METHODS OF PREPARING A CATALYST AND AN ELECTROLYTIC CELL FOR USE IN SYNTHESISING AMMONIA, AND METHODS OF SYNTHESISING AMMONIA

Abstract

One aspect of the present invention is concerned with a method of preparing a catalyst for use in synthesizing ammonia from nitrate. The method has the steps of i) providing a nickel foam material, ii) removing an oxide layer on the nickel foam material, iii) subjecting the nickel foam material to cooper sulphate (Cu2SO4) solution for a predetermined period of time for forming a copper (Cu) dopped nickel foam electrode, and iv) subjecting the Cu dopped nickel foam electrode to an annealing treatment at a predetermined temperature for a predetermined period of annealing time in atmospheric air thus forming a catalyst of copper oxide-nickel oxide nickel foam (CuO/NiO@NF) with a heterostructure. Other aspects of the invention are concerned with a method of preparing an electrolytic cell for use in synthesizing of ammonia, and a method of synthesis of ammonia from nitrate.

Description

TECHNICAL FIELD

The present invention is concerned with a method of preparing a catalyst for use in synthesizing ammonia, a method of preparing an electrolytic cell for use in synthesizing ammonia, and a method of synthesizing ammonia, for example or particularly from a source of nitrate.

BACKGROUND OF THE INVENTION

Ammonia (NH₃), a versatile raw ingredient (industrial chemical), can be used as a feedstock to synthesize fertilizers, dyes, pharmaceuticals, etc. Currently, the main synthetic route for NH₃ is the Haber-Bosch process, which requires harsh reaction conditions and ultra-high energy consumption and releases large amounts of CO₂. Meanwhile, large-scale industrial nitrogen fixation also leads to the most extensive disturbance of the nitrogen cycle. Nowadays, electrochemical nitrogen (N2) reduction reaction (NRR) has been widely considered a sustainable approach for preparing green ammonia via atmospheric nitrogen and water under mild conditions. However, the large bond energy of the N≡N triple bond (941 kJ mol⁻¹), limited N2 solubility and competing hydrogen evolution reaction lead to the low reaction rate, selectivity, and Faradic efficiency (FE) of NRR. In contrast, nitrate-rich wastewater as an attractive nitrogen source for ammonia electrosynthesis (nitrate reduction reaction, NO₃—RR) is energy-efficient (204 kJ mol⁻¹) and contributes to restoring the global imbalance nitrogen cycle.

Since NO₃⁻RR (NO₃⁻+6H₂O+8e⁻→NH₃+9OH⁻, E°=0.69 V vs. RHE) involves eight electrons transfer coupled with nine OH⁻ generation processes and thus suffers from various byproducts (i.e., NO₂—, N₂). To this end, electrocatalysts with higher efficiency and selectivity are required to utilize nitrate to produce value-added NH₃ selectively. In recent years, copper-based catalysts have been intensively investigated due to their favorable nitrate binding ability and poor H intermediate binding ability. However, the fabricated copper-based catalyst is prone to suffer from complex surface reconstruction processes during the reduction environment in the forms of valence variations, anion/cation defects generations, phase transformation/separation, etc. In particular, recent works demonstrated Cu-based catalysts experienced the potential-dependent transformation from bivalent copper to monovalent copper, in which the existence of univalent copper is beneficial to inhibit hydrogen generation and regulate the Gibbs free energy of NO₃⁻RR intermediates. It is worth noting that the active intermediate state of monovalent copper is unstable and can be rapidly reduced to metallic copper under high reduction potential (<−0.5 V vs. RHE), resulting in a large number of absorbed intermediates (e.g., NO₂⁻ and NO) passivated surface, which not only leads to the reduction of catalytic activity but also causes great troubles for the rational design and mechanism analysis of the catalyst. In addition, inspired by the tandem NO₃⁻to NH₃ conversion in nature, multi-phase interactions may circumvent the scaling relations to achieve ammonia production. Even though the Cu alloying strategy (e.g., CuRh, CuNi and CuPd, etc.) can achieve excellent FE (more than 80%) and ammonia selectivity (more than 70%), the overpotential is still greater than −0.5 V because of the scaling relationships (LSR). Such a high overpotential is detrimental to the survival of the real active sites and increases energy consumption. Therefore, it is important to stabilize the real active sites of the catalyst with a low applied overpotential, which is vital for guiding the rational design and mechanism analysis of certain catalytic materials for NO₃—RR.

