Introduction

Hydrogen peroxide is a green energy substance that has garnered attention in various fields, including chemical manufacturing1, biological and medical engineering2,3, energy technology, and environmental remediation3,4,5. The traditional method of synthesizing H2O2 poses safety risks and requires large-scale infrastructure, resulting in energy waste and the need for H2O2 purification4,6,7. In-situ production of hydrogen peroxide, particularly through electrocatalytic methods, is a promising alternative due to its safe operation, ambient reaction conditions, on-demand production, and availability of inexpensive raw materials8,9.

Electrocatalytic synthesis of hydrogen peroxide can be achieved by two kinds of pathways: two-electron cathode oxygen reduction reaction (2e− ORR) and two-electron anodic water oxidation reaction (2e− WOR). 2e− ORR pathway has been widely studied10,11,12,13,14. However, this method requires oxygen as a raw material, and a pumping device is unavoidably required during the entire reaction process because of the low solubility (~8 mg L−1 at 25 °C, 1 atm) and slow diffusion (2.1×10−5 cm−2 s−1) of oxygen in the water, leading to high energy consumption and low oxygen utilization efficiency (<1%)12. By contrast, 2e− WOR does not require oxygen gas and is not limited by low solubility and slow diffusion of oxygen gas, making this method suitable in an oxygen-deficient environment. Moreover, the green energy of hydrogen (H2) can be obtained simultaneously in the water electro-decomposition process (Eq. (1)).

$${{{{\rm{H}}}}}_{2}{{{\rm{O}}}}\to {{{{\rm{O}}}}}_{2}+{{{{\rm{H}}}}}_{2}\uparrow$$
(1)

At present, the most frequently used electrolyte in H2O2 electro-synthesis through 2e- WOR are KHCO3 solution and K2CO315,16,17,18,19,20. Gill et al. reported the role of bicarbonate-based electrolytes in the production of hydrogen peroxide through water oxidation. Specifically, HCO3− is oxidized to HCO4− or C2O62− intermediates at the anode, and these intermediates oxidize water to hydrogen peroxide and then change back to HCO3− simultaneously21,22. Nadar et al. found that BiVO4 with different index planes has different water oxidation performance to produce hydrogen peroxide. BiVO4 with high-index planes is more favorable for the adsorption of HCO3− and thus favors the following synthesis of hydrogen peroxide23. Hence, bicarbonate is beneficial to H2O2 generation.

However, bicarbonate-based electrolytes also suffer from two disadvantages. Firstly, it leads to the decomposition of hydrogen peroxide molecules leading to low hydrogen peroxide generation efficiency23. Jawad et al. demonstrated that bicarbonate consumes hydrogen peroxide. Specifically, bicarbonate can react with ·OH to generate water and ·CO3− radicals, and react with hydrogen peroxide to generate water and HCO4− 24. Sablas et al. believe that both carbonate and bicarbonate can react with hydroxyl radicals to form ·CO3−. ·CO3− consumes hydrogen peroxide to form bicarbonate and HO2·, and HO2· can react with ·CO3− to form carbonate25. That is, the presence of carbonate or bicarbonate decomposes the generated hydrogen peroxide. We have also proved that hydrogen peroxide cannot exist stably at ambient in the presence of bicarbonate even without applied bias. Specifically, the concentration of 1 mM hydrogen peroxide in 2 M KHCO3 solution will decrease by 90% after 9 h. Second, bicarbonate is more soluble in the pH range of 7.0 − 9.0. However, if the pH is greater than 9, it will become carbonate with lower solubility24.

Besides, another problem is that some applications of hydrogen peroxide in the environmental field require an acidic to neutral initial pH, but carbonates are often alkaline. Cao et al. show that effective degradation of tetracycline via H2O2 needs in a pH range of 3–726. Fu et al. demonstrated that the pH value of tetracycline degradation by active H2O2 is 627. Wang et al. degrade norfloxacin by H2O2 at a wide initial pH. Besides, the degradation efficiency of acidic to neutral is better than that of strong alkaline28. Furthermore, many applications of iron-related hydrogen peroxide degraded pollutants are biased towards neutral pH29,30. Nie et al. demonstrated that acid red B (ARB) degradation efficiency by H2O2 is best at neutral pH and the degradation rate increased with the decrease of initial pH value29. Jiang et al. proved that phenol degradation by H2O2 in initial neutral pH can achieve the best efficiency30. Nie et al. degrade dimethyl phthalate by H2O2 presents best at neutral pH28. Therefore, an anode material operating in a non-carbonate environment is very important.

However, there are only a few reports about 2e− WOR H2O2 generation in non-bicarbonate systems. Among them, Zinc (II) porphyrin, and Tin (IV)-porphyrins are water-soluble, and they need to adjust the sodium sulfate electrolyte to an alkaline environment to produce hydrogen peroxide31,32. Moreover, the H2O2 yield can only reach the nmo h−1 level. Although both C, N-doped TiO2 and commercial fluorine-doped SnO2 FTO glass can produce hydrogen peroxide via 2e− WOR in acidic sodium sulfate electrolyte, the hydrogen peroxide production rate is only 0.2900 μM h−1 and 2.2794 μmol h−1, respectively33,34. There are also cases where Na2HPO4 and NaH2PO4 are used as electrolytes in the field of water oxidation. These solutions are alkaline, while both the hydrogen peroxide yields are very low, 2.0589 and 1.3971 μmol h-1, respectively34. It can be seen that the H2O2 yield is very low when there is no carbonate mediated for water oxidation to H2O2.

Li et al. present that H2O2 generation through both direct (via OH−) and indirect (via HCO3−) pathways. *HCO3 will be converted to *HCO4 and the O-O bond form. It proves that percarbonate is the critical species generated through the conversion of bicarbonate at Ga-Ga dual sites because percarbonate will decompose to generate H2O2 and carbonate. Therefore, percarbonate is the crucial intermediate for the generation of H2O2. The peroxyl bond is stable on the surface of the ZnGa2O4 anode. This work demonstrates that the key point for the H2O2 generation is the formation of a crucial intermediate that stabilizes the O-O bond on the surface of the catalyst and the generation of H2O2 is affected by the stability of the crucial intermediate and the oxygen bond on the anode surface22. The presence of oxygen acid like carbonate is important in the process of water oxidation to produce hydrogen peroxide.

Actually, in the presence of carbonate, O-O is formed by the participation of carbonate in the solution and OH part in the water. Therefore, this inspires us that if OH and oxygen-acid can be provided, we can prepare H2O2 by a two-electron water oxidation strategy even without carbonate. If we want to increase the yield of hydrogen peroxide in the system without carbonate, the most important step is the formation of sufficient O-O bonds.

