydrated ionic radius of AMCs35. These\ndisparities imply that the dominant mechanism might be sensitive to specific conditions, such as\ncatalyst identity, cation concentration and operating current density. A coherent picture of how\nAMCs reshape the interfacial environment in acidic 2e⁻ ORR is still lacking, especially in the realm\nof prevailing transition metal-involved carbon electrocatalysts (e.g., metal-carbon nanotube, M-\nCNT).” in Page 3-4.\nAdded references:\nCarbon Electrodes: Influence of Defects, Oxygen Functional Groups, and Alkali Metals in\nthe Electrolyte. ACS Catal. 14, 17675-17689 (2024).\n3. Cs⁺ produced the most pronounced suppression, reaching a maximum factor of 13-fold. It is\nfurther noted that there existed a concentration-dependent variation of this suppression effect\n(Supplementary Fig. 3).\nSome parts of CVs appear to be missing in Sup. Fig 3C.\nWhile it is true that Cs+ has the strongest and Li+ the smallest effect in general. Here we do not see\nthe Cs+>K+>Na+>Li+ trend in e.g. 0.1M M2SO4.\nReply 3: Thanks for your meticulous review and valuable comments. Following these suggestions,\nwe have re-examined the data. Firstly, the complete plateau region (from –0.4 to –1.1 V) of the\nLSV profiles for certain conditions (especially at lower cation concentrations) has been presented\nin New Supplementary Fig. 4. Secondly, the AMC concentration-dependent variation of diffusion-\nlimiting current density has been updated in New Supplementary Fig. 5. In this sense, it could be\nobserved that the dosed AMC with varied concentrations (0.1–0.4 M; i.e., including 0.1 M) all\nexhibits suppression effect (Li+ < Na+ < K+ < Cs+) on proton transport. The strength of suppression\nscales monotonically with both cation type and concentration, reinforcing the proposed mechanism\nthat larger cations (Cs+) could more effectively reconstruct the interfacial water and impede proton\nhopping via H-bond network.\nRevisions to Manuscript:\n“It is further noted that there existed a concentration-dependent variation of this suppression effect\n(Supplementary Figs. 4 and 5).” in Page 5.\nSupplementary Fig. 4 | HER polarization curves of Au microelectrodes in 0.5 M H SO + x M\n2 4\nLi+/Na+/K+/Cs+ (x ranging from 0.1 to 0.4) at 0.02 V s⁻1.\nSupplementary Fig. 5 | Diffusion-limited current density for 1 M H+ and 1 M H+ + x M A+ (x\nranging from 0.1 to 0.4; A = Li, Na, K, or Cs).\n4. An optimized concentration of 0.3 M was selected, since higher dosages were found to impair\nH2O2 generation within the system (Supplementary Fig. 4).\nWhy would this be? The authors claim that H+ concentration suppression is good for 2-electron\nORR selectivity, then why would the trend flim at 0.3M?\nReply 4: Thank you for the professional suggestion on the experimentally optimized concentration\nat 0.3 M and the subsequent decline in H O production performances at higher concentrations. The\n2 2\nquestion refers to a key interfacial reaction engineering: the balance between suppressing parasitic\npathways and maintaining sufficient reaction kinetics.\nOur work demonstrates that the proton-delivery suppression by AMCs boosts 2e⁻ ORR activity.\nHowever, this does not imply that proton availability should be minimized indefinitely. The overall\nORR rate, including the desired 2e⁻ pathway, remains a proton-coupled electron transfer process.\nExcessively hindering the proton transport would hence slow down the entire reaction, leading to\nthe observed decline in H O production rate at quite high cation concentrations. The optimum at\n2 2\n~0.3 M reflects a trade-off between selectivity and activity. Below this concentration, proton\navailability remains too high, favoring 4e⁻ ORR and H O reduction. Above this concentration,\n2 2\nproton supply becomes kinetically limiting, reducing the overall ORR current and hence the H O\n2 2\ngeneration rate.\nIt is worth-noting that such an optimized concentration alters systematically with the cation identity:\nit occurs at ~0.3 M for Cs+, ~0.3 M for K+, ~0.4 M for Na+, and ~0.5 M for Li+ (New Supplementary\nFig. 6). This trend aligns precisely with each typed cation’s capability to perturb the interfacial H-\nbond network and suppress proton transport (Cs+ > K+ > Na+ > Li+). A stronger “suppressor”\nrequires a lower concentration to reach the same level of proton limitation, explaining why the peak\nperformance occurs at lower concentrations for Cs+ than for Li+.\nMore significantly, from an environmental engineering perspective, this concentration-dependent\nbehavior is meaningful. Industrial wastewater streams often contain fluctuating ionic strength; our\nsurvey indicate that even at moderate cation concentrations (≥ 0.2 M), the promoting effect remains\nrobust and follows a predictable order. This provides a practical guideline for electrolyte\nmanagement in real applications, such as using brine or high-salinity wastewater as an electrolyte\nsource for on-site H O electrosynthesis.\n2 2\nRevisions to Manuscript:\n“The optimized AMC concentration refers to a key interfacial reaction engineering of the balance\nbetween suppressing parasitic pathways and maintaining sufficient reaction kinetics. Considering\nthe trade-off between proton availability and reaction kinetics, a medium concentration of 0.3 M\nwas selected to assess the AMC concentration effect on the H O yield (Supplementary Fig. 6).” in\n2 2\nPage 5.\nSupplementary Fig. 6 | H O yield based on Co-CNT at different AMC concentrations in a flow-\n2 2\ncell reactor.\n5. Why did the authors choose to perform all measurements in sulfuric acid? Sulfuric acid is know\nto strongly interact with many surfaces, complicating interpretation. Since the authors claim that\nthe observed trends are a result of H+ availability specifically, the authors should perform\nmeasurements in HClO4 in addition with same amounts of alkaline perchlorates to observe the\ncations effects here as well.\nReply 5: We sincerely thank you for raising this important point. The choice of sulfuric acid as the\nprimary electrolyte was motivated by its widespread environmental relevance. The SO 2⁻ is\n4\namong the most abundant anions in industrial wastewater, mining effluent, and natural saline waters,\nmaking H SO a practically significant medium for evaluating electrochemical processes aimed at\n2 4\nreal-world water treatment applications.