Modulated metal-support interactions for efficient nitrate electroreduction at positive potentials

Abstract

Electrochemical nitrate upgrading presents a sustainable route for repairing unbalanced nitrogen cycle. However, the low energy efficiency caused by high overpotential impedes the industrialization progress. Herein, Ru clusters supported on metal hydroxide are constructed by a universal self-corrosion strategy, and the metal-support interaction is modulated to simultaneously optimize NO₃⁻ adsorption and water dissociation for achieving high energy efficiency at positive potentials. The Co(OH)₂-supported Ru with moderate metal-support interaction exhibits a high energy efficiency of 49.5% and a high NH₃ Faradaic efficiency of ~100% at positive potentials. Furthermore, a long-term stability over 1200 h is achieved at an industrial-scale current density of 200 mA cm⁻². Moreover, the assembled rechargeable hybrid battery system shows a great potential in waste upcycling and energy conversion. This work underscores the significance of metal-support interaction for promoting nitrate electroreduction at positive potentials.

Electrochemical nitrate upgrading is restricted by high overpotential with low energy efficiency. Here, the authors report a metal-support modulation strategy to simultaneously optimize nitrate adsorption and water dissociation for achieving high energy efficiency at positive potentials.

Introduction

Nitrogen serves as the fundamental of modern industry and agriculture. The artificial nitrogen cycle involves the redox transformations across N₂, NH₃, and NO_x, exerting a significant influence on the ecological environment and all of life^1–3. The manufacture of NH₃ from N₂ by the conventional Haber–Bosch process has promoted agricultural progress with boosted fertilizer production^4,5. Moreover, NH₃ also has wide industrial applications in chemical synthesis, clean fuels, and refrigerant⁶. During the agricultural production or industrial manufacture procedures, NH₃ can be transformed into NO₃⁻ with the release of environmentally hazardous wastewater⁷. Consequently, the accumulation of NO₃⁻ in wastewater with nitrogen enrichment can lead to the unbalanced nitrogen cycle (Fig. 1), threatening human life and the global ecosystem^8,9. Although several denitrification methods, such as landfill disposal and biological transformation, have been developed for treating NO₃⁻-rich wastewater, the high treatment expense and potential nitrate leakage are still the challenges¹⁰. Under these circumstances, electrocatalytic NO₃⁻ reduction reaction (NO₃⁻RR), motivated by renewable energy, presents an environmentally benign route for upgrading NO₃⁻ into NH₃^11–13. Beyond enabling the sustainable treatment of NO₃⁻-rich wastewater, the NO₃⁻RR process can also diminish the fossil fuel consumption and hydrocarbon reliance of the Haber–Bosch process for ammonia production¹⁴. Nevertheless, the operating potentials of most NO₃⁻RR electrocatalysts are below 0 V vs. reversible hydrogen electrode (RHE), far away from the theoretical potential (0.69 V vs. RHE), resulting in the low energy efficiency (<40%) associated with high energy consumption^15–17. In order to promote the industrialization progress of electrochemical nitrate upgrading, it is required to develop high-performance catalysts operating at positive potentials, for enhancing the energy efficiency of nitrate electroreduction and reducing energy consumption.

Fig. 1. Schematic diagram of the nitrogen cycle.

Electrochemical NO₃⁻ upgrading pathway for repairing the unbalanced nitrogen cycle.

The prerequisite of catalyst design is understanding the reaction mechanism of nitrate electroreduction. The first and most crucial step in determining whether NO₃⁻RR can proceed is the adsorption of NO₃⁻. It is worth noting that excessively strong adsorption induces active-site poisoning, whereas too weak adsorption impedes the following reduction steps^18–20. The adsorption strength depends on the electronic interaction between the electrocatalyst and NO₃⁻. Therefore, regulating the electronic structure of the electrocatalyst holds the key for achieving suitable NO₃⁻ adsorption capability. Following NO₃⁻ adsorption, the NO₃⁻RR pathway needs to go through consecutive nine steps of hydrogenation reactions, which require continuous active hydrogen (^*H) supply. In neutral and alkaline electrolytes, ^*H originates predominantly from interfacial water dissociation²¹. In fact, the water dissociation process is sluggish dynamically and requires overcoming a high thermodynamic barrier, restricting the ^*H supply and then impeding the NO₃⁻RR hydrogenation steps at low overpotentials^22–24. Rational regulation of the catalyst’s interfacial microenvironment can accelerate water dissociation kinetics, thereby boosting the hydrogenation process of electrochemical NO₃⁻RR²⁵. However, synergistic optimization of NO₃⁻ adsorption and water dissociation by a single regulation strategy is still challenging for catalyst design.

Supported metal catalysts exhibit extensive applications in the field of catalysis, where metal–support interactions drive interfacial charge redistribution, thereby modulating electronic structure, chemical adsorption behavior, and stability^26,27. Recently, supported metal catalysts have spurred increasing attention as NO₃⁻RR electrocatalysts. For example, Li et al. prepared Ru metals supported on nitrogen-doped carbon by electrochemical reduction²⁸. The electronic structure of Ru metal was tuned for repelling hydrated cations and loosening water structure, thus improving the NO₃⁻RR performance with the highest Faradaic efficiency at −1.1 V vs. RHE. Besides, Cho et al. reported the reconstruction-induced Cu metals loaded on Co(OH)₂ and the Co(OH)₂ support can accelerate the hydrogenated conversion from NO₃⁻ to NH₃, contributing an optimal Faradaic efficiency of 96.15% at −0.5 V vs. RHE²⁹. Despite the progress of supported metal catalysts, the ultra-negative potential can lead to low energy efficiency, and the influence of metal–support interaction on NO₃⁻RR performance is still unclear.

Herein, we demonstrate a universal self-corrosion strategy for synthesizing the hydroxide-supported metallic Ru catalyst, in which the metal–support interaction can be modulated for optimizing both NO₃⁻ adsorption and water dissociation processes. In-situ characterizations and theoretical calculations indicate that among different iron-group hydroxides, Co(OH)₂ shows moderate metal–support interaction with metallic Ru, which endows metallic Ru appropriate NO₃⁻ adsorption ability and accelerates the following hydrogenation steps. Benefitted from the modulated metal–support interaction, the metallic Ru supported on Co(OH)₂ (Ru-Co(OH)₂) electrocatalyst exhibits an energy efficiency of 49.5% at a positive potential of +0.2 V vs. RHE and a high NH₃ Faradaic efficiency of ~100% at +0.1 V vs. RHE. Furthermore, in a membrane electrode assembly (MEA) electrolyzer, the Ru-Co(OH)₂ electrocatalyst can keep relatively stable during the 1200 h of NO₃⁻RR stability tests at 200 mA cm⁻². Therefore, the Ru-Co(OH)₂ electrocatalyst establishes industrial-scale electrochemical NO₃⁻RR feasibility due to the outstanding performance at positive potentials, extraordinary stability, and high commercialization value.