In view of the above, there is an unmet need for an effective catalyst and methodologies for the synthesis of ammonia from nitrate. The present invention seeks to provide a solution to the aforementioned challenges, or at least to provide an alternative to the public.

BRIEF DESCRIPTION OF THE INVENTION

According to a first aspect of the present invention, there is provided a method of preparing a catalyst for use in synthesizing ammonia from nitrate, comprising the steps of i) providing a nickel foam material, ii) removing an oxide layer on the nickel foam material, iii) subjecting the nickel foam material to cooper sulphate (Cu₂SO₄) solution for a predetermined period of time for forming a copper (Cu) dopped nickel foam electrode, and iv) subjecting the Cu dopped nickel foam electrode to an annealing treatment at a predetermined temperature for a predetermined period of annealing time in atmospheric air thus forming a catalyst of copper oxide-nickel oxide nickel foam (CuO/NiO@NF) with a heterostructure.

Preferably, the step of removing the nickel oxide (NiO_x) layer on the nickel foam material may be achieved by sonicating the nickel foam material in an acid solution for a predetermined period of time. The acid solution may be hydrochloric acid (HCl) solution, sulfuric acid (H₂SO₄) solution, or nitric acid (HNO₃) solution, the concentration of the acid may preferably be 0.5-1 M, and the period of time of the sonicating may preferably be 5⁻10 minutes. After the step of removing the NiO_X, there may preferably be a step of removal of residuals from the acid solution, wherein the residual removal may be achieved by rinsing with acetone or water.

Suitably, the Cu₂SO₄ solution may have a concentration of 40-60 mM but preferably 50 mM and the period of time may be 30-40 minutes but preferably 30 minutes.

Advantageously, the predetermined temperature of the annealing temperature may range from 300° C. to 600° C., or 300° C., 400° C., 500° C. or 600° C., and the predetermined annealing time may be 30-60 mins. Preferably, the predetermined annealing temperature may be 400° C.

According to a second aspect of the present invention, there is provided a method of preparing an electrolytic cell for use in synthesizing of ammonia, comprising a step of i) providing the catalyst CuO/NiO@NF made according to a method as described above, a platinum foil, and a saturated Ag/AgCl electrode, wherein the catalyst CuO/NiO@NF, the platinum foil, and the saturated Ag/AgCl electrode act as a working electrode, a counter electrode, and a reference electrode, respectively in the electrolytic cell.

Preferably, the catalyst CuO/NiO@NF may be in the formation of a sheet and has a dimension of 1×1 cm.

Suitably, the method may comprise a step of providing a sodium sulfate (Na₂SO₄) solution as a catholyte and an anolyte of the electrolytic cell for use in a cathode chamber and an anode chamber of the electrolytic cell. The concentration of the (Na₂SO₄) solution may be 0.5-1 M. The method may comprise a step of providing a proton exchange membrane between the cathode chamber and an anode chamber of the electrolytic cell.

Advantageously, the proton exchange membrane may be a Nafion® 117 membrane.

According to a third aspect of the present invention, there is provided a method of synthesis of ammonia from nitrate, comprising the steps of i) providing an electrolytic cell made as described above, and ii) prior to using the electrolytic cell, purging all solutions used in preparing the electrolytic cell preparation and the electrode preparation including the Na₂SO₄ solution with pure nitrogen (N₂) gas for a predetermined period of purging time.

Preferably, the purging time may be 10 mins.

Suitably, the method may comprise a step of adding a source of nitrate (NO₃) to the cathode chamber and allowing reaction to take place for generating ammonia (NH₃).

Advantageously, the method may comprise a step of collecting the generated ammonia (NH₃) from the cathode chamber.

BRIEF DESCRIPTION OF THE DRAWINGS

The patent or application file contains at least one drawing executed in color. Copies of this patent or patent application publication with color drawing(s) will be provided by the Office upon request and payment of the necessary fee.

Some embodiments of the present invention will now be explained, with reference to the accompanied drawings, in which:

FIGS. 1a to 1i illustrate the synthesis and morphology of the heterostructure of a CuO/NiO catalyst, in which, FIG. 1a is a schematic diagram showing synthesis of the CuO/NiO heterostructure; FIGS. 1b and 1c are surficial and cross-sectional field-emission scanning electron microscope images of the CuO/NiO heterostructure; FIG. 1d is a HRTEM image; FIG. 1e is an image showing the SAED pattern of CuO/NiO, where different component crystalline phases are marked out; and FIGS. 1f to 1i are a HRTEM image and the corresponding HRTEM-EDS elemental mapping images of CuO/NiO relating to Ni, Cu and 0, respectively.