\nWe fully agree that SO 2⁻ may exhibit specific adsorption or surface interactions that could\n4\npotentially complicate the interpretation of cation effects. To rigorously decouple the role of anions\nand confirm that the observed trends originate primarily from cation-modulated proton availability,\ngiven the low solubility of HClO4, we have followed the suggestion and supplemented a\ncomplete set of control experiments using HClO with corresponding alkali perchlorates\n4\n(0.1 M HClO + 0.05 M AClO , A = Li, Na, K, or Cs). The introduction of AMCs in perchlorate\n4 4\nmedia again enhances the H O production. The cation‑dependent promotion trend follows the same\n2 2\norder: Cs⁺ > K⁺ > Na⁺ > Li⁺, in good agreement with the trend observed in SO 2⁻‑based electrolytes.\nFig. R2. The H O concentration based on Co-CNT for 0.1 M HClO and 0.1 M HClO + 0.05 M\n2 2 4 4\nAClO in a flow-cell reactor after 1 h (A = Li, Na, K, or Cs).\n4\n6. In Fig 2c, as well as Fig 5b, besides a change in the diffusion limiting current, we can also observe\na shift in the onset and halfwave potentials for both reactions. If the activity differences stem purely\nfrom H+ concentration near the surface, this should primarily cause a difference in the current\ndensities.\nIn fact, the cation concentrations used in this work are quite high. Have the authors measured the\npHs of the electrolytes and have any differences been accounted for in the calculation of the applied\npotentials? Cation effects have been observed in much more subtle concentrations of 10-50 mM,\ni.e., two orders of magnitude below what the authors are applying!\nReply 6: Thank you for this insightful comment. The shifts in both onset potential and half-wave\npotential indeed reflect a comprehensive modulation of the ORR process by alkali metal cations\n(AMCs).\nOnset potential: The onset potential (defined here as the potential at which the current density\nreaches –0.05 mA cm⁻2) corresponds to the initial activation of O . Its negative shift indicates that\n2\nthe adsorption and activation of O are hindered in the presence of AMCs, consistent with a\n2\nreduction in proton availability that impedes the formation of the *OOH intermediate. The negative\nshift of the half-wave potential further reflects a slowdown in the overall ORR kinetics, primarily\ndue to the suppression of the 4e⁻ pathway that requires more protons compared to the 2e⁻ pathway,\nwith a lower activation barrier under these conditions, is less affected. This interpretation aligns\nwith the simultaneous decrease in the limiting current density in the LSV curves (Fig. 5b), which\nsignals a shift from the 4e⁻ to the 2e⁻ pathway. The shift in potential in Fig. 2c is also attributed to\nthe suppression of electrode surface reactions caused by AMC.\nCation concentration: We have measured the pH values in varied electrolytes (0.5 M H SO\n2 4\nwith 0.15 M A SO ) and found that the values remained virtually unchanged at ~0.41 (±0.02)\n2 4\n(New Supplementary Fig. 8). In converting the applied potentials to the RHE scale using E(RHE)\n= E(Ag/AgCl) + 0.0592 × pH + 0.197, such minor pH variations introduce a negligible shift (< 2\nmV). Therefore, the observed potential offsets originate principally from cation-induced interfacial\nwater restructuring rather than pH changes.\nThe higher cation concentrations used here (0.3 M) were aimed to emphasize the restructuring effect\non the interfacial H-bond network and its pronounced impact on proton transport under industrially\nrelevant current densities. We have supplemented H O production tests at lower AMC\n2 2\nconcentrations (10 or 50 mM), which confirm that the promotional trend (Cs⁺ > K⁺ > Na⁺ > Li⁺)\npersists even at lower ionic strengths (Fig. R3). This consistency bridges our observations with\nreported cation effects at much lower concentrations (10–50 mM) and underscores the general role\nof AMCs in tuning proton delivery through interfacial water reconfiguration.\nRevisions to Manuscript:\n“To exclude potential interference from other variables, the pH value, ionic conductivity and\ndissolved oxygen in different electrolytes were probed. As shown in Supplementary Figs. 8 and 9,\nthe AMC dosage did not significantly alter the values, indicating such parameters were unlikely to\nbe the primary drivers behind the observed cation-dependent trend.” in Page 6.\nSupplementary Fig. 8 | The pH of electrolytes for 0.5 M H SO and 0.5 M H SO + 0.3 M A+ (A\n2 4 2 4\n= Li, Na, K, or Cs).\nFig. R3. H O yield using Co-CNT at different AMC concentration in a flow-cell reactor.\n2 2\n7. “The addition of AMCs significantly lowered the onset potential and disk current density as\ncompared to the bare electrolyte scenario, which was likely to originate from hindered proton\ntransfer in impeding the HER or 4e⁻ ORR pathways occurring at the catalyst surface4,38” referring\nto Figure 5b.\nIn Figure 5b we can see that the ORR onset is, in fact, delayed after the addition of the cations, i.e.,\nthe reaction onset happens later in the potential sweep.\nReply 7: Thank you for the detailed comment. The onset of the ORR is indeed shifted to more\nnegative potentials upon the addition of AMCs (Figure 5b). This delayed onset indicates that a\nhigher overpotential is required to initiate the reaction in the presence of AMCs.\nWe attribute this shift primarily to the restructuring of the interfacial water network by AMCs,\nwhich hinders proton transport to the catalytic surface. The reduced proton availability raises the\nkinetic barrier for the initial activation of O , thereby delaying the reaction onset. Importantly, this\n2\neffect preferentially suppresses the 4e⁻ ORR pathway, which relies heavily on efficient proton\ndelivery, while the 2e⁻ pathway is less affected under these conditions. This interpretation is\nconsistent with the statement in our manuscript that “the lowered onset potential and disk current\ndensity likely originate from hindered proton transfer, which impedes both the HER and the\ndominant 4e⁻ ORR route”.