Sulfate is a very comment oxygen acid and it will be applicable if we can use sulfate to replace carbonate. However, compared to sulfate is difficult to react with OH from electrolytes to form O-O bond. In our opinion, to achieve this, a suitable catalyst is required to work, which can provide OH to contribute to the reaction.

Herein, we achieve this goal by using basic copper carbonate (Cu2(OH)2CO3), which is commonly found on the surface of copper. We discover that these materials contain lattice OH in its structure, and we achieve the production of H2O2 via the 2e− WOR method in non-carbonate Na2SO4 electrolyte on a tablet basic copper carbonate independent anode. A crystal OH mediating pathway for the generation of hydrogen peroxide in non-carbonate Na2SO4 electrolyte. SO4 in the electrolyte combined with the crystal OH part in Cu2(OH)2CO3 to form SO4OH*. SO4OH* is further transferred to OOH and SO3. The O-O the band is formed. This O-O bond length is calculated to be 1.47 Å by DFT, which is close to the O-O bond length of 1.48 Å in H2O2. The combination of OOH and H from H2O would form H2O2. Meanwhile, the remaining OH in water combines with Cu2(OH)CO3 and returns to Cu2(OH)2CO3. SO3 combined with another H2O to form SO3·H2O. SO3 part in SO3·H2O combines with OH in water to form HSO4* and release a H+. And then the HSO4* ionizes to produce a SO42− and another H+. This pathway provides an idea in the field of hydrogen peroxide production. It provides a method for hydrogen peroxide production via two-electron water oxidation in a non-carbonate system.

Results

H2O2 electro-synthesis via 2e− water oxidation in Na2SO4 electrolyte by Cu2(OH)2CO3 tablet

The production of H2O2 in a non-carbonate electrolyte is a challenging task. For instance, when graphite is used as the anode in a non-carbonate electrolyte of Na2SO4, it cannot oxidize water to generate H2O2. In contrast, loading basic copper carbonate powder onto carbon paper (Fig. 1a) as the anode can produce 26.32 μM of hydrogen peroxide in 120 min under 3.4 V vs. RHE, which is much better than the widely-used graphite anode. However, the loaded powder of basic copper carbonate quickly falls off from the carbon paper (see Text S12 for more details), the comparison as illustrated in Fig. 1a, leading to a stagnant H2O2 concentration in the later stage. To address this issue, the catalyst is directly pressed into a solid tablet as the anode, as shown in Fig. 1b. In our experiment, it is observed that Cu2(OH)2CO3 powder can be pressed into a tablet block (Fig. 1b) with the addition of a 60% PTFE binder. It’s due to that the Cu2(OH)2CO3 powder could not be pressed into a tablet block without PTFE preprocessing regardless of the applied pressure (details shown in Text S4 of the supplementary information). The Cu2(OH)2CO3 tablet before and after use showed no significant variation, as depicted in Fig. 1b. This comparison diagram indicates that the basic copper carbonate tablet block electrode is superior to loading the powder on carbon paper. Additional, physical properties of Cu2(OH)2CO3 are shown in Text S13. When the Cu2(OH)2CO3 tablet is used directly as the anode, the hydrogen peroxide yield (1065.16 μM) under a potential of 3.4 V vs. RHE is significantly improved compared to loading on carbon paper (26.32 μM), as shown in Fig. 1c, with a 40-fold improvement. It is confirmed that the product is indeed hydrogen peroxide, as discussed in Text S6 and Text S7 of the supplementary information.

Fig. 1: Performance of Cu2(OH)2CO3 tablet electrode in H2O2 production by water oxidation pathway.
figure 1

a Comparison of Cu2(OH)2CO3 powder loading on carbon paper as electrode before (left) and after (right) reaction. b Cu2(OH)2CO3 tablet as electrode before (left) and after (right) reaction. c H2O2 yield in 50 mL 0.5 M Na2SO4 electrolyte at 3.4 V vs. RHE. d H2O2 yield in 50 mL 0.5 M Na2SO4 electrolyte under different potential. e H2O2 generation rate and faradic efficiency of Cu2(OH)2CO3 tablet under different potential. f I-t curves and FE at 30 min in 8 h of Cu2(OH)2CO3 tablet in 50 mL Na2SO4 electrolyte by refreshing electrolyte per 30 min (16 cycles). Error bars represent the standard deviation of three measurements. Source data for Fig. 1c–f are provided as a Source Data file. The voltage is not iR corrected in all electro-experiments. The mass of the active materials in all comparative experiments is 2 g.

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Furthermore, the H2O2 yield under different potentials using a basic copper carbonate tablet anode in a 50 mL 0.5 M Na2SO4 electrolyte solution is shown in Fig. 1d. The hydrogen peroxide yield concentration increases with the voltage increase due to the increase in current (Fig. S1). However, when the potential rises to 3.6 V vs. RHE, the hydrogen peroxide yield concentration decreases because the excessive voltage may decompose a large amount of the generated hydrogen peroxide and oxygen evolution. The highest hydrogen peroxide yield concentration is achieved at 3.4 V vs. RHE (1065.16 μM) in the voltage range of 1.8–3.6 V vs. RHE. The corresponding hydrogen peroxide generation rate under different potentials are shown in Fig. 1e. The generation rate at 3.4 V vs. RHE high to 64.35 μmol h−1, and Faradaic efficiency high to 66.76% at 1.8 V vs. RHE. These results are higher than existing reports related to non-carbonate researches (as listed in Table S1). Additionally, experimental results from 16 cycles of 8-hour operation have demonstrated the stability of the performance of Cu2(OH)2CO3 tablet. It is illustrated in Fig. 1f. The current values monitoring in 8 h are stabilized at 20 ~ 22 mA. The FE values are basically stable between 20 ~ 30%. It reflects that the Cu2(OH)2CO3 tablet can be stable in Na2SO4 electrolyte in water oxidation reaction to hydrogen peroxide product process.

Furthermore, we also investigated the performance of Cu2(OH)2CO3 tablet electrodes for water oxidation to produce hydrogen peroxide in other electrolyte solutions without carbonate, such as NaNO3, Na2SiO3, and NaAc solutions. The results are depicted in Fig. S11 and Fig. S12. It indicates that Cu2(OH)2CO3 tablet electrode can approach the production of hydrogen peroxide through water oxidation in both NaAc and NaNO3 solutions. By comparing the current in these four non-carbonate salt electrolyte solutions and the pH values listed in Table S6, it can be observed that our Cu2(OH)2CO3 tablet electrode is more suitable for operation in neutral electrolytes.