Results

Material synthesis and characterizations

The Ru clusters loaded on metal hydroxide were prepared by the self-corrosion method (Fig. 2a). Motivated by the standard potential difference of redox species (Supplementary Table 1), the self-corrosion process can proceed spontaneously³⁰. Specifically, under the dual corrosion effect of Ru³⁺ and O₂, the metal atoms from the metal surface can be oxidized and dissolved as M²⁺ ions (M represents the metal support). For the cathode half-reactions, O₂ from the air was reduced to OH⁻ while Ru³⁺ was converted to metallic Ru. Moreover, M²⁺ ions can be coordinated with OH⁻ to form metal hydroxide (M(OH)₂), and metallic Ru atoms can be further deposited on the surface of M(OH)₂, leading to the synthesis of hydroxide-supported Ru clusters electrocatalysts. As a universal synthesis strategy, the self-corrosion method can be utilized for preparing noble metal clusters supported on different transition metal hydroxides. For investigating the metal–support interactions, Ru cluster loaded on various iron-group hydroxides (including Co(OH)₂, Ni(OH)₂, and Fe(OH)₂) were synthesized by using Co, Ni, and Fe foams as sacrificial metal source, respectively. To illustrate, Ru-Co(OH)₂ catalyst exhibits the microscopic nanosheet morphology of hydroxide as shown in the scanning electron microscope (SEM) image (Fig. 2b). Corresponding energy dispersive spectroscopy (EDS) elemental analysis reveals that the atomic content of Ru can reach 6.22% for the Ru-Co(OH)₂ electrocatalyst (Supplementary Fig. 1). Transmission electron microscopy (TEM) image shows ultra-fine Ru clusters uniformly dispersed on the nanosheet support (Fig. 2c). High-resolution TEM (HRTEM) combined with corresponding Fast Fourier transform (FFT) and inverse FFT (IFFT) images identify the exposed (101) plane of metallic Ru with the lattice spacing of 0.204 nm (Supplementary Fig. 2, Fig. 2d, f), and the lattice spacing of 0.237 nm is ascribed to the (101) plane of Co(OH)₂ (Fig. 2e, g). Selected area electron diffraction (SAED) pattern also identifies dual components of metallic Ru and Co(OH)₂ (Supplementary Fig. 3). High-angle annular dark-field scanning TEM (HAADF-STEM) and corresponding energy dispersive spectroscopy (EDS) mapping display uniform distributions of Ru, Co, and O elements (Fig. 2h). To demonstrate the universality of self-corrosion strategy, control samples with different hydroxide supports can also be successfully synthesized by a similar self-corrosion method (Supplementary Figs. 4 and 5). Compared with the Ru-Co(OH)₂ electrocatalyst, Ru-Ni(OH)₂ and Ru-Fe(OH)₂ electrocatalysts exhibit a similar structure of Ru cluster-incorporated metal hydroxide nanosheet arrays, but with different compositions. X-ray diffraction (XRD) patterns confirm the metal hydroxide crystalline phases across all electrocatalysts (Supplementary Fig. 6), which are consistent with TEM results. The surface elemental composition can be further identified by X-ray photoelectron spectroscopy (XPS) characterization (Supplementary Figs. 7 and 8). As shown in Fig. 2i, j, the negatively shifted Ru⁰ 3p binding energy and positively shifted Co²⁺ 2p binding energy demonstrate that the electrons transfer between Co(OH)₂ support and metallic Ru for the Ru-Co(OH)₂ electrocatalyst. The electronic metal–support interactions between Co(OH)₂ support and metallic Ru can effectively regulate the adsorption property of reactants and intermediates, thereby influencing the electrochemical performance of nitrate electroreduction.

Fig. 2. Synthesis and characterizations of the electrocatalyst.

a Schematic illustration of the self-corrosion method for synthesizing Ru clusters anchored on the metal hydroxides. b SEM, c TEM, d, e FFT (inset) and IFFT, f, g lattice spacing calculations, h HAADF-STEM and EDS mapping images of the Ru-Co(OH)₂ electrocatalyst. i Ru 3p XPS spectra of Ru-Co(OH)₂ and Ru. j Co 2p XPS spectra of Ru-Co(OH)₂ and Co(OH)₂.

Nitrate electroreduction performance measurements

Nitrate electroreduction performance was firstly evaluated by the standard three-electrode system in 1 M KOH containing 0.1 M KNO₃. As shown in Fig. 3a and Supplementary Fig. 9, compared to Ru-Ni(OH)₂ and Ru-Fe(OH)₂ electrocatalysts, the Ru-Co(OH)₂ electrocatalyst exhibits the most positive onset potential (0.49 V, approaching the standard potential of NO₃⁻RR) and the maximal increase of current density with the existence of NO₃⁻, illustrating its high NO₃⁻RR performance. The reduced Tafel slope illustrates the accelerated NO₃⁻RR kinetics of Ru-Co(OH)₂ (Supplementary Fig. 10). Furthermore, electrochemical chronoamperometry tests were conducted in a potential range of +0.3 V ~−0.3 V vs. RHE for 1 h (Supplementary Fig. 11). Electrolytes after chronoamperometry measurements were collected to quantify the N-containing products by ultraviolet-visible (UV-vis) spectrophotometry (Supplementary Figs. 12–15). As a result, Ru-Co(OH)₂ catalyst displays an energy efficiency of 49.5% at a positive potential of +0.2 V vs. RHE (Fig. 3b, Supplementary Fig. 16) and the highest NH₃ Faradaic efficiency of ~100% at +0.1 V vs. RHE (Fig. 3c), outperforming Ru-Ni(OH)₂ and Ru-Fe(OH)₂ electrocatalysts. Moreover, NH₃ yield rate increased proportionally to the applied negative potentials, and the Ru-Co(OH)₂ electrocatalyst exhibits the highest NH₃ yield rate of 45.4 mg h⁻¹ cm⁻² at −0.3 V vs. RHE (Fig. 3d), corresponding to an industrial-scale NH₃ partial current density of 572.32 mA cm⁻² (Supplementary Fig. 17). The Ru-Ni(OH)₂ electrocatalyst shows a lower NH₃ yield rate and a lower NH₃ partial current density than that of Ru-Co(OH)₂ across the wide potential range. Besides, Ru-Co(OH)₂ electrocatalyst displays the lowest Faradaic efficiency of NO₂⁻ (Supplementary Fig. 18), demonstrating its excellent reaction selectivity of NH₃ production. However, the Ru-Fe(OH)₂ electrocatalyst is restricted to the abundance of NO₂⁻ byproduct due to its weak ability for the hydrogenation process of NO₂⁻. Moreover, the NO₃⁻RR performance of Ru-Co(OH)₂ surpasses individual Ru and Co(OH)₂ electrocatalysts (Supplementary Fig. 19), illustrating the synergistic effect of metal and hydroxide support for boosting NO₃⁻RR performance. Base on above analysis, Ru-Co(OH)₂ exhibits the best comprehensive performance of NO₃⁻RR when screening different hydroxide supports for loading active Ru metal clusters. Particularly, Ru-Co(OH)₂ shows the nearly 100% Faradaic efficiency and recorded-high energy efficiency at positive potential, which is competitive with the state-of-art NO₃⁻RR electrocatalyst (Fig. 3e, Supplementary Table 2).

Fig. 3. Electrochemical NO₃⁻RR performance tests.

a LSV tests without iR compensation of different electrocatalysts in 1 M KOH with 0.1 M NO₃⁻. b Energy efficiency, c Faradaic efficiency, and d yield rate of NH₃ of different electrocatalysts at various potentials. The error bars derived from three independent measurements. e Comparison of the NO₃⁻RR performance of the Ru-Co(OH)₂ electrocatalyst in this work with state-of-art electrocatalysts, and the detailed data are listed in Supplementary Data 2. f Time-dependent N concentration of NH₃, NO₂⁻, and NO₃⁻ of Ru-Co(OH)₂ electrocatalyst at 100 mA cm⁻² for upgrading simulated waste water with 0.01 M NO₃⁻. g Schematic diagram of MEA-based electrolyzer. h Long-term stability tests of the Ru-Co(OH)₂ electrocatalyst at 200 mA cm⁻² for 1200 h using a MEA-based electrolyzer.