\nThus, the delayed onset not only reflects the overall kinetic modulation induced by AMCs but also\nprovides further evidence for the precise suppression of the 4e⁻ pathway, in good agreement with\nthe enhanced H O selectivity observed in the present study.\n2 2\n8. “Our catalyst showcased an outstanding turnover frequency (TOF) value of 0.51 s⁻1 at 0.5 VRHE\n(Fig. 5d), which compared favorably with reported electrocatalysts39,40.”\nCalculating TOF requires knowledge of the precise amount of active sites on the surface. How was\nthis determined and how was it calculated?\nReply 8: Thank you for the careful attention regarding the TOF calculation. As noted, an accurate\nassessment of the number of active sites is essential for reliable TOF estimation. In our manuscript,\nthe TOF value of 0.51 s⁻¹ for the Co-CNT catalyst is derived under the assumption that all cobalt\natoms constituting the catalytic cluster architecture function as active sites for H O production. The\n2 2\nTOF is calculated using the following expression:\nTOF = I M / n F m w\nmetal catalyst metal\nHerein, I (A) represents the Faradaic current for H O generation, obtained by multiplying the disk\n2 2\ncurrent by the Faradaic efficiency of H O . M (g/mol) represents the molar mass of cobalt. n =\n2 2 metal\n2 corresponds to the number of electrons transferred per H O molecule formed. F is the Faraday\n2 2\nconstant. m (g) is the total catalyst mass loaded on the electrode, and w is the mass fraction\ncatalyst metal\nof Co in the catalyst, as determined by inductively coupled plasma atomic emission spectroscopy.\nThe number of active sites is accordingly estimated as m ⋅w /M . This approach based\ncatalyst metal metal\non the usage of the total metal content as a proxy for the number of active sites, has been widely\nemployed in the literature for supported metal cluster catalysts, particularly when precise\nquantification of exposed surface atoms is experimentally challenging. Similar methodology has\nbeen adopted in several recent studies with respect to H O electrosynthesis5.\n2 2\nIt is important to emphasize that, since cobalt is present in the form of clusters rather than fully\nexposed single atoms, a portion of the Co sites may not be accessible at the electrolyte–catalyst\ninterface. Therefore, the reported TOF of 0.51 s⁻1 should be regarded as a conservative estimate; the\nactual TOF per exposed active site is likely to be higher. This interpretation is in line with the\nfavorable H O production performance observed in our flow-cell tests. To ensure full transparency\n2 2\nand reproducibility, we have now expanded the Methods section to include a detailed description\nof the active-site estimation protocol and the underlying assumptions in the revised manuscript.\nRevisions to Manuscript:\n“This approach, which presumed all Co atoms in the clusters are active, yielded an overestimated\ncount of apparent active site. This, in turn, resulted in a calculated TOF that systematically\nunderestimated the actual catalytic activity.” in Page 25.\n9. “To uncover the influence of different electrolyte additives upon catalyst activity, we performed\nelectrochemical double-layer capacitance measurements. As anticipated, the catalyst exhibited the\nhighest electrochemical active surface area value in the Cs⁺-dosed electrolyte, indicative of\nfacilitated mass transport at ORR active sites (Supplementary Fig. 37)43,44”\nThis is, likely, a misinterpretation. The authors do in fact measure higher capacitance, however,\nsince the cations are participating in the double layer the specific capacitance is likely different in\nthe presence of different cations, while the ECSA is likely the same.\nReply 9: Thank you for the professional comment regarding the electrochemical metrics. We agree\nwith you that there exists different electrochemical double-layer capacitance with the same\nECSA, owing to the participation of varied cations. Therefore, employing such capacitance value\nto estimate ECSA trends could introduce quantitative discrepancies if the intrinsic capacitance per\nunit area varies with cation identity. The capacitive trend analysis, focusing on relative rather than\nabsolute ECSA comparisons, has been utilized in several recent studies addressing cation-mediated\nelectrolyte effects in electrocatalytic systems6,7. To ensure rigorous and unambiguous discussion,\nwe have revised the relevant section in the revised manuscript to refrain from quantitative\nECSA comparison.\nRevisions to Manuscript:\n“In addition, electrochemical double-layer capacitance measurements demonstrate that the catalyst\nexhibits the highest double-layer capacitance in Cs+-dosed electrolyte, suggesting an optimized\ninterfacial environment that promotes reaction activity and electron transfer.” in Page 11.\n10. The authors measure the generated H2O2 by oxidation at 1.2 V at the ring electrode, but later,\nwhen they measure H2O2 decomposition (also by oxidation) they perform the measurement at 1V.\nWhat is the reason for this? Is the trend the same at 1.2V?\nReply 10: Thanks for your thoughtful question regarding the selection of applied potentials for H O\n2 2\ndetection and decomposition measurements. The difference in the set-up potentials stems from the\ndistinct objectives and underlying electrochemistry of each test.\nRRDE measurements: The Pt ring was held at 1.2 V to quantitatively detect H O generated at\nRHE 2 2\nthe disk electrode. Although the thermodynamic equilibrium potential for H O oxidation to O is\n2 2 2\napproximately 0.695 V , a higher overpotential is required to achieve near 100% detection\nRHE\nefficiency due to kinetic limitations. The choice of 1.2 V is an empirical standard that ensures\nRHE\nrapid and complete oxidation of H O while avoiding interference from the oxygen evolution\n2 2\nreaction (OER) at higher potentials.\nH O reduction reaction (H O RR): H O RR was evaluated based on LSV profiles on a catalyst-\n2 2 2 2 2 2\ncoated glassy carbon electrode in an electrolyte containing 10–50 mM H O . Here, the potential\n2 2\nrange of 0 to 0.8 V is selected to focus on the reduction process (H O + 2H⁺ + 2e⁻ → 2H O)\nRHE 2 2 2\nwhile minimizing contributions from H O oxidation, which becomes significant above ~1.0 V .