FIGS. 2a to 2h illustrate the chemical and electronic structures of the composite of the CuO/NiO heterostructure, in which FIG. 2a is a Raman spectra and FIG. 2b is a XRD pattern of CuO/NiO-300, CuO/NiO-400, CuO/NiO-500, and CuO/NiO-600, FIG. 3c are energy band diagrams of CuO and NiO before contact and after formation of CuO/NiO heterojunction, FIG. 2d and FIG. 2e are a high-resolution XPS spectra of Cu 2p for CuO/NiO and CuO, and Ni 2p for CuO/NiO and NiO, respectively, FIG. 2f is a schematic diagram showing a calculated charge density difference of CuO/NiO, in which the Cu, Ni, and O atoms are marked in blue, grey, and red, respectively, and the yellow and indigo blue regions represent the accumulation and depletion of charge density, respectively; FIG. 2g is a graph showing an electronic density of states (DOS) calculated for CuO/NiO (CuO) and CuO; and FIG. 1h is a graph showing a DOS calculated for CuO/NiO (NiO) and NiO.

FIGS. 3a to 3i illustrate the identification of the reconstructuion of CuO/NiO catalysts, in which FIG. 3a is schematic illustration of in situ electrochemical Raman spectra of phase transformation of CuO/NiO during electrochemical NO₃⁻RR process; FIGS. 3b and 3c are graphs showing in situ Raman spectra of the CuO/NiO catalyst; FIGS. 3d and 3e graphs showing in situ Raman spectra of the CuO and NiO, respectively, at different applied potentials in 0.5 M Na₂SO₄ electrolyte with 200 ppm NO₃—N; FIG. 3f is a graph showing in situ Raman spectra of the CuO/NiO catalyst at different applied potentials in 0.5 M Na₂SO₄ electrolyte without NO₃—N; FIG. 3g is a surficial field-emission scanning electron microscope image of CuO/NiO catalysts after 2 h electrochemical reduction; and FIG. 3h are FIG. 3i illustrate HRTEM and (i) SAED pattern, respectively, of CuO/NiO catalysts after 2 h electrochemical reduction, where different component crystalline phases are marked out.

FIGS. 4a to 4h illustrate the electrocatalytic NO₃⁻— to —NH₃ conversion over CuO/NiO; in which FIG. 4a is a schematic illustration of electrochemical NO₃⁻RR in an H-type electrolytic cell; FIG. 4b is a graph showing LSV curves of CuO/NiO, CuO, and NiO in 0.5 M Na₂SO₄ electrolyte with 200 ppm NO₃⁻—N upon a scan rate of 5 mV s⁻¹; FIG. 4c is a graph showing time-dependent concentration change of NO₃—, NO₂⁻ and NH₃ over CuO/NiO; FIG. 4d is a graph showing the FE and yield rate of NH₃ at different applied potentials in 0.5 M Na₂SO₄ electrolyte with 200 ppm NO₃⁻—N over a CuO/NiO cathode; FIG. 4e is a graph showing the FE and yield rate of NH₃ in consecutive recycling tests over a CuO/NiO cathode; FIG. 4f is a graph showing the ¹H NMR spectra of the electrolyte after NO₃⁻RR by using ¹⁴NO₃⁻—¹⁴N/¹⁵NO₃⁻—¹⁵N as the nitrogen source; FIG. 4g is a graph showing the standard curve of the integral area against ¹⁵NH₄+—¹⁵N concentration; and FIG. 4h is a graph showing a comparison of NH₃ Faradaic efficiency, NO₃⁻to NH₃ selectivity and NH₃ yield rate for recently reported state-of-the-art NO₃RR electrocatalysts including Cu/Cu₂O, PdCu/Cu₂O, Cu/Cu—Mn₃O₄, Co/CoO, CuPd, TiO_2-x, Fe/Cu, and Fe single atom catalyst.