Characterization of Cu2(OH)2CO3 tablet

The morphology of the Cu2(OH)2CO3 tablet is characterized by the SEM images, as shown in Fig. 2a. The filaments in all images originate from the PTFE binder, which acts to better bond the basic copper carbonate microspheres together so that they can be pressed into a tablet solid. There is no significant change in the size of the microspheres before and after the reaction, and they are all about 250 nm. The overall morphology and arrangement of the microspheres do not change significantly before and after the reaction. Figure 2b shows the comparison of FTIR spectra of Cu2(OH)2CO3 tablets before and after use. The emerging band at 1379, 558, 497, 428 cm−1 are assigned to SO3·H2O35,36,37,38,39. The emerging bands at 1090 and 506 cm−1 correspond to γ3-SO42− vibration which originates from the residual SO42− from aqueous electrolyte40,41. The band at 1205 cm−1 is owing to the antisymmetric stretching vibration of CF2, which is from the addition of PTFE binder in tablet anode42,43,44. The band at around 1150 cm−1 is attributed to the superposition of symmetric stretching vibration of CF2 and γ3-SO42− from residual Na2SO4 crystal40,42,43. The crystallinity of the tablet is characterized by XRD which shown in Fig. 2c, the tablet before and after reaction show typical 2θ peaks at 14.8°, 17.6°, 24.1°, 31.2° and 35.6°, which are assigned to the (0 2 0), (1 2 0), (2 2 0), (2 0 -1), (2 4 0) planes of Cu2(OH)2CO3 (PDF#41-1390). It proves that the composition in the pressed solid anode is indeed basic copper carbonate. Fig. 3.

Fig. 2: Characterization of Cu2(OH)2CO3 tablet.
figure 2

a SEM images of Cu2(OH)2CO3 tablets before (left) and after (right) reaction. b Fourier infrared absorption spectra (FTIR) of Cu2(OH)2CO3 tablets before and after reaction in the wave number range of 400–4000 cm−1. c XRD patterns of Cu2(OH)2CO3 tablets in the 2-theta range of 5–90o. d S 2p high-resolution XPS survey spectra of Cu2(OH)2CO3 tablets before and after reaction. e C 1 s high-resolution XPS survey spectra of Cu2(OH)2CO3 tablets before and after reaction. f Cu 2p high-resolution XPS survey spectra of Cu2(OH)2CO3 tablet before and after reaction. Source data for Fig. 2b–f are provided as a Source Data file.

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Fig. 3: Electro-chemical performance of Cu2(OH)2CO3 tablet.
figure 3

a LSV curves recorded in 2 M KHCO3 at a scan rate of 50 mV s−1 of Cu2(OH)2CO3 tablet anode. b Galvanostatic charge/discharge behaviors of Cu2(OH)2CO3 tablet and graphite oxide with an applied constant current density of 1 (Cu2(OH)2CO3 tablet) and 0.5 (graphite) mA cm−2 measured by Chronopotentiometry measurement. c Nyquist plots of Cu2(OH)2CO3 tablet anode in 50 mL 0.5 M Na2SO4 with different pH and corresponding fitting circuit diagram. d LSV curves were recorded in 0.5 M Na2SO4 electrolyte with or without 30 ppm H2O2 addition at a scan rate of 50 mV s−1 of Cu2(OH)2CO3 tablet. e Electro-decomposition rate of H2O2 on Cu2(OH)2CO3 tablet anode at different potentials in 0.5 M Na2SO4. f Self-decomposition rate constants of H2O2 with 1 mM initial concentrations in 2 M KHCO3 or 0.5 M Na2SO4 solution. pH error bars represent the standard deviation of three measurements. Acidic (4.17 ± 0.1), basic (11.33 ± 0.07), and the initial pH of 0.5 M Na₂SO₄ (7.15 ± 0.15). Source data for Fig. 3a–f are provided as a Source Data file. The voltage is not iR corrected in all electro-experiments.

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Figure 2d–f are S 2p, C 1 s, and Cu 2p high-resolution XPS survey spectra of Cu2(OH)2CO3 tablets before and after reaction, respectively. Since there is no S element on the Cu2(OH)2CO3 tablet anode before the reaction, the S 2p high-resolution XPS survey spectrum here only has the result after the reaction (Fig. 2d). The S 2p spectrum is fitted into two peaks at the binding energy of 167 eV and 169 eV, which are corresponding to SO3 species and metal sulfate45,46. The existence of SO3 species corresponds to the occurrence of SO3 after the reaction shown in Fig. 2b, which confirms the existence of SO3 · H2O residue after the reaction. The emergence of metal sulfate is due to the residual electrolyte sodium sulfate. The C 1 s spectra are fitted into three peaks with the binding energies of 284.5, 288.5, and 291 eV as shown in Fig. 2e, which can be assigned to elementary species, CO32− carbon species, and C-F2, respectively47,48,49,50. Among them, elementary species refers to the skeleton carbon such as CH2, which and C-F2 are derived from the binder PTFE in the tablet. CO32− is origin from basic copper carbonate. Carbon in PTFE accounts for a large proportion of the overall C 1 s spectra because PTFE is an organic substance, which contains a large amount of carbon. While basic copper carbonate is an inorganic substance, so the proportion of carbon is relatively small. The C 1 s spectra before and after the reaction present similar results. In the Cu 2p spectra as depicted in Fig. 2f, Cu 2p normally spin-orbit couples into two peaks as well as several shake-up satellite peaks at 941 eV, 944 eV, and 963 eV. The binding energy of Cu 2p3/2 and Cu 2p1/2 at around 935 eV and 955 eV are addressed to copper ions Cu2+, which confirmed the existence of Cu2(OH)2CO351,52,53. There is no emerging peak after the reaction in the C 1 s spectrum. The above characterization results show that the surface electronic states and the chemical elemental composition of the Cu2(OH)2CO3 tablet are stable.

Electrochemical performance of Cu2(OH)2CO3 tablet

LSV curves recorded in 0.5 M Na2SO4 at the scan rate of 50 mV s−1 are shown in Fig. 3a. It shows that the current increases dramatically when it exceeds 1.8 V vs. RHE. Figure 3b is the capacitance-potential (C-P) curves of Cu2(OH)2CO3 tablet and graphite concluded by Chronopotentiometry measurement. It is used to determine the electrochemically active surface areas (EASA) of the Cu2(OH)2CO3 tablet and graphite anode. Specifically, a constant current density of 1 mA was applied to the Cu2(OH)2CO3 tablet with a potential window of 1.2 V, while the graphite anode was carried out with a constant current density of 0.5 mA in the potential windows of 0.4 V. The following equation can be used to estimate the specific capacitance C of Cu2(OH)2CO3 tablet and graphite anode from the potential-time curves54,55.

$$C=\frac{I*\Delta t\,}{\Delta V}$$
(2)

Where I present the current density during Chronopotentiometry measurement processes, t presents the time, and ΔV presents the potential window.