\n2 2 RHE\nThis approach allows clear observation of the cathodic current associated with H O consumption.\n2 2\nRegarding the trend at 1.2 V , we note that the oxidation current measured at the ring electrode\nRHE\n(at 1.2 V) directly reflects the amount of H O produced, and thus the trend in H O generation\n2 2 2 2\nacross electrolytes remains consistent. However, if H O reduction is to be measured at 1.2 V, the\n2 2\nsignal would be convoluted by simultaneous oxidation, leading to misinterpretation. Therefore, the\nselected potential windows have been optimized for their respective purposes, with the\nreported trends remaining robust within the applied conditions.\n11. “The on-site generation of H2O2 in our system is highly attractive for targeted applications\nsuch as wastewater treatment.”\nHow would such high alkaline cation concentrations (0.3M) be applied in practical wastewater\ntreatment. Seems like this would be arguably a worse or equally bad contamination as the\ncontaminants being removed by H2O2 generation.\nReply 11: We are grateful for you to raise this critical comment regarding the practical applicability\nof our electrolyte engineering strategy. The concern about potential secondary contamination from\nadded salts is well-noted and highly relevant for real implementation.\nOur systematic investigation demonstrates that AMCs exert a concentration-dependent promotional\neffect on H O electrosynthesis across a wide cation concentration range (0.1–0.5 M), accompanied\n2 2\nby a systematic investigation at 0.3 M. Importantly, this concentration range is not only effective\nbut also directly compatible with existing high-salinity wastewater, particularly reverse osmosis\n(RO) concentrates, which typically contain total dissolved solids well above 0.3 M. Following this\ncomment, we have supplemented further discussion in the revised manuscript.\nRevisions to Manuscript:\n“Our foregoing evaluation confirmed that dosed AMC concentrations ranging from 0.1–0.5 M\nwould enhance H O electrosynthesis, with core investigation at 0.3 M that could be encountered in\n2 2\nindustrial effluents such as reverse osmosis concentrates. Under such conditions, endogenous ions\nserve as inherent promoters, eliminating the external additives.” in Page 13.\n12. Figure 2A the voltammograms are too small to be assessed. Why are the potential windows for\nthe blank and the other CVS different?\nCaption for Figure 5f is missing.\nReply 12: Thanks for the valuable suggestions. Following this comment, we have rephrased the\nvoltammogram profile to improve readability and ensure the clear visibility of the key features of\ndiffusion-limited current plateaus (New Fig. 2a). The apparent difference in potential windows\nbetween the blank and cation-modified systems stems from the marked suppression of proton\ndiffusion upon AMC addition, which shifts the HER onset to more negative potentials. To facilitate\ndirect comparison, we have now aligned all polarization curves within a common potential window\nin the updated figure. Furthermore, we have included additional comparative LSV curves for each\ncation system in the revised Supplementary Information file to provide view of the cation-dependent\nsuppression behavior (New Supplementary Fig. 3). In addition, we have supplemented the figure\ncaption for Fig. 5f.\nRevisions to Manuscript:\nSupplementary Fig. 3 | HER polarization curves for Au microelectrodes in Ar-saturated 0.5 M\nH SO + 0.15 M A SO at 0.02 V s⁻1 (A = Li, Na, K, or Cs).\n2 4 2 4\nReviewer #3 (Remarks to the Author):\nAuthors have presented a comprehensive report on the effects of alkali metal ions on H2O2\nelectrosynthesis and found an interesting size-dependent trend. They have also furnished\nmechanistic insights into the transformation. While the quantify of experimental work is quite\nsignificant, I have some concerns about the impact of the main claims, which renders the current\nversion not par with the Nature Comms level. However, there may be some ways to bridge the gap.\nBelow, I present my major points and invite the co-authors to address them in their revised version,\nwhose suitability I will be happy to advise the Editor on.\nResponse: Thank you for your approval to our work. All requested revisions based on your\ninstructive comments have been carried out.\nMain points:\n1. The authors claim that they are the first to investigate the efficacy trend of alkali metal cations\n(AMC) on the electrosynthesis of H2O2, and they suggest that AMC effectiveness follows the trend:\nLi+ < Na+ < K+ < Cs+. This claim seems misleading, as other researchers have investigated the\nrole of AMC on the electrosynthesis of H2O2 in acidic media already, and, interestingly, found\ndissimilar trends.\na. For Instance, Peike Cao et. al. (Reference 14) have also investigated the effect of AMCs and\nreported that “the effective H2O2 currents followed the trend of K+ > Na+ > Li+ ≈ Rb+ ≈\nCs+ > NH4 +, which was related to their hydrated ionic radius,” which seems to contradict the\nmain finding of this manuscript.\nb. Also, Zhang et. al. (Reference 13) has also examined the effect of the same (Li+, Na+, K+, Cs+)\ncations on H2O2 production and suggested that the “production rates of H2O2 are relatively\nunaffected by the size of the alkali metal cations in the electrolyte.”\nThese differing views and the underlying factors should be discussed in the manuscript to truly\nadvance the field forward.\nReply 1: Much appreciated for raising this crucial point on the consistency of reported trends for\ncation-effect. We admit that it is not objective to claim “the first” to report the trend without\nanalyzing the prior studies in this field. In the revision, we have conducted a detailed analysis by\nfully acknowledging the important work by Cao et al. (Ref. [35]) and Zhang et al. (Ref. [33]),\nshowing a distinct cation-effect trend.