To confirm the adaptability in real waste water containing different NO₃⁻ concentrations, the nitrate electroreduction performance was further tested with the NO₃⁻ concentration range of 0.01–0.5 M. Linear sweep voltammetry (LSV) measurements demonstrate that the current density and NH₃ yield rate of Ru-Co(OH)₂ electrocatalyst are enhanced gradually with the increase of NO₃⁻ concentration (Supplementary Fig. 20a, b). The Faradaic efficiency for NH₃ production is closely related to the applied potential in different NO₃⁻ concentrations (Supplementary Fig. 20c). At +0.1 V vs. RHE, Ru-Co(OH)₂ electrocatalyst could sustain >90% NH₃ Faradaic efficiency over a broad NO₃⁻ concentration range of 0.01–0.1 M, and the reduced NH₃ Faradaic efficiency in 0.5 M NO₃⁻ can be attributed to the impeded conversion of NO₂⁻ to NH₃ in high NO₃⁻ concentration (Supplementary Fig. 20d). Moreover, continuous NO₃⁻RR experiments were further performed in the simulated nitrate-containing waste water to assess the NO₃⁻ removal capability of Ru-Co(OH)₂ electrocatalyst. As shown in Fig. 3f, following 7 h of sustained electrolysis at 100 mA cm⁻², the Ru-Co(OH)₂ electrocatalyst achieves 93.4% NO₃⁻ removal efficiency with residual NO₃⁻ concentrations below the World Health Organization (WHO) drinking water standard, demonstrating its application value for wastewater treatment. Besides, control experiments and ¹⁵N isotopic labeling studies were systematically conducted to confirm the origin of NH₃. As a result, control experiments detected negligible NH₃ background signals in electrolytes without adding NO₃⁻ throughout electrochemical testing, excluding non-catalytic ammonia contamination sources (Supplementary Fig. 21). Isotopic labeling experiments with ¹⁵NO₃⁻ and ¹⁴NO₃⁻ feedstocks yielded definitive ¹H nuclear magnetic resonance (NMR) signals of characteristic doublet peaks for ¹⁵NH₄⁺ and triplet peaks for ¹⁴NH₄⁺, respectively (Supplementary Figs. 22, 23a), confirming the origin of NH₃ from electrocatalytic NO₃⁻RR while excluding ambient nitrogen contamination. Furthermore, the produced NH₃ concentration measured by ¹H NMR shows a negligible difference with the result based on the indophenol blue method (Supplementary Fig. 23b), demonstrating the experimental consistency and reliability of this work. Besides, the consecutive cycling test shows the stable NO₃⁻RR performance of Ru-Co(OH)₂ electrocatalyst (Supplementary Fig. 24). TEM and XPS analyses of Ru-Co(OH)₂ electrocatalyst after cycling tests revealed negligible structural degradation and compositional change, illustrating the structural stability (Supplementary Figs. 25, 26). With the aim of evaluating the potential of the Ru-Co(OH)₂ electrocatalyst in practical applications, a MEA electrolyzer was further applied to investigate the long-term NO₃⁻RR stability (Fig. 3g). As shown in Fig. 3h, NH₃ yield rate, Faradaic efficiency, and cell voltage of the Ru-Co(OH)₂ electrocatalyst can keep relatively stable during the 1200 h of NO₃⁻RR tests at a large current density of 200 mA cm⁻², revealing the significant utility of Ru-Co(OH)₂ for upgrading industrial effluents abundant in nitrate. In order to evaluate the commercialization value of Ru-Co(OH)₂ catalyst in MEA electrolyzer, techno-economic analysis (TEA) was further performed (Supplementary Fig. 27, Supplementary Note 1). As a result, the NH₃ production cost of the Ru-Co(OH)₂ electrocatalyst is $ 979.92 per tonne, which is lower than the potential profits from selling NH₃ at the current market price and treating NO₃⁻-rich waste water³¹.

Reaction mechanism analysis

Electrocatalytic nitrate reduction involves the 8 electrons and 9-proton-coupled process, while the continuous proton supply is the guarantee of completed hydrogenation conversion from NO₃⁻ to NH₃³². In alkaline conditions, active hydrogen species (^*H) are derived from the H-OH bond cleavage of interfacial water, thus the rational modulation of interfacial water behavior with increased ^*H supply holds the key for boosting the hydrogenated process of nitrate electroreduction. For investigating the influence of hydroxide support on the water dissociation and ^*H behavior, in-situ electrochemical impedance spectroscopy (EIS) was conducted at different potentials. Bode phase analysis can reveal the electrochemical interface processes between electrode and electrolyte³³. In the absence of NO₃⁻, Ru-Co(OH)₂ catalyst exhibits a rapid decrease of the phase angle with simultaneous shift towards high frequency when potential is below +0.1 V vs. RHE, revealing the accelerated kinetics of Volmer step (H₂O + ^* + e⁻ → ^*H + OH⁻) through enhanced water dissociation (Supplementary Fig. 28). With the introduction of 0.1 M NO₃⁻, the phase angle shows attenuation when potential reaching +0.1 V vs. RHE, revealing the rapid consumption of ^*H by the dominant NO₃⁻RR process (Fig. 4a). Compared with the Ru-Ni(OH)₂ and Ru-Fe(OH)₂ counterparts, Ru-Co(OH)₂ electrocatalyst presents lower amplitude of phase angle in a wide potential range (Supplementary Figs. 29–31), implying the rapid reaction kinetics of NO₃⁻RR motivated by the accelerated H₂O dissociation. Moreover, the formation and consumption process of ^*H was explored by cyclic voltammetry (CV) test³⁴. In the absence of NO₃⁻, the Ru-Co(OH)₂ electrocatalyst displays an increased peak intensity of ^*H characteristic peak than that of Ru-Ni(OH)₂ and Ru-Fe(OH)₂ (Fig. 4b), indicating the stronger ^*H production ability of Co(OH)₂ support. Besides, Ru-Co(OH)₂ shows a reduced slope of the linear fitting between the CV scan rate and ^*H peak position (Supplementary Fig. 32), which can be ascribed to the improved ^*H production kinetics³⁵. With the introduction of NO₃⁻, the characteristic peak of ^*H totally disappears in the CV curve of Ru-Co(OH)₂ electrocatalyst; however, ^*H signals of Ru-Ni(OH)₂ and Ru-Fe(OH)₂ can be partly retained, revealing the rapid consumption of ^*H and accelerated NO₃⁻RR kinetics of Ru-Co(OH)₂. Moreover, the ^*H quenching experiment reveals dramatically decreased NO₃⁻RR performance after adding tertiary butanol (TBA) into the electrolyte (Supplementary Fig. 33), implying the essential involvement of ^*H during the hydrogenated process of nitrate electroreduction. Furthermore, the KSCN poisoning experiment shows a significant suppression of both current density and NH₃ production rate (Supplementary Fig. 34), attributed to the competitive adsorption of SCN⁻ ions on Ru active sites³⁶. This observation confirms the critical role of Ru clusters as the active site for electrochemical nitrate electroreduction. Above analysis confirms that Co(OH)₂ support can accelerate the water dissociation with increased ^*H provision, and modulated metal–support interactions endow Ru active sites with faster ^*H consumption kinetics for boosting the hydrogenation reaction process.

Fig. 4. Reaction mechanism analysis.

a In-situ EIS spectra of the Ru-Co(OH)₂ electrocatalyst. b CV curves of different electrocatalysts in 1 M KOH with and without 0.1 M NO₃⁻ for determining surface-adsorbed ^*H. c In-situ Raman contour maps of the Ru-Co(OH)₂ electrocatalyst during NO₃⁻RR performance measurements. d In-situ Raman spectra of the O–H stretching mode of interfacial H₂O on Ru-Co(OH)₂ surface. e In-situ ATR-FTIR spectra of Ru-Co(OH)₂ electrocatalyst. f In-situ DEMS measurements of Ru-Co(OH)₂ electrocatalyst during NO₃⁻RR.