During this measurement, I and ΔV are considered constants. After substituting the corresponding values into the above Eq. (2), the specific capacitance C of Cu2(OH)2CO3 tablet and graphite can be calculated to be 16.67 and 50 mF, respectively. Usually, the electrochemically active surface areas can be calculated by using a ratio between capacitance and specific surface area. Taking a value of 14 μF cm−2 54, the electrochemical specific surface area of the Cu2(OH)2CO3 tablet and graphite is 1189 and 3571 cm2 per geometric cm2, respectively. Although the electrochemically active surface areas of graphite are much larger than that of tablet basic copper carbonate. However, a relative experiment shows that graphite cannot produce hydrogen peroxide through water oxidation in Na2SO4 electrolyte solution. Therefore, even if graphite has large electrochemically active surface areas, there is no active site that can produce hydrogen peroxide in Na2SO4 solution. Figure 3c shows the Electrochemical impedance of Cu2(OH)2CO3 tablet in Na2SO4 with different pH. The AC impedance is smaller in an environment with a pH of 11.33 than in an environment with a pH of 4.17, which may be because there are more mobile ions in the environment of 11.33. The AC impedance was fitted with the analog circuit illustrated in Fig. 3c. The fitted solution resistance (R1) and charge transfer resistances (R2) in the Nyquist plot were 1258 ± 10 Ω (11.33) and 14978 ± 38 Ω (4.17), respectively.

Besides, the decomposition of H2O2 on the electrode surface is investigated by the following experiments according to related references 56,57. Specifically, the LSV curves of the basic copper carbonate tablet anode were scanned in the Na2SO4 electrolyte solution with or without 30 ppm H2O2 addition, and the obtained results are shown in Fig. 3d. The I-V curve exhibits a negative shift when with 30 ppm H2O2 addition compared with the curve obtained without H2O2 addition. It is due to the anodic oxidation decomposition of hydrogen peroxide on the electrode surface. Figure 3e is the electro-decomposition rate of H2O2 on Cu2(OH)2CO3 tablet anode at different potentials. This calculation refers to the method in the relevant literature58 (see Text S8 of supplement information). It shows that the electro-decomposition rate on Cu2(OH)2CO3 tablet anode increases with the increase of voltage. This is why the concentration of H2O2 decreases when the voltage is too high. The difference in the production of hydrogen peroxide in the carbonate-free electrolyte system comes from the difference in the active site, which is not found in graphite.

Besides, Na2SO4 solution as an electrolyte is still superior to KHCO3 because the self-decomposition ratio in Na2SO4 is much smaller than in KHCO3. The experimental details are shown in Text S5. The results shown in Fig. 3f indicates that the decomposition rate constants are determined to be 0.098 h−1 in KHCO3, which is indeed smaller than 0.002 h−1 in Na2SO4. Subsequently, to further explore the mechanism of tablet basic copper carbonate producing hydrogen peroxide in Na2SO4 solution, the following series of experiments are carried.

Influence of carbonate in the liquid phase

To determine whether the H2O2 produced by Cu2(OH)2CO3 tablet is derived from the effect of carbonate dissolved in the solution, a control experiment is conducted. 2 g Cu2(OH)2CO3 powder and the Cu2(OH)2CO3 tablet are put into two separate 50 mL of 0.5 M Na2SO4 solutions. These two solutions remain colorless after 2 h, as shown in Fig. 4a. Solid basic copper carbonate powder is visible in the aqueous phase. Copper ions in aqueous solutions typically appear green, but no green color is observed in this experiment. Additionally, UV-vis adsorption of solutions at different times showed no peaks, and the entire curve was almost horizontal (Fig. S3). By contrast, the UV spectra of the Cu2+ solution in Fig. S4 show obvious peaks in the same wavelength scanning range, indicating that Cu2+ is present in the solution. These results suggest that basic copper carbonate is not dissolved in a neutral Na2SO4 solution.

Fig. 4: Investigation of the active component of the Cu2(OH)2CO3 tablet Electrode.
figure 4

a Picture of 50 mL 0.5 M Na2SO4 solution with Cu2(OH)2CO3 tablet or Cu2(OH)2CO3 powder addition. b H2O2 production concentration of graphite under the potential of 3.4 V vs. RHE in different electrolyte. c H2O2 concentration of Cu2(OH)2CO3 tablet in different concentration of Na2SO4 electrolyte under the potential of 3.4 V vs. RHE. d H2O2 concentration of Cu2(OH)2CO3 tablet in Na2SO4 solution with different pH under the potential of 3.4 V vs. RHE. e FTIR spectra of Cu2(OH)2CO3 obtained from Isotope O18 tracing experiments. f H2O2 generation properties of Cu2(OH)2SO4 tablet in 50 mL Na2SO4 under 3.4 V vs. RHE. The results in Fig. 4b–f are all obtained from a 50 mL electrolyte system. Source data for Fig. 4b–f are provided as a Source Data file. The voltage is not iR corrected in all electro-experiments. The mass of the active materials in all comparative experiments is 2 g.

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Furthermore, another control experiment is conducted using a graphite plate as the anode and 2 g basic copper carbonate powder directly dispersed into the 50 mL 0.5 M Na2SO4 electrolyte solution in the anode cell under the same potential of 3.4 V vs. RHE. It is known that graphite can oxidize water to produce H2O2 in an electrolyte system with carbonate, which reached 832 μM under the potential of 3.4 V vs. RHE at 120 min in 50 mL KHCO3 as shown in Fig. 4b. However, the results showed that no H2O2 was produced in this Na2SO4 and basic copper carbonate powder mixture electrolyte (Fig. 4b). It indicates that basic copper carbonate and sodium sulfate solution cannot produce carbonate that has an effectual effect in the water oxidation process to produce hydrogen peroxide. In summary, it is confirmed that the production of H2O2 in a Na2SO4 electrolyte solution by a Cu2(OH)2CO3 tablet is not derived from carbonate in the solution.

Experiments are conducted using a basic copper carbonate tablet as an anode in different concentrations of Na2SO4 at 3.4 V vs. RHE. The results shown in Fig. 4c indicate that the hydrogen peroxide yield increases with an increase in the concentration of Na2SO4 electrolyte. It suggests that sulfate in electrolyte may play a role in the production of hydrogen peroxide in the system, and a higher concentration of sulfate is more conducive to the production of hydrogen peroxide.