\nIn further contexts, the core innovation of our work is not merely reporting a cation-effect trend but\nin providing a unique mechanistic explanation rooted in the cation-specific restructuring of the\ninterfacial hydrogen-bond (H-bond) water network. Previous studies primarily attributed cation-\neffects to electrostatic shielding (Zhang et al.) or intermediate stabilization influenced by hydrated\nionic radius (Cao et al.). Our combined experimental and theoretical evidence reveals that larger\ncations (Cs⁺) more effectively disrupt the interconnected H-bond network at the electrode-\nelectrolyte interface, thereby imposing a greater kinetic barrier to proton transport. The monotonic\ntrend we observe emerges clearly when this interface-specific water restructuring becomes the\ndominant factor. According to your kind suggestion, we have supplemented further discussion in\nthe revised Introduction section.\nRevisions to Manuscript:\n“Particularly using metal-free carbon-based electrocatalysts, Zhang et al. observed that H O\n2 2\nproduction rates can be improved by an AMC-induced shielding effect, yet were relatively\nunaffected by the AMC types33. A recent study indicated a graphitic carbon electrode exhibited an\nincrease in 2e⁻ ORR selectivity following the trend of Cs+ > K+ > Na+ > Li+, where different cations\nwould interfere in the *OOH binding and influence the local electric field34. It is interesting to note\nthat the AMC roles in acidic 2e⁻ ORR manifest inconsistent trends. In further contexts, the cation\neffect is effective across other catalyst types apart from carbonaceous materials. For instance, Cao\net al. found the H O production followed the trend of K⁺ > Na⁺ > Li⁺ ~ Cs⁺ with the employment\n2 2\nof earth-abundant TiC catalyst, which was dictated by the hydrated ionic radius of AMCs35. These\ndisparities imply that the dominant mechanism might be sensitive to specific conditions, such as\ncatalyst identity, cation concentration and operating current density. A coherent picture of how\nAMCs reshape the interfacial environment in acidic 2e⁻ ORR is still lacking, especially in the realm\nof prevailing transition metal-involved carbon electrocatalysts (e.g., metal-carbon nanotube, M-\nCNT).” in Page 3-4.\n2. I wish to draw the co-authors’ attention to the recent reports on the spontaneous H2O2 production\nat water’s interfaces with solids (Refs.1-3), contradicting many groups’ experimental and\ncomputational claim that the H2O2 forms at the air-water interface and has much hitherto\nunrealized environmental and applied significance (4-12). This specific topic has in fact become\nvery active – a number of reports corroborating the experimental work from the Mishra group (13-\n17)(also, see https://arxiv.org/abs/2506.23988 & https://arxiv.org/abs/2508.15422. Therefore, these\nimportant findings could be highlighted to underscore the importance of this problem (and perhaps\nan opinion on the right answer!).\nReply 2: Thanks for guiding us to the fascinating and rapidly evolving body of work on\nspontaneous H O formation at solid-water interfaces (e.g., Refs. 1-3 & preprints\n2 2\narxiv:2506.23988, arxiv:2508.15422). This is a paradigm-shifting discovery with profound\nimplications for environmental chemistry and catalysis.\nWhile our study focuses on electrocatalytic H O generation via the 2e⁻ ORR pathway, we fully\n2 2\nrecognize the significance of these recent reports demonstrating non-electrochemical, interface-\ninduced spontaneous generation. This phenomenon underscores the inherent redox activity of\nwater-derived intermediates at confined interfaces, a concept that profoundly enriches the context\nof our work.\nIn the revised Introduction, we have incorporated a discussion of these findings. We posit that the\nspontaneous generation and the electrocatalytic generation of H O might share a common\n2 2\nunderlying thread: the critical role of interfacial water structure and dynamics in modulating\noxygen reduction pathways. In spontaneous processes, unique interfacial water orientations and\nreactivities may facilitate O activation. In our electrocatalytic system, we demonstrate that alkali\n2\nmetal cations can be used as a “tool” to precisely engineer this interfacial water network, thereby\nsteering the selectivity towards H O .\n2 2\nThus, rather than contradicting, these recent discoveries and our work mutually reinforce the\noverarching importance of the interfacial aqueous microenvironment in governing oxygen\nreduction chemistry. Our study provides a tunable, electrochemical strategy to manipulate this\nmicroenvironment for technological gain, directly bridging fundamental interfacial science with\napplied environmental engineering. We thank you for this insightful connection, which has allowed\nus to position our work within a broader and more impactful scientific narrative. We have\nincorporated the advised references in the revised manuscript (New Ref. [15–26]).\nRevisions to Manuscript:\n“Beyond deliberate electrocatalysis, the fundamental chemistry of aqueous interfaces has garnered\nintense interest for its role in spontaneous chemical transformations15-18. Notably, the purported\nspontaneous formation of H O has sparked significant debate. While early reports emphasized\n2 2\nremarkable accelerations at the air-water interface19-23, a growing body of evidence, including\nrigorous work from the Mishra group, suggests that such phenomena might be more pronounced at\nsolid–water interfaces24-27. This ongoing controversy underscores that the intrinsic drivers of\nreactivity in confined aqueous environments remain a fundamental question in interfacial science.\nResolving this question is critical not only for understanding natural processes but also for the\nrational manipulation of electrolyte–catalyst interfaces.” in Page 3.\n3. In this work, the authors have exploited the potassiumtitanyloxalate (PTO) reagent for the\ndetection and quantification of H2O2. Unfortunately, this reagent unspecific to H2O2, i.e., it can\ngive a false positive signal with reactive oxygen species (ROS) such as hydroxyl radicals,\nsuperoxides, and hydroperoxyl radicals, etc. Therefore, the quantification and formation of H2O2\nshould be confirmed by a more reliable and direct technique, such as 1 H-NMR spectroscopy (See\nRefs.18). It is possible that the reported signals originate from different ROS, and the actual H2O2\nconcentrations are much lower than the reported ones.