In-situ Raman measurements were further conducted to track reaction intermediates and elucidate the interfacial water structure. As shown in Fig. 4c, the distinguishable Raman signals at 509 and 575–680 cm⁻¹ are assigned to A_2u and E_g mode of Co(OH)₂³⁷, and the steady Raman peaks at different potentials reveal the structural stability of Co(OH)₂ support. Additionally, the Raman peak at 1050 cm⁻¹ is attributed to the symmetric stretching of ^*NO₃, indicating the strong adsorption of ^*NO₃ on the electrocatalyst surface³⁸. Moreover, the wide O-H stretching peak of interfacial water observed within 2900–3750 cm⁻¹ is deconvoluted into three characteristic peaks (Fig. 4d), assigned to four hydrogen bonds coordinated water (4-HB·H₂O), two hydrogen bonds coordinated water (2-HB·H₂O), and K⁺ ion hydrated water (K·H₂O), respectively^39,40. In comparison with the robust hydrogen–bonding interactions of 4-HB·H₂O, the 2-HB·H₂O with fewer coordinated hydrogen bonds enable rapid dissociation process with facilitated proton transfer kinetics⁴¹. Compared to the Ru-Ni(OH)₂ and Ru-Fe(OH)₂, the Ru-Co(OH)₂ electrocatalyst displays the highest 2-HB·H₂O content and lowest 4-HB·H₂O content (Supplementary Figs. 35–37), indicating that the Co(OH)₂ support can weaken the hydrogen–bonding interactions between interfacial water and accelerate H₂O dissociation with elevated availability of ^*H species for boosting hydrogenated conversion of nitrate.

In-situ attenuated total reflection Fourier transformed infrared (ATR-FTIR) spectroscopy was applied to explore the electrocatalytic NO₃⁻RR mechanism by detecting reaction intermediates. As shown in Fig. 4e, a prominent H-O-H bending vibration of interfacial water is located at 1631 cm⁻¹, demonstrating the robust water adsorption on the Ru-Co(OH)₂ surface. Moreover, the upward absorption band at 1679 cm⁻¹ is ascribed to the N-O stretching vibration of ^*NO intermediate, and peaks located at 1339 and 1374 cm⁻¹ are the asymmetric stretching vibrations of ^*NO₂ and ^*NO₃⁴². Besides, the characteristic peaks at 1538 and 1508 cm⁻¹ correspond to ^*NH₂ and ^*NOH^43,44, and the N-H stretching vibration peak of ^*NH₃ is located at 1456 cm⁻¹. The detection of ^*NOH, ^*NH₂, and ^*NH₃ intermediates reflects the superior hydrogenation ability of the Ru-Co(OH)₂ electrocatalyst. Furthermore, the decreased signal intensity of ^*NO₃ and the increased signal intensity of ^*NH₃ with the negative shift of potential indicate improved transformation from NO₃⁻ to NH₃ over the Ru-Co(OH)₂ electrocatalyst. However, Ru-Ni(OH)₂ and Ru-Fe(OH)₂ electrocatalysts exhibit weak characteristic peaks for different intermediates in the in-situ ATR-FTIR spectra (Supplementary Figs. 38 and 39), corresponding to the inferior NO₃⁻RR performance. Furthermore, in-situ differential electrochemical mass spectrometry (DEMS) was performed for five consecutive cycles to systematically elucidate the NO₃⁻RR reaction pathway of Ru-Co(OH)₂. Different mass-to-charge ratio (m/z) signals of 30, 14, 15, 16, 17 were detected, assigning to ^*NO, ^*N, ^*NH, ^*NH₂ and ^*NH₃ intermediates, respectively (Fig. 4f). Importantly, the absence of ^*NO₂ signal (m/z = 46) confirms fast conversion of the ^*NO₂ intermediate during the NO₃⁻RR process. Moreover, the weak peak intensity of ^*NOH (m/z = 31) relative to other intermediates could be attributed to slow production and quick consumption of ^*NOH intermediate during the NO₃⁻RR process. With the mechanism substantiated by above in-situ characterizations, the reaction pathway is proposed as “^*NO₃ → ^*NO₂ → ^*NO→ ^*NOH→ ^*N→ ^*NH→ ^*NH₂ → ^*NH₃”.

Metal–support interaction investigation

The modulation mechanism of metal–support interaction on the NO₃⁻RR performance was explored by density functional theory (DFT) calculations. Ru-Ni(OH)₂ shows a more negative binding energy than that of Ru-Co(OH)₂ and Ru-Fe(OH)₂ (Supplementary Fig. 40). Charge density difference and corresponding planar-average charge density plot show that more electrons are transferred from Ru clusters to the Ni(OH)₂ support (Fig. 5a), indicating the strong metal–support interaction of Ru-Ni(OH)₂. Medium electron transfer occurs between Ru clusters and Co(OH)₂ support (Fig. 5b), revealing the moderate metal–support interaction of Ru- Co(OH)₂. The least electrons are transferred across the Ru-Fe(OH)₂ interface (Fig. 5c), corresponding to the weak metal–support interaction. The Bader charge analysis demonstrates that Ru clusters of Ru-Ni(OH)₂ show the highest valence state because more electrons are transferred to the Ni(OH)₂ support (Supplementary Fig. 41). In contrast, the valence electrons of Ru clusters are more abundant for the Ru-Fe(OH)₂ due to the weak metal–support interaction. Moreover, the Ru-Fe(OH)₂ with the smallest work function (Φ) of 3.52 eV can easily donate electrons for improving the NO₃⁻ adsorption (Supplementary Fig. 42). As a result of the different metal–support interactions, Ru-Fe(OH)₂ exhibits the strongest NO₃⁻ adsorption while the Ru-Ni(OH)₂ exhibits the weakest NO₃⁻ adsorption (Supplementary Figs. 43, 44, 45). Based on the Sabatier principle, the reactant adsorption can be neither too strong nor too weak. Accordingly, Ru-Co(OH)₂ with the moderate NO₃⁻ adsorption is more promising for achieving the efficient NO₃⁻RR process.

Fig. 5. Metal–support interaction investigation.

Charge density difference and corresponding planar-average charge density plot of a Ru-Ni(OH)₂, b Ru-Co(OH)₂, and c Ru-Fe(OH)₂. The cyan and green regions represent electron depletion and accumulation, respectively. The isosurface of charge density is set to 0.005 e/Bohr³. d Free energy diagrams of NO₃⁻RR on different electrocatalysts when U_RHE = 0 V. The insets show the adsorption structures of each intermediate on the Ru-Co(OH)₂ electrocatalyst. e Free energy diagrams of ^*H formation and HER process on different electrocatalysts when U_RHE = 0 V. f Schematic illustration of modulated metal–support interaction for improving NO₃⁻RR performance. Ru, Ni, Co, Fe, N, O, and H atoms are colored in yellow, gray, purple, green, blue, red, and white, respectively.