Another difference between Na2SO4 and KHCO3 solutions is the pH of the solution. An alkaline electrolyte is known to be conducive to water oxidation to produce hydrogen peroxide. To investigate the effect of pH of the electrolyte on the electro-synthesis of the hydrogen peroxide process on the basic copper carbonate tablet, experiments were designed. The pH of the 0.5 M Na2SO4 solution in the anode cell was adjusted to 11.33 by NaOH, but no hydrogen peroxide production was observed in this non-carbonate basic electrolyte by graphite anode, as shown in Fig. 4b. It indicates that for graphite anodes, only an alkaline environment is not enough, and the presence of carbonates is necessary to water oxidation to produce hydrogen peroxide. By contrast, the basic copper carbonate tablet can produce hydrogen peroxide in a system without carbonate in a neutral Na2SO4 electrolyte.

The effect of the pH of the electrolyte solution on the production of hydrogen peroxide on a Cu2(OH)2CO3 tablet is investigated (Fig. 4d). At pH = 11.33, the yield of hydrogen peroxide is almost the same as that when the pH is not adjusted. At pH = 4.17 by H2SO4, the basic copper carbonate sheet could still produce hydrogen peroxide in an acidic electrolyte, although the yield of hydrogen peroxide is reduced by half to 501 μM at 120 min. The basic copper carbonate tablet can produce H2O2 in a solution without carbonate, but the acidic environment makes the production of H2O2 less than that of neutral or alkaline, which is consistent with previous literature conclusions56,57. Because the 2e− WOR water oxidation process will release H+ (Eq. (3)). The presence of excessive products can hinder the forward progress of the reaction. Therefore, Eq. (3) can reflect that the presence of excessive H+ is unfavorable for the formation of hydrogen peroxide. Moreover, the results of Nyquist plots show that the AC impedance of the Cu2(OH)2CO3 tablet anode in the acidic electrolyte is larger than that in the alkaline electrolyte. It suggests that alkaline electrolytes can neutralize the generated H+ from water oxidation21, which is beneficial to the forward progress of the reaction. If it is an acidic electrolyte, the generated H+ accumulates continuously, which is not conducive to the forward progress of the reaction. From this perspective, it will be more difficult to produce hydrogen peroxide by water oxidation in an acidic electrolyte.

$${2{{{\rm{H}}}}}_{2}{{{\rm{O}}}}\to {{{{\rm{H}}}}}_{2}{{{{\rm{O}}}}}_{2}+{2{{{\rm{H}}}}}^{+}+{2{{{\rm{e}}}}}^{-}$$
(3)

Furthermore, graphite as the anode could not produce hydrogen peroxide in Na2SO4 solutions at any pH, as shown in Fig. 4b and Table S5. It could only produce H2O2 in KHCO3 solution. It indicates that the difference in the production of hydrogen peroxide in the carbonate-free electrolyte system comes from the difference in the active site, which is not found in graphite.

Influence of composition in Cu2(OH)2CO3 tablet solid anode

The role of the compositions of basic copper carbonate tablet block in the 2e− WOR process to produce hydrogen peroxide is further investigated. The following experiments are conducted and the results are shown in Table S4. Firstly, to study the role of the carbonate structure in Cu2(OH)2CO3 tablet solid anode, CaCO3 was used as the anode for comparison. As the structure of CaCO3 is similar to that of the copper carbonate in basic copper carbonate. This result shows that the CaCO3 anode cannot catalyze water oxidation to produce H2O2, which indicates that the CO3 part in the solid anode should not play a key role in the production of hydrogen peroxide in a non-carbonate electrolyte system. Subsequently, the role of Cu part is investigated. We utilize CuO as the anode in a comparative experiment. The result shows that CuO as an anode could not catalyze water oxidation to produce H2O2 either. It turns out that it may not be Cu part in Cu2(OH)2CO3 that determines this process. However, the previous results showed that there is indeed an active part in the basic copper carbonate. Therefore, we believe that the OH part of the Cu2(OH)2CO3 tablet plays a certain role in the H2O2 production by 2e− water oxidation process.

Evidence of cyclic exchange between O in crystal OH and O in H2O

To demonstrate the role of the crystal OH component, we have designed three sets of experiments with the following thematic approach:

Firstly, to demonstrate the exchange between the O in H2O and the O within the crystal of Cu2(OH)2CO3 (Isotope O18 tracing method). Results are shown in Table 1, the results indicate that after the electrochemical reaction in the O18 electrolyte, samples were thoroughly washed, and O18 was detected in the electrolyte after subsequent electrochemical reaction in the O18-free electrolyte. This observation suggests the presence of O18 alternating between the Cu2(OH)2CO3 tablet and the electrolyte solution. Furthermore, samples that were soaked for the same duration in the O18 electrolyte, followed by thorough washing and subsequent electrochemical reaction in the O18-free electrolyte, did not exhibit the presence of O18 in the electrolyte. It eliminates the potential influence of adsorption effects during the experimental process. These results provide evidence that O18 has indeed entered the crystal of Cu2(OH)2CO3. It is demonstrated that the O from the Cu2(OH)2CO3 tablet electrode undergoes a cyclic exchange process with the O in H₂O.

Table 1 O18 abundance results
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Secondly, to prove that the exchanged O originates from the OH component within the crystal of Cu2(OH)2CO3. GC-MS analysis confirms the absence of O18 in the CO2 gas produced during high-temperature decomposition, while FTIR analysis demonstrates the presence of O18H in the crystal OH groups in Cu2(OH)2CO3. The results are shown in Fig. 4e and Fig. S9 in Supporting information. Using PY-GCMS analysis, it was determined that there was no presence of O18 in the generated gas (Fig. S9). The gas produced during the high-temperature formation of Cu2(OH)2CO3 contained CO2, indicating that the oxygen in CO2 did not contain O18. The O18 in CO2 originated from the CO3 component within Cu2(OH)2CO3, suggesting that the CO3 component did not contain O18. Combining this with experimental findings from isotope O18 tracing experiments, where O18 was found to enter the lattice of Cu2(OH)2CO3, it can be inferred that the O18 is located in the O of the OH component within the lattice of Cu2(OH)2CO3. For FTIR spectrum (Fig. 4e), specific chromophores in molecules exhibit unique absorption capabilities at different wavelengths. The absorption wavelength of a specific vibrational mode depends on the mass of its atoms, which can be altered by substituting atoms with isotopes of different masses59. O18 significantly alters the absorption wavelength of this isotope-edited vibrational mode, allowing for the examination and orientation of individual transition dipole moments, with a negative frequency shift59,60,61. The spectrum of Cu2(OH)2CO3 displays characteristic peaks in the range of 2500–4000 cm−1, which can be assigned to the single-bond OH groups in the Cu2(OH)2CO3 crystal62,63. It is observed in the Fig. 4e that the absorption peak at 3316 cm−2 is not prominent in the spectrum of #Soaking sample, but is pronounced in the spectrum of #Electro sample at that position. Introduction of O18, as indicated Fig. 1 in Reference59, results in a negative shift of the peak, suggesting that this is a consequence of O18 entering the crystal OH portion.