FIGS. 5a to 5g illustrate DFT calculations for the electronic structures and reaction mechanism; in which FIG. 5a is a schematic diagram showing the calculated charge density difference of Cu₂O/NiO, in that the Cu, Ni, and 0 atoms are marked in blue, grey, and red; FIGS. 5b to 5d are schematic diagrams showing the charge density difference for NO₃* adsorption on Cu₂O/NiO, CuO, and NiO, respectively; FIG. 5e is a free-energy diagram for NO₃⁻RR on the interface of Cu₂O/NiO (red line), CuO region in Cu₂O/NiO (blue line), and NiO region in Cu₂O/NiO (yellow line) at U=0 V; FIG. 5f is a graph showing partial electronic density of states (PDOS) calculated for NO₃* and NH₃* adsorption on Cu₂O/NiO; and FIG. 5g are free-energy diagrams for NO₃⁻RR on the surfaces of Cu₂O/NiO (red line), Cu (blue line), and Ni (yellow line) at U=0 V (insets show the optimized configurations for NO₃⁻RR adsorbates. Blue, grey, red, white, and pink spheres represent Cu, Ni, 0, N, and H atoms, respectively).

FIG. 6 is a graph illustrating a XRD pattern of Cu-doped Ni foam synthesized via a galvanic displacement reaction and pristine Ni foam.

FIGS. 7a to 7d are optical images of four materials of pristine NF (FIG. 7a), CuO/NiO (FIG. 7b), CuO (FIG. 7c), and NiO (FIG. 7d), respectively.

FIG. 8 is a Field-emission scanning electron microscope image of CuO/NiO.

FIGS. 9a to 9f are surficial and cross-sectional field-emission scanning electron microscope images of CuO/NiO-300 (FIGS. 9a, b), CuO/NiO-500 (FIGS. 9c, d), and CuO/NiO-600 (FIGS. 9e, f), respectively.

FIGS. 10a to 10c are graphs showing high energy cut-off of UPS spectra recorded, in which FIG. 10a is the high energy cut-off of UPS spectra for CuO; FIG. 10b is the high energy cut-off of UPS spectra for NiO; and FIG. 10c is the high energy cut-off of UPS spectra for CuO/NiO. The inserts in the graphs shows the onset energy region.

FIGS. 11a to 11j are charts showing in situ Raman spectra of the pre-synthesized CuO/NiO catalyst in consecutive recycling tests.

FIGS. 12a to 12c are charts showing in situ Raman spectra of the CuO/NiO catalyst in 2 h electrolysis.

FIG. 13 is a chart illustrating XRD pattern of CuO/NiO catalysts after 2 h electrochemical reduction.

FIGS. 14a to 14c are graphs showing LSV curves of CuO/NiO (FIG. 14a), CuO (FIG. 14b), and NiO (FIG. 14c), respectively, in 0.5 M Na₂SO₄ electrolyte with and without 200 ppm NO₃⁻—N upon a scan rate of 5 mV s⁻¹.

FIG. 15a is a graph showing Faradaic efficiencies for NH₃, and FIG. 15b is a graph showing and current densities of NH₃, on CuO/NiO, CuO and NiO at various applied potentials.

FIG. 16a to 16c are graphs showing cyclic voltammetry (CV) cycles at different scan rates from 10 to 60 mV s⁻¹ CuO/NiO, CuO and NiO, respectively; FIG. 16d is a graph showing the corresponding current density differences vs. scan rates of the samples to calculate electrochemical double-layer capacitance (C_dl) (The C_dl is proportional to the electrochemical surface area (ECSA); and FIG. 16e is a graph showing ECSA normalized partial current densities of NH₃ in LSV curves.

FIG. 17a is a graph showing UV-Vis absorption curves using a series of standard concentration of KNO₃ solutions; FIG. 17b is a graph showing the concentration-absorbance calibration curve for determining the concentration of NO₃⁻(the inset in is a photograph of KNO₃ solutions mixing with the color reagent; FIG. 17c is a graph showing UV-Vis absorption curves of the Griess method using a series of standard concentration KNO₂ solutions; FIG. 17d is a graph showing the concentration-absorbance calibration curve for determining the concentration of NO₂⁻(the inset is a photograph of KNO₂ solutions mixed with the color reagent and sitting for 20 min); FIG. 17e is a graph showing UV-Vis absorption curves of the indophenol blue method using a series of standard concentration NH₄(SO₄)₂ solutions; and FIG. 17f is a graph showing the concentration-absorbance calibration curve for determining the concentration of NH₄+(the inset in is a photograph of (NH₄)₂SO₄ solutions mixed with the color reagent and sitting for 2 h).