Based on the constructed reaction pathway by in-situ characterizations, the Gibbs free energy of adsorbed intermediates were calculated as shown in Fig. 5d. Both Ru-Co(OH)₂ and Ru-Ni(OH)₂ electrocatalysts shows the highest energy barrier of the ^*NO → ^*NOH among the whole elementary steps, confirming that the ^*NO → ^*NOH step is the rate-determining step. Ru-Co(OH)₂ electrocatalyst with moderate metal–support interaction displays a relatively lower energy barrier (0.69 eV) of the RDS. The free energy change of the RDS is elevated to 1.12 eV for Ru-Ni(OH)₂ due to the weak adsorption of ^*NOH (Supplementary Fig. 46). Different from Ru-Co(OH)₂ and Ru-Ni(OH)₂, the RDS of Ru-Fe(OH)₂ is changed into ^*NH₂ → ^*NH₃ with a high energy barrier 3.07 eV (Supplementary Fig. 47), due to the too strong adsorption of the intermediate and impeded hydrogenation process. Considering that H₂O serves as the source of ^*H involved in the hydrogenation process, the adsorption and dissociation process of H₂O can significantly influence the NO₃⁻RR performance. It is worth noting that Ru clusters serve as active sites for NO₃⁻ adsorption, and hydroxide support contributes to H₂O adsorption (Supplementary Fig. 48). Among different hydroxide supports, Co(OH)₂ and Ni(OH)₂ are more beneficial to the H₂O adsorption compared with Fe(OH)₂ (Supplementary Fig. 49), due to the up-shifted 3d-states of Co and Ni closer to Fermi level (Supplementary Fig. 50). In addition, the projected crystal orbital Hamilton population (pCOHP) results reveal the weak interaction between adsorbed H₂O molecule and Fe(OH)₂ surface, due to the increased integrated COHP value (−0.17 eV) with less bonding state occupancy (Supplementary Fig. 51). Furthermore, compared with the Fe(OH)₂, the free energy changes of ^*H₂O dissociation are much reduced for Co(OH)₂ and Ni(OH)₂ for boosting ^*H supply (Fig. 5e). Therefore, the modulation mechanism of metal–support interaction on the NO₃⁻RR performance can be described in Fig. 5f. Ru-Fe(OH)₂ with weak metal–support interaction shows too strong NO₃⁻ adsorption and limited ^*H supply, restricting the hydrogenation process for ammonia production. However, too strong metal–support interaction can weaken the NO₃⁻ adsorption on Ru-Ni(OH)₂, thus lowering its catalytic activity for nitrate reduction. Compared to Ru-Fe(OH)₂ and Ru-Ni(OH)₂, the moderate metal–support interaction can endow Ru-Co(OH)₂ with proper NO₃⁻ adsorption and ^*H provision for promoting the hydrogenation conversion from nitrate to ammonia during NO₃⁻RR process.

Rechargeable hybrid battery measurements

In addition to upgrading nitrate-rich wastewater, electrocatalytic NO₃⁻RR also shows potential in energy conversion. For example, Zn-NO₃⁻ battery coupling cathodic NO₃⁻RR and anodic Zn oxidation has spurred recent interest due to the simultaneous NH₃ production and energy output. However, the present discharge potential and power density are unsatisfactory because of the high overpotential of present NO₃⁻RR electrocatalyst⁴⁵. Moreover, the charge process of Zn-NO₃⁻ battery is usually impeded by inappropriate oxidation reaction at the nitrate-containing electrolyte side, because oxygen evolution reaction (OER) or ammonia oxidation reaction tends to elevate charge plateaus or deplete value-added ammonia. In our previous work, the introduction of alcohols into the catholyte has been proved as an effective strategy for sustaining long-term charge-discharge stability³¹. Furthermore, the added alcohols can be extended to ethylene glycol (EG), which is a main hydrolysis product of depolymerized polyethylene terephthalate (PET) plastic. Critically, the replacement of OER by EG oxidation in the cathode side can not only promote the charge process, but also upgrade PET plastic. The electrochemical oxidation of EG in PET hydrolysate exhibits substantially lower overpotential than OER, enabling reduced charge plateaus (Supplementary Fig. 52)⁴⁶. Therefore, the rechargeable Zn-NO₃⁻/EG hybrid battery was designed to simultaneously achieve nitrate upgrading, PET plastic upcycling, and energy conversion. As shown in Fig. 6a, cathodic NO₃⁻RR occurs concurrently with the oxidation of anodic Zn into Zn(OH)₄²⁻ during the discharge process. During the charge process (Fig. 6b), EG is oxidized into acetic acid, which can be coupled with NH₃ to generate value-added ammonium acetate.

Fig. 6. Rechargeable hybrid battery measurements.

Schematic diagram of a discharge process and b charge process of Zn-NO₃⁻/EG hybrid battery. c Polarization curve and discharge power density tests of different cathodes. d Multi-step chronoamperometric curves of different cathodes at various current densities during the discharge process. e Schematic diagram of PET upcycling. f High-rate charge-discharge cycling test of Zn-NO₃⁻/EG hybrid battery using the Ru-Co(OH)₂ cathode at 10 mA cm⁻². g ¹H NMR spectra of catholyte before and after the charge-discharge cycling test.

Based on the superior performance of nitrate electroreduction of Ru-Co(OH)₂ electrocatalyst, a Zn-NO₃⁻ battery was firstly assembled with the self-supported Ru-Co(OH)₂ electrocatalyst as the cathode and Zn metal as the anode. The open-circuit voltage (OCV) of Zn-NO₃⁻ battery assembled with Ru-Co(OH)₂ electrocatalyst is 1.39 V vs. Zn²⁺/Zn, higher than that of the Ru-Ni(OH)₂ and Ru-Fe(OH)₂ based Zn-NO₃⁻ battery, revealing the reduced NO₃⁻RR overpotentials of Ru-Co(OH)₂ cathode (Supplementary Fig. 53). Significantly, the Zn-NO₃⁻ battery using Ru-Co(OH)₂ electrocatalyst exhibits a peak power density of 46.47 mW cm⁻² at a large current density of 143.88 mA cm⁻² (Fig. 6c). Moreover, the Ru-Co(OH)₂ cathode demonstrates exceptional rate performance with higher discharge voltage than Ru-Ni(OH)₂ and Ru-Fe(OH)₂ across a wide current density range of 10–160 mA cm⁻² (Fig. 6d). It is noteworthy that the power density and corresponding current density of Ru-Co(OH)₂ cathode in this work outperform the majority of recently reported cathodes in Zn-NO₃⁻ battery in alkaline environment (Supplementary Table 3). With the aim of lowering charge plateaus and turning waste into treasure, the EG originated from PET plastic is added into the catholyte to construct the Zn-NO₃⁻/EG hybrid battery. As shown in Fig. 6e, the upcycling of PET plastic can be achieved by coupling hydrolysis and electrolysis processes, in which EG can be oxidized into acetic acid by the charge process of Zn-NO₃⁻/EG hybrid battery. As a result, the Zn-NO₃⁻/EG battery with Ru-Ni(OH)₂ cathode maintains a high-rate charge-discharge stability over 150 cycles at a large current density of 10 mA cm⁻² (Fig. 6f). After charge-discharge cycling test, ¹H NMR characterization of the catholyte shows obvious signals of NH₄⁺ and acetic acid and minor signal of formic acid, confirming the main product of ammonium acetate (Fig. 6g). Based on above results, the constructed Zn-NO₃⁻/EG battery shows a great potential in waste upcycling and energy conversion by upgrading nitrate and PET plastic into high-valued products.

Discussion

In summary, the hydroxide-supported metallic Ru electrocatalyst was synthesized by a universal self-corrosion strategy. The well-designed Ru-Co(OH)₂ electrocatalyst exhibits a high NH₃ Faradaic efficiency of ~100% at a positive potential of +0.1 V vs. RHE and a high energy efficiency of 49.5% at +0.2 V vs. RHE. Furthermore, Ru-Co(OH)₂ electrocatalyst can keep relatively stable during the 1200 h of NO₃⁻RR stability tests at a large current density of 200 mA cm⁻² in a MEA electrolyzer. The TEA assessment illustrates the industrial-scale electrochemical NO₃⁻RR feasibility due to the high performance at positive potentials, high stability, and high commercialization value. In-situ EIS and in-situ Raman spectra show that Co(OH)₂ support can weaken the hydrogen–bonding interactions between interfacial water and accelerate H₂O dissociation for boosting the hydrogenation process of nitrate. In-situ ATR-FTIR spectra and in-situ DEMS elucidate the reaction intermediates and the reaction pathway. Theoretical calculations reveal that compared with other hydroxides, Co(OH)₂ shows moderate metal–support interaction with metallic Ru, which endows metallic Ru appropriate NO₃⁻ adsorption ability and accelerates the following hydrogenation steps. In order to simultaneously achieve waste nitrate upgrading, PET plastic upcycling, and energy conversion, a rechargeable Zn-NO₃⁻/EG hybrid battery system is designed and assembled by utilizing Ru-Co(OH)₂ cathode. The constructed hybrid battery system exhibits a high-rate charge-discharge stability over 150 cycles at a large current density of 10 mA cm⁻². Therefore, this work presents a universal modulation strategy of metal–support interaction for facilitating NO₃⁻RR at positive potentials, which can further inspire the rational design of supported metal catalysts in extensive hydrogenation reactions.