Finally, to verifying experiments, a confirmatory experiment to show that if the OH component is active rather than the O within the CO3 component, then under similar conditions, replacing CO3 with other components should exhibit similar hydrogen peroxide production performance. As shown in Fig. 4f, Cu2(OH)2SO4 tablets as the anode can produce hydrogen peroxide in Na2SO4, consistent with our expected results. Furthermore, the hydrogen peroxide yield is nearly identical to that of Cu2(OH)2CO3. It suggests that replacing the CO3 component with SO4 can also lead to hydrogen peroxide production when the results are similar, indicating that CO3 is not a critical component in the hydrogen peroxide production process. In addition, DFT calculation reveals that the SO4* cannot form HSO5 intermediates if removing the OH from the Cu2(OH)2CO3 structure (Fig. S17a), while can observed in the presence of OH (Fig. S17b). Combining the results above, it can be concluded that the OH component within the crystal undergoes cyclic exchange with O in H2O.

DFT calculation of the interaction between Cu2(OH)2CO3 and the electrolyte

Since CuO and CaCO3 cannot produce hydrogen peroxide in Na2SO4 electrolyte solution under the same conditions. Then, it is speculated that the OH part plays an important role in the water oxidation process. Therefore, density functional calculation is conducted to study the interaction between Cu2(OH)2CO3 and the electrolyte. The electrolyte mainly consists of OH− and SO42− ions. We study the OH− from the electrolyte. According to calculation by density functional theory, there are two cases for the adsorption of OH* from the electrolyte on the substrate: one with the OH* adsorbed on the Cu atoms (Fig. 5a), and the other with the OH* adsorbed on the O atoms via hydrogen bond (Fig. 5b). The adsorption energy is calculated by the following Eq. (4):

$${E}_{{ad}}={E}_{{tot}}-{E}_{{substrate}}-{E}_{OH}$$
(4)

where Ead is the adsorption energy, Etot is the total energy of the molecule on the substrate, Esubstrate is the energy of the substrate, and EOH is the energy of the OH adsorbed. The Ead for the above-mentioned two cases is -3.3 eV and -4.8 eV, respectively. The negative value suggests that these two cases are possible to occur.

Fig. 5: Direct pathway via OH* and indirect pathway via SO4*.
figure 5

a Structure of OH* adsorbed on Cu atom and O atom on Cu2(OH)2CO3 before structure relaxation. b Structure of OH* adsorbed on Cu atom and O atom on Cu2(OH)2CO3 after structure relaxation. c Structure of SO4* adsorbed on Cu2(OH)2CO3 before structure relaxation. d Structure of SO4* adsorbed on Cu2(OH)2CO3 after structure relaxation.

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According to previous reports, the change of Gibbs free energy of *OH (ΔGOH*) is a crucial parameter for H2O2 production. The optimized range for directing oxidation of OH is between 1.6 to 2.4 eV19. However, the ΔGOH* is calculated to be about 0.70 eV when OH* is adsorbed on Cu, and 0.82 eV when OH* is adsorbed on O. These two values are not in the optimum range of 1.6 to 2.4 eV, suggesting directing oxidation of OH should not be the pathway for the generation of H2O2.

In addition to OH−, SO42− ions are the dominating ions in the electrolyte. Therefore, we further investigate the interaction between SO42− and the substrate. As shown in Fig. 5c, an SO42− ion is put near the substrate. Then, after structure relaxation, the structure of the SO42− total changes (Fig. 5d). It is clear that the four O atoms rearrange. Interestingly, two O atoms link with each other to form the O-O bond, which is a very important bond for the production of H2O2. Importantly, some O atoms from the SO42− are bonded with the OH from the substrate Cu2(OH)2CO3, suggesting there are interactions between the SO42− and the Cu2(OH)2CO3 (Supplementary Data 1). Importantly, the O-O bond distance is found to be about 1.47 Å, which is very close to the O-O bond distance of 1.48 Å in H2O2. Therefore, this rearrangement of the SO42− should be a crucial step for the production of H2O2. Therefore, to conclude, the generation of H2O2 should not be a direct pathway via OH* adsorption. Instead, it should be an indirect pathway via SO42−.

Crystal OH mediating pathway for hydrogen peroxide production

The O18 isotope tracing series of experiments demonstrated that the O from the crystal OH in the Cu2(OH)2CO3 tablet electrode undergoes a cyclic exchange process with the O in H₂O. DFT calculations indicate that this process occurs via an indirect pathway mediated by SO42− in the solution. To this end, we propose a crystal OH mediating pathway for hydrogen peroxide production via two-electron water oxidation. The crystal OH in the tablet electrode interacts with SO₄ in the electrolyte to form SO₄OH* intermediate. This intermediate subsequently decomposes into HO₂* and SO₃. The SO₃ intermediate contacts with H₂O, releasing two H⁺ ions and SO₄ back into the solution. HO₂* combines with H from the H2O to generate hydrogen peroxide, while the remaining OH returns to the crystal structure of the electrode, regenerating the Cu2(OH)2CO3.

Figure 6 provides a detailed explanation of this process of H2O2 generation on Cu2(OH)2CO3. Firstly, the SO42− is adsorbed near the OH of the Cu2(OH)2CO3. The structure of the SO42− total change. It is clear that the four O atoms rearrange. Interestingly, two O atoms link with each other to form the O-O bond, which is a very important bond for the production of H2O2. The O-O bond distance is found to be about 1.47 Å, which is very close to the O-O bond distance of 1.48 Å in H2O2. This process is shown as Eq. (5) below. Therefore, this rearrangement of the SO42− should be a crucial step for the production of H2O2.