FIGS. 18a to 18c are graphs showing UV-Vis absorption curves of NO₃—; FIGS. 18d to 18f are graphs showing UV-Vis absorption curves of NH₃; and FIGS. 18g to 18i are graphs showing UV-Vis absorption curves of NO₂, all in time-dependent 2 h electrolysis.

FIGS. 19a to 19c are graphs showing UV-Vis absorption curves of NH₃ at different applied potentials.

FIG. 20 is a graph showing UV-Vis absorption curves of NH₃ in consecutive recycling tests.

FIG. 21 is a graph of the i-t curves of CuO/NiO in 2 h electrolysis at −0.2 V vs RHE.

FIG. 22a is a graph showing the 1H NMR spectra of the electrolyte after NO₃⁻RR by using 15NO₃⁻—¹⁵N as the nitrogen source; and FIG. 22b is a graph showing the 1H NMR spectra of ¹⁵NH₄+ using 15(NH₄)₂SO₄ with different known concentrations and the same concentration of C₄H₄O₄ as standards.

FIG. 23a and FIG. 23b are schematic diagrams showing optimized geometric structure of intermediates absorbed on CuO region in Cu₂O/NiO (FIG. 23a) and NiO region in Cu₂O/NiO (FIG. 23b).

FIGS. 24a to 24b are graphs showing partial electronic density of states (PDOS) calculated for NO₃* (FIG. 24a) and NH₃* (FIG. 24b) adsorption on Cu₂O and NiO.

FIG. 25 is a table showing NO₃⁻—N Removal Rate in 2 h electrolysis.

DETAILED DESCRIPTION OF PREFERRED EMBODIMENTS OF THE INVENTION

Electrochemically converting waste nitrate (NO₃—) into ammonia (NH₃) is a green route for both wastewater treatment and high-value-added ammonia generation. However, the NO³⁻— to —NH₃ reaction involves multi-step electron transfer and complex intermediates, making it a grand challenge to drive efficient NO³⁻ electroreduction with high NH₃ selectivity. Herein, an in-operando electrochemically synthesized Cu₂O/NiO heterostructure electrocatalyst is proven for efficient NH₃ electrosynthesis. In situ Raman spectroscopy reveals that the obtained Cu₂O/NiO, induced by the electrochemistry-driven phase conversion, is the real active phase. This electronically coupled phase can modulate the interfacial charge distribution, dramatically lower the overpotential in the rate-determining step and thus requiring lower energy input to proceed with the NH₃ electrosynthesis. The orbital hybridization calculations further identify that Cu₂O is beneficial for NO³⁻ adsorption, and NiO could promote the desorption of NH₃, forming an excellent tandem electrocatalyst. Such a tandem system leads to NH₃ Faradaic efficiency of 95.6%, a super-high NH₃ selectivity of 88.5% at −0.2 V vs. RHE, surpassing most of the NH₃ electrosynthesis catalysts at an ultralow reaction voltage.

The present invention has made use of a in situ electrochemistry-driven phase conversion strategy to reconstruct and stabilize the active Cu₂O/NiO intermediate phase via a binary metal oxide CuO/NiO. In situ Raman spectra, Auger spectra, and high-resolution transmission electron microscopy (HRTEM) characterizations revealed that the Cu₂O/NiO active phase was reconstructed and stabilized at low overpotentials during the whole NO₃⁻RR process. Density functional theory (DFT) calculations exhibit that the Cu₂O/NiO interface is the optimal active site for NO₃⁻RR due to the electron-coupled interactions between Cu₂O and NiO. Those two active phases can also be combined into a tandem system to break the adsorption-energy scaling relationship, which exhibits strong NO₃⁻ adsorption at Cu₂O and rapid NH₃ release at NiO. As a result, the intermediate active phase of Cu₂O/NiO reveals a superior electrocatalytic NO₃⁻— to —NH₃ activity at −0.2 V vs. RHE with a FE of 95.6%, a high NH₃ selectivity of 88.5% and the ammonia-evolving rate of 2.1 mol m⁻¹ m⁻², surpassing most of the NO₃⁻RR catalysts at an ultralow reaction voltage.