Methods

Chemicals

Ruthenium chloride hydrate (Aladdin, RuCl₃ ∙ xH₂O, 99.95%), sodium chloride (Aladdin, NaCl, AR), ethanedioic acid dihydrate (Sinopharm Chemical Reagent Co., Ltd, C₂H₂O₄ ∙ 2H₂O, AR), ammonium chloride (Aladdin, NH₄Cl, AR), ammonium chloride-15N (Aladdin, ¹⁵NH₄Cl, 99%), potassium nitrate (Sinopharm Chemical Reagent Co., Ltd, KNO₃, AR), potassium nitrite (Aladdin, KNO₂, 97%), potassium hydroxide (Aladdin, KOH, 85%), sodium hydroxide (Aladdin, NaOH, 99%), salicylic acid (Sinopharm Chemical Reagent Co., Ltd, C₇H₆O₃, AR), sodium citrate (Sinopharm Chemical Reagent Co., Ltd, C₆H₅Na₃O₇, 98%), sodium hypochlorite aqueous solution (Sinopharm Chemical Reagent Co., Ltd, NaClO, CP), sodium nitroferricyanide dihydrate (Aladdin, C₅FeN₆Na₂O ∙ 2H₂O, AR), N-(1-Naphthyl)ethylenediamine dihydrochloride (Sinopharm Chemical Reagent Co., Ltd, C₁₂H₁₄N₂ ∙ 2HCl, AR), sulfanilamide (Aladdin, C₆H₈N₂O₂S, AR), sulfamic acid (Aladdin, H₃NO₃S, AR), phosphoric acid (Aladdin, H₃PO₄, 85%), ethanol absolute (Sinopharm Chemical Reagent Co., Ltd, C₂H₆O, AR), hydrazine hydrate (Sinopharm Chemical Reagent Co., Ltd, N₂H₄ ∙ H₂O, 85%), sulfuric acid (Sinopharm Chemical Reagent Co., Ltd, H₂SO₄, AR), hydrochloric acid (Sinopharm Chemical Reagent Co., Ltd, HCl, AR), cis-Butenedioic acid (Sinopharm Chemical Reagent Co., Ltd, C₄H₄O₄, AR), tert-Butanol (Macklin C₄H₁₀O, AR), potassium thiocyanate Analytical Titrant (Aladdin, KSCN, AR), zinc acetate (Aladdin, C₄H₆O₄Zn ∙ H₂O, AR), sodium deuteroxide (Aladdin, NaOD/D₂O, AR), p-Dimethylaminobenzaldehyde (Aladdin, C₉H₁₁NO, AR). Ultrapure water used throughout all experiments was purified through a Millipore system.

Preparation of Ru-Co(OH)₂ electrocatalyst

Ru-Co(OH)₂ electrocatalyst was prepared by the self-corrosion method with the assistance of Ru³⁺ and O₂. In details, a piece of Co foam (1 × 2 cm²) was firstly sonicated in HCl solution (25 vol.%) for 15 min to remove surface oxides, and then washed by ultrapure water. After the pretreatment, the clean Co foam was immediately immersed in 10 mM RuCl₃ solution with continuous agitation at 60 °C for 15 min. During the self-corrosion process, metallic Co was involved in the spontaneous redox reaction with RuCl₃ and O₂, resulting in the formation of Ru clusters loaded on Co(OH)₂ support. As-prepared Ru-Co(OH)₂ was washed by ultrapure water to remove the residual RuCl₃, and then dried in a vacuum oven at 60 °C for 24 h.

Preparation of control samples

Ru-Ni(OH)₂ and Ru-Fe(OH)₂ electrocatalysts were prepared by a similar approach by replacing Co foam with Ni foam and Fe foam, respectively. Ru electrocatalyst supported on Co foam was synthesized by a similar approach under Ar atmosphere. Co(OH)₂ electrocatalyst supported on Co foam was prepared by a similar approach by replacing RuCl₃ solution with NaCl solution.

Characterizations

X-ray diffraction (XRD) patterns were acquired using a Bruker D8 Advance diffractometer with Cu Kα radiation (40 kV, 40 mA). Scanning electron microscopy (SEM) images were captured on a Thermo Scientific Verios G4 instrument equipped with an energy-dispersive spectroscopy (EDS) detector. Transmission electron microscopy (TEM) characterization was performed on an FEI Tecnai F20 microscope (200 kV accelerating voltage), with elemental mapping conducted via EDS. X-ray photoelectron spectroscopy (XPS) measurements were carried out on a Thermo Scientific K-Alpha system, and all spectra were charge-corrected by referencing the C 1 s peak to 284.8 eV. Raman data were collected on the Witec Alpha 300 R Raman system using the excitation wavelength of 532 nm. UV–visible (UV–vis) absorption spectra were recorded on a UV7600 spectrophotometer.

Electrochemical nitrate reduction tests

All electrochemical measurements were performed using a CHI760E electrochemical workstation. Electrocatalytic nitrate reduction experiments were carried out on an H-type cell equipped with an anion exchange membrane (FAB-PK-130, thickness of 130 μm, size of 1.5 × 1.5 cm²), employing a standard three-electrode configuration. The electrolyte was prepared by dissolving KNO₃ into 1 M KOH with the KNO₃ concentration of 0.1 M (pH 14). In order to investigate the influence of NO₃⁻ concentration on the NO₃⁻RR performance, an electrolyte containing different KNO₃ concentrations (0.01~ 0.5 M) was prepared. The as-prepared self-supported electrocatalysts (1 × 0.5 cm²) were directly utilized as the working electrode, the Hg/HgO electrode and Pt foil (1 × 1 cm²) were used as the reference electrode and counter electrode. All potentials were converted to RHE based on the Nernst equation (E_RHE= E_Hg/HgO + 0.059 × pH + 0.098). Before electrochemical tests, the electrolyte was saturated by Ar to remove the dissolved O₂. LSV tests were conducted from +0.6 V to −0.6 V vs. RHE at a scan rate of 5 mV s⁻¹. For electrocatalytic nitrate reduction experiments, chronoamperometry tests were performed for 1 h, and then the electrolyte was collected to analyze the liquid products.

In-situ Raman spectroscopy measurements

In-situ Raman spectroscopy measurements were performed using a WITec Alpha 300 R system with a 532 nm excitation wavelength. Spectra were acquired at 2 s exposure time averaged over 10 accumulations (Supplementary Fig. 54). Chronoamperometry tests were applied across a potential range of +0.3 to −0.3 V vs. RHE in a custom-designed Teflon electrochemical cell (EC-Raman-H, Beijing Scistar Technology Co., Ltd) featuring a quartz optical window, with continuous electrolyte circulation maintained via a peristaltic pump. The self-supported electrocatalyst was immersed in the flowing electrolyte while the electrode plane was kept perpendicular to the laser. Raman spectra were recorded after 1-min of the chronoamperometry test.