$${{{{{\rm{Cu}}}}}_{2}({{{\rm{OH}}}})}_{2}{{{{\rm{CO}}}}}_{3}+{{{{{\rm{SO}}}}}_{4}}^{2-}={{{{\rm{Cu}}}}}_{2}({{{\rm{OH}}}}){{{{\rm{CO}}}}}_{3}-{{{{{\rm{HSO}}}}}_{5}}^{2-}$$
(5)
$${{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}}){{{\rm{CO}}}}}_{3}-{{{{\rm{HSO}}}}}_{5}={{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}}){{{\rm{CO}}}}}_{3}+{{{{\rm{HO}}}}}_{2}-{{{{\rm{SO}}}}}_{3}$$
(6)
$${{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}}){{{\rm{CO}}}}}_{3}+{{{{\rm{HO}}}}}_{2}-{{{{\rm{SO}}}}}_{3}={{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}}){{{\rm{CO}}}}}_{3}+{{{{\rm{HO}}}}}_{2}+{{{{\rm{SO}}}}}_{3}$$
(7)
Fig. 6: Diagram of H2O2 generation process on Cu2(OH)2CO3.
figure 6

a Structural model of SO4 on the surface of Cu2(OH)2CO3. b SO4 combines with OH in Cu2(OH)2CO3 to form HSO5. c HSO5 structural reorganization forming HO2-SO3. d The bond in HO2-SO3 breaks, resulting in the formation of HO2 and SO3 intermediates. e HO2 and SO3 interact with two H2O molecules to generate H2O2 while releasing two H+, returning the SO4 and Cu2(OH)2CO3 original structure.

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Subsequently, the HSO5 part release from the Cu2(OH)CO3-HSO5 and causes the following formation of HO2 and SO3 immediately. The process is demonstrated as Eqs. (6) and (7).

Then, two H2O will take part in the reaction and produce H2O2, SO42− and 2H+. In this step, one OH from H2O will fill into Cu2(OH)CO3 and form Cu2(OH)2CO3 again. The process is illustrated by Eq. (8).

$${{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}}){{{\rm{CO}}}}}_{3}+{{{{\rm{HO}}}}}_{2}+{{{{\rm{SO}}}}}_{3}+{2{{{\rm{H}}}}}_{2}{{{\rm{O}}}}= {{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}})}_{2}{{{{\rm{CO}}}}}_{3}+{{{{\rm{H}}}}}_{2}{{{{\rm{O}}}}}_{2}+{{{{{\rm{SO}}}}}_{4}}^{2-}\\ +{2{{{\rm{H}}}}}^{+}+{2{{{\rm{e}}}}}^{-}$$
(8)

According to our calculation, the change of energy is about 2.625 eV. At the applied potential of 2.4 V vs. RHE, the change of energy is calculated to be about -2.17 eV, suggesting the reaction is thermodynamically favorable at this potential.

Therefore, to conclude, the generation of H2O2 should not be a direct pathway via OH* adsorption. Instead, it should be an indirect pathway via SO42−, and the SO42− and the OH from the Cu2(OH)2CO3 contribute simultaneously to the formation of H2O2. The total reaction is expressed as Eq. (9):

$${{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}})}_{2}{{{{\rm{CO}}}}}_{3}+{{{{{\rm{SO}}}}}_{4}}^{2-}+{2{{{\rm{H}}}}}_{2}{{{\rm{O}}}}= {{{{\rm{Cu}}}}}_{2}{({{{\rm{OH}}}})}_{2}{{{{\rm{CO}}}}}_{3}+{{{{\rm{H}}}}}_{2}{{{{\rm{O}}}}}_{2}+{{{{{\rm{SO}}}}}_{4}}^{2-}\\ +{2{{{\rm{H}}}}}^{+}+{2{{{\rm{e}}}}}^{-}$$
(9)

Discussion

In this study, the Cu2(OH)2CO3 tablet anode is successfully prepared, which solved the problem that Cu2(OH)2CO3 powder is easy to fall off during electrochemical experiments. More importantly, the use of Cu2(OH)2CO3 tablet electrode successfully produced hydrogen peroxide by 2e− water oxidation reaction in a carbonate-free electrolyte solution Na2SO4 solution, and the hydrogen peroxide production rate is as high as 64.35 μmol h−1, and the Faradaic efficiency is up to 66.76%. Both the yield and Faradaic efficiency are higher than other the literature in non-carbonate electrolyte field. In addition, this study also verifies that the OH part in Cu2(OH)2CO3 and sulfate radicals in electrolytes play key roles in the production of hydrogen peroxide through a series of comparative experiments. It’s concluded that SO42− attaches near the OH groups of Cu2(OH)2CO3, causing a significant rearrangement of its structure and the Cu2(OH)CO3-HSO5 is formed. Moreover, the release of HSO5- from Cu2(OH)CO3-HSO5 prompts the formation of HO2 and SO3. O-O bond is crucial for H2O2 generation. The O-O bond distance of this HO2 is approximately 1.47 Å, similar to that in H2O2. Subsequently, two H2O molecules react, resulting in the production of H2O2, SO42−, and 2H+. During this step, one OH from H2O fills into Cu2(OH)CO3, return reforming Cu2(OH)2CO3. This method enables hydrogen peroxide production in a neutral setting, expanding its environmental application potential. Although current density limitations exist, future research aims to enhance performance through anode modifications like nano-engineering and heterojunction construction. This study can inspire people to pay more attention to the contribution of lattice elements to the two-electron water oxidation reaction.

Methods

Chemicals

Cu2(OH)2CO3 powder Cu2(OH)2CO3, KHCO3, CaCO3, KCl, Cu2(OH)2SO4 and DPD (N, N-diethyl-p-phenylene amine) are procured from Macklin Chemical Co. Na2SO4, Na2HPO4, NaH2PO4, and NaOH are purchased from Aladdin Chemical Co., Ltd. POD (horseradish peroxidase) with an enzyme activity of ≥ 300 U mg−1 is obtained from Sigma-Aldrich, Germany and stored in a dark environment at −2–8 °C. A 60% PTFE solution (Polytetrafluoroethylene) was purchased from Shanghai D&B Biological Science and Technology Co., Ltd. H2SO4 was procured from Merck, Germany, and copper oxide was purchased from Hebei Badu Metal Materials Co., Ltd. H2O18 was purchased from Energy Chemical Co., Ltd. All chemicals are analytical grade and used without further purification, and ultrapure water is utilized in all experiments. The graphite plate is obtained commercially.

Preparation of Cu2(OH)2CO3 tablet

A slurry was prepared by mixing 4 g of Cu2(OH)2CO3 and 1 mL of 60% PTFE, which was then dried in an oven at 70 °C for 12 h to obtain PTFE-treated powder. 2 g of the dried powder was pressed into a Cu2(OH)2CO3 pie-shaped tablet using a mold under a pressure of 10 kPa for 2 min, as depicted in Fig. 1b. For comparison, Cu2(OH)2CO3 powder was loaded onto carbon paper by ultrasonically dispersing 2 g of the powder evenly in 1 mL 60% PTFE and 100 μL 0.1% Nafion mixture solution for 5 min, followed by drying in an oven at 70 °C for 12 h to form the anode as depicted in Fig. 1a. Additionally, when attempting to press the basic copper carbonate powder that was not treated with PTFE into a pie-shaped tablet using the same conditions, it remained in a powder state and could not be formed into a tablet. The electrode is depicted as a circular pie-shaped tablet with a radius of 1 cm. Please see Text S14 in Supporting Information for more details.