Experiments

Preparation of the CuO/NiO@NF electrode: The CuO/NiO@NF electrode was prepared through two steps: (i) The Cu—NF electrode was synthesized through galvanic replacement between Ni and CuSO₄ on a NF matrix driven by the difference in the reduction potentials of Cu versus Ni. (ii) The Cu-doped NF electrode was converted into CuO/NiO@NF via an annealing treatment. Before galvanic replacement, the nickel foam (3.0×1.0 cm) was sonicated in HCl solution (1 M) for 10 min to remove the NiO surface layer and followed by rinsing with water and acetone for three times, then allowed to blow-dry in with the nitrogen gas. The galvanic replacement was carried out by dipping the cleaned NF into 50 mM CuSO₄ solutions and kept for 30 min. The obtained Cu—NF electrode was finally annealed at 300° C., 400° C., 500° C., and 600° C. for 30 min in an air atmosphere, where they were abbreviated as CuO/NiO-300, CuO/NiO-400, CuO/NiO-500 and CuO/NiO-600, respectively. The CuO@NF and NiO@NF electrode were prepared through directly annealing copper foam (CF) and NF for 30 min at 400° C. under the air atmosphere.

Characterization: SEM images and EDX spectra were obtained with FEI Quanta 450 FESEM analysis system. The cross-sectional sample is artificially pre-cut and prepared vertically with a sharp scissor. Moreover, high-resolution transmission electron microscopy (HRTEM), selected area electron diffraction (SAED), and elements mapping were obtained with JEOL JEM-2100F equipment. Raman spectroscopy was employed in a WITec alpha 300 R Raman System with a 532 nm laser excitation source. The crystalline structure was confirmed by Powder X-ray diffraction with a Bruker D8 Focus Diffraction System with a monochromatized Cu-Kα radiation source (λ=0.1541 nm). The elemental composition and valance state were measured on X-ray photoelectron spectroscopy (XPS) analyses equipped with a PHI 5000 Versaprobe system and monochromatic AI Kα radiation. All binding energies were calibrated on the C 1s peak at 284.8 eV. Ultraviolet photoelectron spectroscopy (UPS) measurements were performed with an He 1 (21.22 eV) line as an excitation source. The work function (0) was examined by the secondary electron cut-off (E_cutoff) in the UPS spectrum (equivalent to 21.22 eV-E_cutoff). The valence band maximum (E_V) was determined from the low binding energy onset. The conduction band energy (E_C) positions of CuO and NiO are determined by E_V and the energy gap (E_g) obtained in the reported literature.^[34] Isotope labeling experiments were carried out on Nuclear Magnetic Resonance measurement (Bruker 400-MHz system).

Electrochemical Measurements: All the electrochemical measurements were conducted on a CHI660D electrochemical workstation in an H-type electrolytic cell separated by a Nafion 117 membrane. The CuO/NiO@NF (1×1 cm), platinum foil, and saturated Ag/AgCl electrode acted as the working electrode, counter electrode, and reference electrode, respectively. 0.5 M Na₂SO₄ solution (40 mL) was used as catholyte and anolyte. 200 ppm nitrate-N was also added into the cathode compartment for the nitrate reduction test. All the solutions used in the electrochemical nitrate reduction were purged with high-purity nitrogen (N₂) for 10 min. Before the nitrate electroreduction test, linear sweep voltammetry (LSV) curves were recorded at a rate of 5 mV s⁻¹. For nitrate electroreduction experiments, potential-controlled measurements were conducted for 2 h in each potential at a stirring rate of 400 rpm. Three replicate electrochemical experiments were conducted on each sample to determine the error bar. The electrolyte was collected to identify the product of nitrate reduction by the ultraviolet-visible (UV-Vis) spectrophotometer.

The potential (versus saturated Ag/AgCl) was converted to RHE using the Nernst equation (3):

$\begin{array}{cc}{E}_{\mathrm{RHE}}=E_{\mathrm{Ag}/\mathrm{AgCl}}+0.0591\times \mathrm{pH}+E_{\mathrm{Ag}/\mathrm{AgCl}}^0\left(E_{\mathrm{Ag}/\mathrm{AgCl}}^0=0.197\right)& \left(3\right)\end{array}$

Determination of NO₃⁻: An amount of electrolyte was collected and diluted to the range of detection. Then, the diluted electrolyte (10 mL) was mixed well with 100 μL 1 M HCl and 10 μL 0.8 wt % sulfamic acid solution. The absorption spectrum of NO₃ was determined at 220 nm and 275 nm by an ultraviolet-visible spectrophotometer. The absorbance value at 220 nm detected both organic nitrate and chemical signals, while the absorbance value at 275 nm only contained signals from the organics. The final absorbance value was calculated by the absorbance at 220 nm minus the absorbance at 275 nm. In order to quantify the concentration, the concentration-absorbance curve was plotted using a series of standard KNO₃ solutions.