In-situ ATR-FTIR measurements

In-situ ATR-FTIR measurements were conducted on the Thermo Scientific Nicolet iS50 spectrometer equipped with a liquid N₂-cooled MCT-A detector (Supplementary Fig. 55). Chronoamperometry tests were applied across a potential range of +0.3 to −0.3 V vs. RHE in a custom-designed Teflon electrochemical cell (EC-ATR-H, Beijing Scistar Technology Co., Ltd) featuring a quartz optical window, with continuous electrolyte circulation maintained via a peristaltic pump. The self-supported electrocatalyst was press against the silicon crystal covered by a gold film, establishing a nanometer-scale electrolyte interface between the working electrode and substrate. FTIR spectra at each applied potential were recorded after 1-min of chronoamperometry test, and the background spectrum was taken at the open-circuit potential. The spectra were obtained from an average of 32 scans with a resolution of 4 cm⁻¹.

In-situ DEMS measurements

In-situ DEMS measurements were performed on PM-DEMS with a customized electrochemical cell filled with flowing 1 M KOH + 0.5 M KNO₃ electrolyte (Supplementary Fig. 56). The catalyst loaded on a gold-coated hydrophobic polytetrafluoroethylene (PTFE) membrane serves as the working electrode. The Hg/HgO electrode and the Pt wire were used as the reference electrode and the counter electrode, respectively. LSV tests were performed from +0.4 V to −0.3 V vs. RHE, and the mass signals were recorded during the LSV tests, five consecutive tests were conducted to avoid experimental errors. During the electrochemical tests, the reaction occurred at the catalyst/electrolyte interface on one side of the PTFE membrane. The generated volatile intermediate molecules/species penetrated crossover the PTFE membrane into the vacuum chamber, which were then recorded by the MS detector.

Zn-NO₃⁻ battery measurements

A Zn-NO₃⁻ battery was assembled with the self-supported Ru-Co(OH)₂ electrocatalyst as the cathode and Zn metal as the anode. 1 M KOH with 0.5 M KNO₃ was used as catholyte and 1 M KOH was used as anolyte during the discharge measurements. Polarization curves from OCV to 0 V (vs. Zn²⁺/Zn) were acquired at 5 mV s⁻¹ using a CHI760E workstation. Chronoamperometry discharge tests were performed at varied current densities from 10 to 160 mA cm⁻².

PET hydrolysate was prepared from the discarded plastic bottles. An amount of 8 g PET plastic bottles were cut into pieces and then added to 200 mL 4 M KOH solution. The mixed solution was heated in an oil bath for 96 h at 90 °C. Finally, the solution was diluted to a PET hydrolysate containing 1 M KOH. For the charge-discharge cycling test of Zn-NO₃⁻/EG hybrid battery, catholytes contain 0.5 M KNO₃ and PET hydrolysate, and anolytes contain 1 M KOH and 0.02 M Zn(CH₃COO)₂. Long-term cycling on LAND systems was coupled with post-test catholyte analysis via ¹H NMR spectroscopy.

Ammonia quantification

The ammonia concentration was quantified using the indophenol blue method. Specifically, 2 mL of diluted electrolyte was combined with 2 mL of NaOH solution (1 M, containing 5 wt.% salicylic acid and 5 wt.% sodium citrate), 1 mL of NaClO (0.05 M), and 0.2 mL of C₅FeN₆Na₂O aqueous solution (1 wt.%) to prepare the reaction mixture. After allowing the mixture to stand in darkness for 2 h, the absorption spectrum was recorded using a UV–vis spectrophotometer. The indophenol blue formation was quantified by measuring the absorbance at 655 nm. Calibration curves were established using standard ammonia solutions with known concentrations. The standard curve (y = 0.072 + 0.474x, R² = 0.999) demonstrated excellent linearity between absorbance and NH₃ concentration.

Determination of hydrazine

The hydrazine (N₂H₄) concentration was determined via the Watt and Chrisp method. The chromogenic reagent was prepared by dissolving 0.998 g of p-Dimethylaminobenzaldehyde in a mixture of 5 mL concentrated HCl and 50 mL ethanol. Subsequently, 3 mL of the prepared reagent was mixed with 3 mL of electrolyte and allowed to stand in darkness for 10 min. The absorbance of the resulting solution was measured at 460 nm. Calibration was performed using standard hydrazine hydrate solutions of known concentrations, yielding a linear standard curve ((y = 0.069 + 1.133x, R² = 0.999) that demonstrated excellent correlation between absorbance and N₂H₄ concentration.

Determination of NO₃⁻

In a standard procedure, 5.0 mL of either standard or diluted sample solutions were combined with 0.10 mL of 1.0 M HCl and 0.01 mL of 0.8 wt.% sulfamic acid solution. Following thorough mixing and standing for 5 min, the NO₃⁻ concentration was quantified via UV–vis spectrophotometry across the 200–300 nm wavelength range. The NO₃⁻ concentration was determined based on absorbance measurements at 220 and 275 nm. Calibration was performed using KNO₃ standard solutions of known concentrations, yielding a linear standard curve (y = −0.002 + 0.066x, R² = 0.999) that exhibited excellent correlation between the calculated absorbance (A = A_220nm– 2A_275nm) and NO₃⁻ concentration.

Determination of NO₂⁻

The NO₂⁻ concentration was determined through the Griess reaction. The chromogenic reagent was prepared by dissolving 0.2 g of N-(1-Naphthyl) ethylenediamine dihydrochloride and 4 g of sulfanilamide in a mixture containing 10 mL of phosphoric acid and 90 mL of ultrapure water. Subsequently, 0.10 mL of the prepared reagent was added to 5 mL of standard or sample solutions, followed by a 20-min incubation period. NO₂⁻ quantification was performed via UV–vis spectrophotometry across the 450–650 nm wavelength range, with specific absorbance measurements taken at 540 nm. Calibration was established using KNO₂ standard solutions of known concentrations, yielding a linear standard curve (y = 0.049 + 0.194x, R² = 0.999) that demonstrated excellent correlation between absorbance and NO₂⁻ concentration.

¹⁵N isotopic labeling measurement

The ¹⁵N isotopic labeling experiment was performed using K¹⁵NO₃ (Aladdin, 99 atom%) as the ¹⁵N source. Following electrochemical nitrate reduction at +0.1 V vs. RHE for 1 h in 1 M KOH electrolyte containing 0.1 M ¹⁵NO₃⁻, the resulting ¹⁵NH₃ was quantified by ¹H nuclear magnetic resonance spectroscopy (¹H NMR, Bruker 600 MHz). Specifically, the pH of the diluted electrolyte was first adjusted to 2–3, after which 425 μL of electrolyte was combined with 50 μL of DMSO-d₆ and 25 μL of maleic acid (5000 ppm) as an internal standard for NMR analysis.

Calculations of NH₃ yield rate and Faradaic efficiency

The Faradaic efficiency (η) was determined using the following equation:

$$\eta =\frac{n\cdot F\cdot C\cdot V}{M\cdot Q}$$ 1

where n is the number of electrons required for producing one NH₃ molecule (n = 8), F is the Faraday constant (F = 96485.33), C is the measured NH₃ concentration, V is the volume of the electrolyte, M is the relative molecular mass of NH₃ (M = 17), and the Q is the total charge passed through the electrodes.

The NH₃ yield rate (R) was determined using the following equation:

$$R=\frac{C\cdot V}{t\cdot S}$$ 2

where C is the measured NH₃ concentration, V is the volume of the electrolyte, t is the reaction time, and S is the geometric area of the electrocatalyst.

The NH₃ partial current density (j) was determined using the following equation:

$$j=\frac{Q\cdot FE}{t\cdot m}$$ 3

where Q is the total charge passed through the electrodes, FE is the Faradaic efficiency, t is the reaction time, and m is the geometrical area of the working electrode.