Characterizations

H2O2 was quantified by the DPD-POD method19. Specifically, a 0.3 mL sample was added into a 1.7 mL buffer. After 50 μL of DPD and POD addition, the mixture was scanned for UV curve and the absorption peak at 551 nm was used as a sign of H2O2 quantification. The morphology of samples was characterized by scanning electron microscope conducted by a QUANTAN FEG250 scanning electron microscope (SEM) at the accelerating voltages of 15–20 kV. The attenuated total reflection Fourier transform infrared (ATR-FTIR) spectra were collected in the ATR-FTIR mode using a Multimode In-Situ Infrared Spectrometer (Nicolet iS50) which was equipped with Golden Gate ATR attachment with diamond crystal and OMNIC software in the range of 400–4000 cm−1. The X-ray diffraction (XRD) patterns were examined by using an UltimaIV X-ray diffractometer (Japan) with Cu Kα irradiation scan within 2-theta 5-90° to determine the crystal structure of the anode tablet. The X-ray photoelectron spectroscopy (XPS) was performed on a Thermo ESCALAB 250XI multifunctional imaging electron spectrometer (American Thermal Power, USA) to determine the surface electronic states and the chemical elemental composition of the Cu2(OH)2CO3 tablets before and after the reaction. The instrument model of the stable water isotope analyzer is Picarro2140-i.

Electrochemical measurement

Preparation of Na2SO4 electrolyte solution

71.021 g of Na2SO4 powder was dissolved in ultrapure water using magnetic stirring until fully dissolved. Then it was transferred to a 1 L volumetric flask and dilute to a final volume of 1 L. This solution is a 0.5 M Na2SO4 solution. The prepared electrolyte solution should be stored at room temperature in the dark as a reserve solution. Normally, if it is used within three days, a fresh solution will be prepared. If it is not used within three days, a fresh solution will be re-prepared.

Electrochemical measurement

The brand of the quartz electrochemical cell is Gaoshiruilian. The brand of the reference electrode is Chenhua. Specifically, the electrochemical system setup involved an H-type quartz electrolysis cell separated by Nafion 117 membranes (DuPond Co., the radius is 1.5 cm and the thickness is 183 μm, and it is sandwiched between two rubber rings and secured in the middle of two quartz electrolysis cells.), with 50 mL of 0.5 M Na2SO4 solution on each side serving as the electrolyte. The electrochemical setup consisted of a three-electrode system, with a carbon electrode as the counter electrode, an Ag/AgCl electrode (saturated potassium chloride solution) as the reference electrode and Cu2(OH)2CO3 tablets as the working electrode. The experiments were conducted using a CHI 760E electrochemical workstation (Shanghai Chenhua Instrument Co., Ltd.). All tests were conducted at an ambient temperature of 25 °C. Potentials set against Ag/AgCl were converted to the RHE scale according to ERHE = EAg/AgCl + 0.197 + 0.059 × pH, and the pH values of electrolytes value were tested by a pH meter (Shanghai LeiMag Co., Ltd.). A graphite rod and a Ag/AgCl electrode (in a Luggin capillary) were used as the counter, and reference electrode, respectively. Cyclic voltammetry (CV) tests were carried out in this system at a scan rate of 50 mV s−1. The charging/discharging behavior of the electrode is measured through chronopotentiometry. The Nyquist plot of the electrode is conducted by using the A.C. Impedance mode on the software connected to the electrochemical workstation.

Reference electrode calibration

In a 0.5 M H₂SO₅ electrolyte solution saturated with H2, a Pt foil serves as the working electrode, while a carbon rod acts as the counter electrode, with an Ag/AgCl reference electrode immersed in saturated KCl solution. The scan rate is set at 50 mV s−1, and the onset potential for hydrogen evolution obtained from the monitored hydrogen evolution curve is considered the zero potential relative to the RHE.

Faradaic efficiency calculation details

The Faradaic efficiency calculation details are shown below:

$${FE}\left(\%\right)=\frac{m\, {{{\rm{n}}}}\, 96485}{Q}$$

Where, m represents the number of moles of produced hydrogen peroxide (mol), and n denotes the number of electrons transferred. This reaction is a two-electron process, thus n is a constant equal to 2. The value 96485 corresponds to Faraday’s constant. Q represents the amount of charge transferred during the reaction, calculated as the product of time t (s) and current I (A).

Details of the isotope O18 tracing experiment

The aim of this experiment is to demonstrate that the oxygen in water indeed enters the crystal structure of Cu₂(OH)₂CO₃. Experimental Group: The Cu₂(OH)₂CO₃ pellets will operate at a voltage of 2.8 V vs. RHE in an electrolyte solution containing O18 for 2 h. After thorough washing, they will be placed in an electrolyte solution without O18 and operated at the same voltage for another 2 h. If the O18 from the electrolyte enters the Cu₂(OH)₂CO₃ crystal, it should be released back into the electrolyte upon further pressurization. The electrolyte solution obtained from the second electrochemical experiment will be collected and characterized using stable isotope analysis to determine the presence of O18. Additionally, a control experiment will be conducted to eliminate any interference from adsorption effects. Control Group: The Cu₂(OH)₂CO₃ pellets will be soaked in an electrolyte solution containing O18 at room temperature for 2 h, followed by thorough washing. They will then be placed in an electrolyte solution without O-18 and operated at 2.8 V vs. RHE for 2 h. The electrolyte solution from the second operation will be analyzed for the presence of O18 isotopes.

Details for resistance measurements

The AC impedance fitting of sodium sulfate electrolyte solutions at different pH levels was performed. The Rs-(CPE-Rp) model from Instant Fit was selected, and the corresponding simulated circuit is illustrated in Fig. 3c for the fitting of the resistance parameters. The error bar is derived from the results of the cubic simulation.

Theoretical calculations

The mechanism for H2O2 generation is studied19,64. Projector-augmented-wave method with the Perdew–Burke–Ernzerh of GGA functional was used65,66,67. The electronic convergence limit was set to be 1 × 10−5 eV68. Optimization of atomic coordinates was considered to be converged when the Hellmann–Feynman force was smaller than 1 × 10−2 eV Å−1. The slab consisting of Cu2(OH)2CO3 was established. Intermediate like SO4* are placed on the top of the surface. The vacuum region is about 15 Å in height. Detailed information is presented in Supplementary Data 1.