Determination of NO₂: The concentration of NO₂⁻ was determined by Griess method. The color reagent of Griess method was prepared as follows: Adding 4 g p-aminobenzene sulfonamide and 0.2 g N-(1-Naphthyl) ethylenediamine dihydrochloride to the mixed solution of 10 mL phosphoric acid and 50 mL deionized water. A certain amount of electrolyte was diluted to the detection range and 5 mL diluted electrolyte was added with 100 μL color reagent. After standing for 20 min in a dark place, the absorbance value at a wavelength of 540 nm was determined by UV-Vis absorption spectra. The concentration-absorbance curve was calibrated using a series of standard concentration KNO₂ solutions.

Determination of NH₃: Similar to the method of confirming the concentration of NO₃⁻ and NO₂—, NH₃ was determined by the indophenol blue approach. First, a certain amount of electrolyte was diluted to the detection range. Then, 3 mL diluted electrolyte was taken out into a test tube. Next, 1 mL 2.0 M NaOH solution containing 10.0 wt % sodium citrates (C₆H₅Na₃O₇·2H₂O) and 10.0 wt % salicylic acid (C₇H₆₀₃) was added and mixed thoroughly, followed by adding 1 mL 0.05 M NaClO solution and 0.2 mL 1.0 wt % sodium nitroferricyanide (Na₂[Fe (NO)(CN)₅]·2H₂O) solution. The absorbance value at a wavelength of 658 nm was recorded after sitting for 2 h. The concentration-absorbance calibration curve was calibrated using a series of standard concentration NH₄(SO₄)₂ solutions.

Calculation of the Faradaic efficiency and yield rate: Faradaic efficiency of NH₃ was calculated by using equation (4):

$\begin{array}{cc}{\mathrm{FE}}_{\mathrm{NH}3}\left(\%\right)=\frac{8F\times c_{{\mathrm{NH}}_3}\times V}{17\times Q}& \left(4\right)\end{array}$

NH₃ yield rate was calculated by using equation (5):

$\begin{array}{cc}{\mathrm{NH}}_3\mathrm{yield}\mathrm{rate}\left(\mathrm{mg}{h}^{-1}{\mathrm{cm}}^{-2}\right)=\frac{c_{{\mathrm{NH}}_3}\times V}{t\times s}& \left(5\right)\end{array}$

The percentage of NO₃⁻ conversion was calculated by using equation (6):

$\begin{array}{cc}{\mathrm{NO}}_{3^{-}}\mathrm{conversion}\left(\%\right)=\frac{c_0\times c_t}{c_0}\times 100& \left(6\right)\end{array}$

The NH₃ selectivity was evaluated by using equation (7):

$\begin{array}{cc}{\mathrm{NH}}_3\mathrm{selectivity}=\frac{X_i}{c_0\times c_t}\times 100& \left(7\right)\end{array}$

where F is Faraday constant (96485 C mol⁻¹), c_NH3 (mg L⁻¹) is the NH₃ concentrations after electrolysis at the reduction time t, V (mL) is the volume of the cathodic electrolyte, Q (C) is the overall charge passing the electrode, s is the geometric area of catalyst (cm²), c₀ and c_t are the measured NO₃⁻ concentrations before and after electrolysis at the reduction time t, and X_i is the concentration of NH₃ after electrolysis.

In situ Raman spectroscopy: In situ Raman spectroscopy was performed by the aforementioned Raman microscope and a three-electrode electrochemical cell. A Pt wire and Ag/AgCl electrode were used as counter and reference electrodes, respectively. CuO/NiO and CuO electrodes were used as working electrodes and immersed into the 0.5 M Na₂SO₄ electrolyte to ensure sufficient ionic conductivity. The existence and non-existence of 200 ppm nitrate-N were also tested as a contrast.