\nReply 3: Thank you for raising this important methodological point regarding the specificity of the\npotassium titanyl oxalate (PTO) assay for H O quantification. We acknowledge that, in principle,\n2 2\nPTO can react with other reactive oxygen species (ROS), which could potentially lead to an\noverestimation of H O concentration in complex matrices. In the context of our work, several\n2 2\nfactors assure the reliability of the reported H O yields:\n2 2\nSystem specificity: Our electrochemical system is designed for the selective 2e⁻ ORR to generate\nH O . The reaction environment (acidic electrolyte, controlled potential) and the use of a selective\n2 2\nCo-CNT catalyst strongly favor the direct production of H O over other ROS. The substantial and\n2 2\nstable generation of H O is further corroborated by its successful on-site application in electro-\n2 2\nFenton reactions for pollutant degradation, which specifically requires H O as the precursor\n2 2\nfor ·OH generation.\n1H-NMR spectroscopy validation: As advised, we have constructed the corresponding standard\ncurve and conducted 1H-NMR quantification on key samples to provide a validation for the\nconsistency between the potassiumtitanyloxalate (PTO) reagent detection and 1H-NMR\nspectroscopy method (New Supplementary Fig. 49). These results provide strong evidence\nsupporting the accuracy of our reported H O production data.\n2 2\nCross-validation with Ce(IV) titration: To address the concern of assay specificity, we performed\nparallel quantification using the established cerium(IV) sulfate titration method. As shown in the\nsupplementary figure below (Fig. R4), the H O concentrations determined by both methods are in\n2 2\ngood agreement across the measured range. This cross-validation confirms that the signal measured\nby PTO in our system originates predominantly from H O .\n2 2\nWidespread use in electrocatalysis tests: The PTO method is a standard, widely adopted technique\nfor H O quantification in electrocatalytic studies, including recent high-impact publications, due to\n2 2\nits sensitivity, simplicity, and suitability for rapid, high-throughput analysis—a practical necessity\nfor evaluating performance under various reaction conditions. In addition, The PTO method agrees\nwell with 1H-NMR quantification and cerium(IV) sulfate titration method in our system,\ndemonstrating its practical feasibility.\nSupplementary Fig. 49 | Quantification of H O by 1H-NMR spectroscopy. Standard samples of\n2 2\nknown H O (aq) concentrations were prepared by diluting a concentrated 30% (v/v) stock solution\n2 2\nusing HPLC-grade deionized water. All 1H-NMR tests were carried out at 2℃.\nFig. R4. The linear absorbance-concentration calibration curve of different concentrations of\nstandard H O solutions titrated with cerium(IV) sulfate.\n2 2\n4. Lines 76-78: “We observe that the 2e⁻ ORR selectivity in acid follows a monotonically increasing\ntrend from Li+ to Cs+, which is not revealed by prior studies (Figure 1).” Figure 1 is an illustration;\nit does not represent any experimental observation.\nReply 4: Much appreciated for this detailed suggestion. Figure 1 is a conceptual scheme designed\nto summarize the proposed mechanism of cation-mediated interfacial water restructuring, which is\nnot suitable to represent the experimental trend as you pointed out. Following this comment, we\nhave repositioned Figure 1 in the revised manuscript to ensure a clear distinction between the\npresentation of experimental data (in subsequent figures) and explanatory scheme.\nRevisions to Manuscript:\n“The varying influences of different AMCs stem from their distinct capabilities to perturb the H-\nbond network, associated with their hydration properties and coordination strengths. Such a\nperspective is essentially different from the general recognition of AMC roles in this field (Fig. 1).”\nin Page 4.\n5. Supplementary Fig. 3: Concentration-dependent variation of suppression effect—How does the\nconcentration of electrolyte affect the solubility of oxygen in the electrolyte? It is essential to\nquantify the dissolved oxygen concentrations in the electrolyte as a function of the electrolyte\nconcentration. Some other important parameters that need to be kept comparable are ionic\nconductivity and dissolved oxygen levels.\nReply 5: Thank you for raising this critical point. We fully agree that these are essential variables\nto control and characterize when comparing electrochemical performance across different\nelectrolytes, as they exert direct impact upon mass transport and reaction kinetics.\nFollowing this comment, we have supplemented systematic measurements to quantify these key\nparameters in all electrolyte systems used in this study (0.5 M H SO + 0.15 M Li SO , Na SO ,\n2 4 2 4 2 4\nK SO or Cs SO ). The relating ionic conductivity and dissolved oxygen concentrations results have\n2 4 2 4\nbeen summarized in the New Supplementary Fig. 9.\nIonic conductivity: Conductivity was measured using a calibrated conductivity meter at 25 °C. The\naddition of 0.3 M AMC caused only a marginal variation in overall conductivity (167–173 mS cm⁻1)\ncompared to the baseline 0.5 M H SO (~168 mS cm⁻1). This minor variation is unlikely to be the\n2 4\nprimary driver of the cation-specific trends observed in ORR selectivity and activity.\nDissolved oxygen concentration: Dissolved oxygen levels were quantified using a Clark-type\noxygen electrode in O -saturated electrolytes at 25°C. The data confirm that for a given total\n2\nmolarity of added sulfate (e.g., 0.15 M A SO ), the dissolved oxygen concentrations were nearly\n2 4\nidentical (± 3%) across the different AMCs (Li+, Na+, K+, or Cs+). This indicates that the relative\ncomparison of cation effects presented in our work (e.g., the trend Li+ < Na+ < K+ < Cs+) was\nperformed under effectively equivalent O availability.\nand H O production rate are attributed to the specific cationic modulation of the interfacial water\n2 2\nstructure and proton transport.