The energy efficiency (EE) was defined as the ratio of fuel energy to applied electrical, which was determined using the following equation:

$$E{E}_{N{H}_3}=\frac{\left(1.23-E_{N{H}_3}^{\theta }\right)\cdot F{E}_{N{H}_3}}{1.23-E}$$ 4

where E^Ɵ_NH3 represents the equilibrium potential of electrochemical nitrate reduction to ammonia (0.69 V vs. RHE), FE_NH3 is the Faradaic efficiency of ammonia, and E is the applied potential vs. RHE without iR correction.

Theoretical calculations

All DFT calculations in this work were performed by using the Vienna Ab initio Simulation Package, combining generalized gradient approximation with the Perdew–Burke–Ernzerhof pseudopotentials. The cut-off energy of the plane wave function was 450 eV. All spin-polarized computational models were constructed in a box with dimensions of 13 × 13 × 19 Å, which included a 15 Å vacuum layer to eliminate the effects of Z-axis periodicity. All models have been fully relaxed with the energy convergence criterion of 10⁻⁵ eV and the force convergence criterion of 0.05 eV/Å, respectively. The K-point mesh precision used in the Brillouin zone is 0.04 × 2π/Å. DFT-D3 method has been used to correct the van der Waals interaction. The binding energy (E_b) between the substrate and the Ru cluster was calculated using the following equation:

$$E_{\mathrm{b}}={\mathrm{E}}_{\mathrm{Ru}/\mathrm{S}}-{\mathrm{E}}_{\mathrm{Ru}}-{\mathrm{E}}_{\mathrm{S}}$$

where E_Ru/S, E_Ru, and E_S are the total energies of the Ru cluster anchored on the substrate, the Ru cluster, and the substrate, respectively.

The free energy change of all reactions was calculated based on the computational hydrogen electrode (CHE) model. The hydrogen evolution reaction and its standard electrode potential are shown below:

$${\mathrm{H}}^{+}+{\mathrm{e}}^{-}\to 1/2{\mathrm{H}}_2$$

The ΔG for this reaction at a given pH and electrode potential U_SHE is given by:

$$\mathrm{\Delta }\mathrm{G}=1/2\mathrm{G}\mathrm{gas}\left({\mathrm{H}}_2\right)-\mathrm{G}\left({\mathrm{H}}^{+}\right)-\mathrm{G}\left({\mathrm{e}}^{-}\right)=0.059\times \mathrm{pH}+{\mathrm{e}\mathrm{U}}_{\mathrm{SHE}}$$

$$\mathrm{G}\left({\mathrm{H}}^{+}\right)+\mathrm{G}\left({\mathrm{e}}^{-}\right)=1/2{\mathrm{G}}_{\mathrm{gas}}\left({\mathrm{H}}_2\right)-0.059\times \mathrm{pH}-{\mathrm{e}\mathrm{U}}_{\mathrm{SHE}}$$

The chemical potential of the hydrogen–electron pair can be expressed at the RHE scale:

$$\mathrm{G}\left({\mathrm{H}}^{+}\right)+\mathrm{G}\left({\mathrm{e}}^{-}\right)=1/2{\mathrm{G}}_{\mathrm{gas}}\left({\mathrm{H}}_2\right)-{\mathrm{eU}}_{\mathrm{RHE}}$$

The Gibbs free energy (ΔG) is calculated as follows:

$$\mathrm{\Delta }\mathrm{G}=\mathrm{\Delta }\mathrm{E}+\mathrm{\Delta }{\mathrm{E}}_{ZPE}-\mathrm{T}\mathrm{\Delta }\mathrm{S}$$

where ΔE represents the total energy obtained. T is temperature (298.15 K). ΔE_ZPE and ΔS are the correction of zero-point energy and entropy, respectively.

The process of NO₃⁻ aqueous phase adsorption onto the electrode surface can be described as follows:

$${\mathrm{NO}}_3^{-}\left(\mathrm{l}\right)+*+\left({\mathrm{H}}^{+}\right)\to *{\mathrm{NO}}_3+{\mathrm{e}}^{-}+\left({\mathrm{H}}^{+}\right)\left(\mathrm{\Delta }\mathrm{G}\left({\mathrm{NO}}_3^{-}\right)\right)$$

The overall free energy of NO₃⁻ aqueous phase adsorption onto the electrocatalyst surface (ΔG(NO₃⁻)) can be calculated by the thermodynamic cycle as shown in Supplementary Fig. 43, and three steps of thermodynamic cycle are described as follows:

$${\mathrm{NO}}_3^{-}\left(\mathrm{l}\right)+*+\left({\mathrm{H}}^{+}\right)\to {\mathrm{H}\mathrm{NO}}_3\left(\mathrm{l}\right)+*\left(\mathrm{\Delta }{\mathrm{G}}_{\mathrm{protonation}}=0.317\mathrm{eV}\right)$$

$${\mathrm{HNO}}_3\left(\mathrm{l}\right)+*\to {\mathrm{HNO}}_3\left(\mathrm{g}\right)+*\left({\mathrm{\Delta }\mathrm{G}}_{\mathrm{vaporization}}=0.074\mathrm{eV}\right)$$

$${\mathrm{H}\mathrm{NO}}_3\left(\mathrm{g}\right)+*\to *{\mathrm{NO}}_3+{\mathrm{e}}^{-}+\left({\mathrm{H}}^{+}\right)\left({\mathrm{\Delta }\mathrm{G}}_{\mathrm{ads}}\left({*\mathrm{NO}}_3\right)\right)$$

Based on above thermodynamic cycle, the ΔG(NO₃⁻) can be calculated as follows:

$\mathrm{\Delta }\mathrm{G}\left({\mathrm{NO}}_3^{-}\right)={\mathrm{\Delta }\mathrm{G}}_{\mathrm{protonation}}+{\mathrm{\Delta }\mathrm{G}}_{\mathrm{vaporization}}+{\mathrm{\Delta }\mathrm{G}}_{\mathrm{ads}}\left(*{\mathrm{NO}}_3\right)=0.391\mathrm{eV}+\mathrm{G}\left(*{\mathrm{NO}}_3\right)+\mathrm{G}\left({\mathrm{H}}^{+}\right)+\mathrm{G}\left({\mathrm{e}}^{-}\right)-\mathrm{G}\left(*\right)-{\mathrm{G}}_{\mathrm{gas}}\left({\mathrm{HNO}}_3\right)$

And the ΔG(NO₃⁻) can be referenced by RHE based on the hydrogen-electron pair:

$$\mathrm{\Delta }\mathrm{G}\left({\mathrm{NO}}_3^{-}\right)=0.391\mathrm{eV}+\mathrm{G}\left({}^{*}\mathrm{N}{\mathrm{O}}_3\right)+1/2{\mathrm{G}}_{\mathrm{gas}}\left({\mathrm{H}}_2\right)-{\mathrm{eU}}_{\mathrm{RHE}}-\mathrm{G}\left(*\right){\mathrm{G}}_{\mathrm{gas}}\left({\mathrm{HNO}}_3\right)$$

In spite of that, the DFT calculations are limited by the difficulties in modeling the exact material under real catalytic conditions.

Author contributions

Y.W., R.L., J.Z., and W.Y. conceived and supervised the project. Y.T. performed the experiments and collected the data. Y.W. conducted the theoretical calculations. Y.T. and Y.W. co-wrote the manuscript. All authors edited the paper.

Peer review

Peer review information

Nature Communications thanks the anonymous reviewers for their contribution to the peer review of this work. A peer review file is available.

Data availability

The data supporting the conclusions of this study are presented in the paper and its supplementary information. Source data are provided with this paper.

Competing interests

The authors declare no competing interests.

Supplementary information

The online version contains supplementary material available at 10.1038/s41467-026-69802-5.