\nRevisions to Manuscript:\n“To exclude potential interference from other variables, the pH value, ionic conductivity and\ndissolved oxygen in different electrolytes were probed. As shown in Supplementary Figs. 8 and 9,\nthe AMC dosage did not significantly alter the values, indicating such parameters were unlikely to\nbe the primary drivers behind the observed cation-dependent trend. These results motivated us to\ninvestigate the restructuring of interfacial H-bonding water networks in the presence of AMCs.” in\nPage 6.\nSupplementary Fig. 9 | (a) Electrolyte ionic conductivity (EC) and (b) electrolyte dissolved oxygen\n(DO) for 0.5 M H SO and 0.5 M H SO + 0.3 M A+ (A = Li, Na, K, or Cs).\n2 4 2 4\n6. Use the right symbol for free radicals: replace *OOH, *OH, *O by •OOH, •OH, •O\n(Asterisk by Dot).\nReply 6: We sincerely thank you for the meticulous attention to notation. Indeed, the asterisk *\nexplicitly represents an active site on the catalyst, which is distinct from the notation for free radicals\nin solution (e.g., •OH, •OOH) employing a dot (•) to indicate an unpaired electron. To avoid any\nambiguity, we have carefully reviewed the entire manuscript to ensure this notation is used correctly\nand consistently.\n“*OOH” in Page 9-14, “*OH” in Page 9, “*O” in Page 9,14, “•OH” in Page 13.\n___________________\nEmployed Reference:\nN–C Catalysts? J. Am. Chem. Soc. 145, 25352-25356 (2023).\nOxygen Reduction. J. Am. Chem. Soc. 143, 9423-9428(2021).\nElectrocatalytic Hydrogen Peroxide Production. Adv. Mater. 37, 2502197(2025).\nToward Excellent Electrochemical H O Production and Biomass Upgrading Angew. Chem.\n2 2\nInt. Ed. 64 e202502383(2025).\nover Dual-Atom Catalysts for H O Electrosynthesis. J. Am. Chem. Soc. 147, 11465-\n2 2\n11476(2025).\n2\nin strongly acidic medium. Nat. Catal. 5, 268–276(2022).\nKinetics. JACS Au 1, 1674-1687(2021)."}, {"round": 2, "reviewer_comments": {"reviewer_1": "I recommend acceptance of the revised manuscript. The authors have adequately addressed my previous comments, and the revisions have improved the clarity, transparency, and reproducibility of the work."}, "author_response": "Response to Reviewers’ Comments (NCOMMS-25-75578A)\nReviewer #1 (Remarks to the Author):\nThe revision has addressed most of my previous concerns and the manuscript is now close to\nacceptance. I believe the work can be accepted after the authors clarify the following minor points\nto improve transparency and reproducibility.\nResponse: Thank you for your approval to our work. All requested revisions based on your detailed\ncomments have been carried out.\nSpecific comment:\n1. In the rebuttal, the authors state that adsorption energies of alkali metal cations (AMCs) on the\ncatalyst surface were calculated to exclude direct surface binding. It would be helpful to explicitly\nclarify whether the positive charge state of the cations (e.g., Li⁺/Na⁺/K⁺) was considered in these\ncalculations, and how charge compensation was handled in the computational model (e.g.,\nbackground charge, counter ions, solvation treatment, etc.), since this may influence adsorption\nenergetics and interfacial electrostatics.\nReply 1: Thank you for your inspiring suggestions regarding the theoretical calculation of charged\nspecies. In the adsorption energy calculations for AMCs (Li⁺, Na⁺, K⁺, Cs⁺) on the catalyst surface,\nwe explicitly accounted for their charge state. This was done by manually adjusting the number of\nvalence electrons in the computational model to reflect the cationic charge, consistent with standard\npractice for simulating charged adsorbates in periodic calculations.\nTo maintain the interfacial electrostatic environment in actual experiments, we adopted the\nfollowing strategy in the simulation:\nCharge compensation: A uniform background charge (jellium model) was applied to neutralize the\nsystem, which is a common and efficient approach for periodic DFT calculations involving isolated\ncharged species.\nSolvation treatment: The implicit solvation model (VASPsol++) was employed throughout, which\nself-consistently accounts for the dielectric response of the electrolyte and helps stabilize the\ncharged interface.\nModel validation: The adsorption energies obtained under this setup consistently showed positive\nvalues for all AMCs, confirming that the cations remain preferentially solvated rather than adsorbing\ndirectly onto the catalyst surface, which is consistent with our experimental interpretation.\nThese methodological details, including the handling of cation charge states and the charge\ncompensation scheme, have now been incorporated into the Methods section of the revised\nmanuscript (please see “DFT Calculations” subsection) to ensure full transparency and\nreproducibility.\nRevisions to Manuscript (in Page 27):\n“We applied a constant potential (+ 0.15 V vs SHE, consistent with experiments) to the ORR\nsimulations by CP-VASP.90,91 The targeted Fermi level referenced to the electrolyte was set to -4.75\nV with a convergence criterion of 0.01 V for the electrolyte-referenced Fermi level. The solvation\nmodel during in calculations was considered by using the VASPsol++ framework implemented in\nVASP.92\nEach AMC was put on the surface of Co-CNT and the adsorption energies were calculated by DFT:\nE = E – E – E\nad total Co-CNT AMC\nwhere E is the total energy of an AMC adsorbed on the catalyst, while E and E represent\ntotal Co-CNT AMC\nthe energies of Co-CNT and AMC, separately. Positive-charged calculations were performed by\nmanually setting the number of valence electrons to consider the positive charge state of AMCs.”\n2. The authors performed constant-potential DFT calculations at 0.15 V vs. SHE to account for\nelectrochemical conditions. While this value is described as being consistent with experimental\nconditions, it would strengthen the manuscript to provide a clearer electrochemical/theoretical\nrationale for selecting this potential, or to briefly discuss how representative it is within the relevant\nORR operating window and whether the conclusions are expected to remain robust over a reasonable\npotential range.\nReply 2: We thank the reviewer for this thoughtful question regarding the choice of the applied\npotential in our constant-potential DFT calculations. The potential of +0.15 V vs. SHE was selected\nbecause it corresponds directly to the working electrode potential measured during our flow-cell\nexperiments at a current density of 50 mA cm–2, a commonly employed current density in\nlaboratory-scale studies of H O electrosynthesis (Ref: J. Am. Chem. Soc. 2024, 146, 9434; Adv.\n2 2\nFunct. Mater. 2025, 35, 2417090).\nA brief discussion of the potential selection and the associated validation have now been added to\nthe revised manuscript (Methods section).\nRevisions to Manuscript (in Page 27):\n“Each AMC was put on the surface of Co-CNT and the adsorption energies were calculated by DFT:\nE = E – E – E\nad total Co-CNT AMC\nwhere E is the total energy of an AMC adsorbed on the catalyst, while E and E represent\ntotal Co-CNT AMC\nthe energies of Co-CNT and AMC, separately. Positive-charged calculations were performed by\nmanually setting the number of valence electrons to consider the positive charge state of AMCs.”\n3. In the Methods section, particularly in the “DFT calculations” subsection near the end of the\nmanuscript, there are currently no references cited for the computational methods and models\nemployed. Please add appropriate methodological references (e.g., DFT setup,\nelectrochemical/solvation model, constant-potential method) to improve completeness and\nreproducibility.\nReply 3: We apologize for the oversight in not including the methodological references in the\noriginal submission. We have now carefully added all relevant references in the “DFT Calculations”\nsubsection of the Methods section. These include citations for:\n(1) The DFT software package and computational setup (VASP, PAW pseudopotentials, plane-wave\ncutoff, exchange-correlation functional, and dispersion correction).\n(2) The implicit solvation model (VASPsol++) used to describe the electrolyte environment.\n(3) The constant-potential (CP-VASP) methodology employed to incorporate the electrochemical\npotential in our ORR simulations.\n(4) Key methodological papers supporting the thermodynamic analysis (e.g., free-energy\ncorrections, treatment of adsorbates, and determination of reaction pathways).\nBy providing these references, we aim to offer a fully transparent and reproducible account of our\ncomputational approach. The related references have been updated as New Ref. [70-93].\nRevisions to Manuscript (Added References):\n70. Berendsen, H. J. C. et al. A message-passing parallel molecular-dynamics implementation.\nComput. Phys. Commun. 91, 43–56 (1995).\n71. Lindahl, E. et al. GROMACS 3.0: A package for molecular simulation and trajectory\nanalysis. J. Mol. Model. 7, 306–317 (2001).\n72. Hess, B. et al. GROMACS 4: Algorithms for highly efficient, load-balanced, and scalable\nmolecular simulation. J. Chem. Theory Comput. 4, 435–447 (2008).\n73. Berendsen, H. J. C. et al. Molecular dynamics with coupling to an external bath. J. Chem.\nPhys. 81, 3684–3690 (1984).\n74. Darden, T. et al. Particle mesh Ewald: An N log(N) method for Ewald sums in large systems.\nJ. Chem. Phys. 98, 10089–10092 (1993).\n75. Markland, T. E. et al. Nuclear quantum effects enter the mainstream. Nat. Rev. Chem. 2,\n0109 (2018).\n76. Rebstock, J. A. et al. Comparing interfacial cation hydration at catalytic active sites and\nspectator sites on gold electrodes. Chem. Sci. 13, 7634–7643 (2022).\n77. Ringe, S. et al. Understanding cation effects in electrochemical CO reduction. Energy\n78. Rossi, M. et al. Nuclear quantum effects in H⁺ and OH⁻ diffusion along confined water\nwires. J. Phys. Chem. Lett. 7, 3001–3007 (2016).\n79. Varma, S. et al. Coordination numbers of alkali metal ions in aqueous solutions. Biophys.\nChem. 124, 192–199 (2006).\n80. Kühne, T. D. et al. CP2K: An electronic structure and molecular dynamics software\npackage—Quickstep. J. Chem. Phys. 152, 194103 (2020).\n81. Ozkanlar, A. et al. ChemNetworks: A complex network analysis tool for chemical systems.\nJ. Comput. Chem. 35, 495–505 (2014).\n82. Kresse, G. et al. Ab initio molecular-dynamics simulation of the liquid–metal amorphous–\nsemiconductor transition in germanium. Phys. Rev. B 49, 14251–14269 (1994).\n83. Kresse, G. Ab initio molecular dynamics for liquid metals. J. Non-Cryst. Solids 193, 222–\n229 (1995).\n84. Kresse, G. et al. Efficiency of ab initio total energy calculations for metals and\nsemiconductors using a plane-wave basis set. Comput. Mater. Sci. 6, 15–50 (1996).\n85. Kresse, G. et al. Efficient iterative schemes for ab initio total-energy calculations using a\nplane-wave basis set. Phys. Rev. B 54, 11169–11186 (1996).\n86. Kresse, G. et al. From ultrasoft pseudopotentials to the projector augmented-wave method.\nPhys. Rev. B 59, 1758–1775 (1999).\n87. Perdew, J. P. et al. Atoms, molecules, solids, and surfaces: Applications of the generalized\ngradient approximation for exchange and correlation. Phys. Rev. B 48, 4978–4978 (1993).\n88. Hammer, B. et al. Improved adsorption energetics within density-functional theory using\nrevised Perdew–Burke–Ernzerhof functionals. Phys. Rev. B 59, 7413–7421 (1999).\n89. Grimme, S. et al. Effect of the damping function in dispersion-corrected density functional\ntheory. J. Comput. Chem. 32, 1456–1465 (2011).\n90. Zhao, X. H. et al. Origin of selective production of hydrogen peroxide by electrochemical\noxygen reduction. J. Am. Chem. Soc. 143, 9423–9428 (2021).\n91. Yu, S. R. et al. What is the rate-limiting step of oxygen reduction reaction on Fe–N–C\ncatalysts? J. Am. Chem. Soc. 145, 25352–25356 (2023).\n92. Islam, S. M. R. et al. An implicit electrolyte model for plane wave density functional theory\nexhibiting nonlinear response and a nonlocal cavity definition. J. Chem. Phys. 159, 204103\n(2023).\n93. Bochevarov, A. D. et al. Jaguar: A high-performance quantum chemistry software program\nwith strengths in life and materials sciences. Int. J. Quantum Chem. 113, 2110–2142 (2013).\nReviewer #2 (Remarks to the Author):\nDear Authors, Dear Editors,\nI feel like all my remarks have been answered appropriately.\nBest regards.\nResponse: Thank you for your approval of our work.\nReviewer #3 (Remarks to the Author):\nI am satisfied by the additional work the co-authors have done to elevate their contribute to Nature\nComms. level. I recommend it's acceptance.\nResponse: Thank you for your approval of our work."}]

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