¹⁵N and ¹⁴N isotope-labeling experiment: The isotope-labeling nitrate reduction experiments were conducted on a Bruker 400 MHz Nuclear Magnetic Resonance System. 98.5% Na¹⁵NO3 was used as the feeding N-source and 0.5 M Na₂SO₄ was used as the electrolyte. After electroreduction, the obtained electrolyte (25 ml) containing ¹⁵NH₄⁺—¹⁵N was taken out and adjusted to be weak acid with H₂SO₄, followed by mixing with 0.01 g maleic acid (C₄H₄O₄) as an internal standard. Then, 50 μL deuterium oxide (D20) was added to 0.5 mL above solution for the NMR detection. 50, 100, 150, 200, and 250 ppm ¹⁵NH₄⁺-¹⁵N were prepared for achieving the calibration curve. Similarly, Na¹⁴NO₃/Na¹⁵NO₃ equals to 1:1 was also used as N-source in the nitrate reduction methods, and then the ratio of products ¹⁴NH₄⁺-¹⁴N/¹⁵NH₄⁺-¹⁵N was quantified by NMR detection to further clarify the source of ammonia.

Computational details: The first principles density functional theory (DFT) calculations were carried out using the Vienna ab initio Simulation Package (VASP) and the Perdew-Burke-Ernzerhof (PBE) functional within generalized gradient approximation (GGA).^[621] The projector-augmented wave (PAW) method was applied in plane waves with a cutoff energy of 500 eV. The 3×3×1 k-point mesh set was used for geometry optimization, self-consistent, and the density of states (DOS) calculations. A vacuum layer of 15 Å was used for Cu₂O/NiO (200), Cu (100), and Ni (100) slab models. The crystal Cu₂O/NiO was constructed by cleavage of (200) surface according to the TEM results. Cu (100) and Ni (100) are the most active surface for the electrochemical NO₃⁻ reduction.^[63] In all calculations, the electronic iterations convergence was optimized to 10⁻⁵ eV with the Hellmann-Feynman forces of 0.02 eV Å⁻¹. VESTA and VASPKIT were applied to obtain the DOS diagrams.^[64]

The NO₃ reduction process can be described as following steps (8-15):

*+NO₃⁻→*NO₃⁺e⁻  (8)

*NO₃⁺H₂O⁺e⁻→*NO₂⁺2OH⁻  (9)

*NO₂⁺H₂O+2e⁻→*NO+2OH⁻  (10)

*NO+H₂O+2e⁻→*N+2OH⁻  (11)

*N+H₂O⁺e⁻*NH+OH⁻(12)

*NH+H₂O⁺e⁻→*NH₂⁺OH⁻  (13)

*NH₂⁺H₂O⁺e⁻→*NH₃⁺OH⁻  (14)

*NH₃→NH₃⁺*  (15)

where * represents the active site. The reaction free energy change was calculated as follows (16):

$\begin{array}{cc}\Delta G=\Delta E+\Delta E_{\mathrm{ZPE}}-T\Delta S& \left(16\right)\end{array}$

where ΔE is the total energy difference. ΔE_ZPE and ΔS are the corrections of the zero-point energy and entropy, which are obtained from the vibrational frequency calculations. T is the temperature and equals to 298.15 K.To avoid calculating the energy of charged NO₃⁻ directly, HNO₃(g) is chosen to be a reference as below equation (17-18):

HNO₃(g)→HNO₃(l)  (17)

HNO₃(l)→H⁺+NO₃⁻  (18)

Hence, the equation (6) can be rewritten as below (19):

*+HNO₃-→*NO₃⁺H⁺+e⁻  (19)

The Gibbs free energy of NO₃ (ΔG(*NO₃)) is described as (20):

$\begin{array}{cc}\Delta G\left(\mathrm{NO}{3}^{*}\right)=G{(}^{*}{\mathrm{NO}}_3)-G\left({\mathrm{HNO}}_3\right)+\frac{1}{2}G\left(H_2\right)-G{(}^{*})+\Delta G_{\mathrm{correct}}& \left(20\right)\end{array}$

in which G(*NO₃), G(HNO₃), and G (*) are the DFT-calculated Gibbs free energy of adsorbed nitrate, HNO₃, and catalysts, respectively. ΔG_correct is the correction of adsorption energy.

Results and Discussion

The fabrication of CuO/NiO heterostructures on Ni foam is illustrated in FIG. 1a. First, the Cu-doped Ni foam (NF) was synthesized via a galvanic displacement reaction. The CuO/NiO heterostructures on NF were then obtained through annealing in the air (a detailed fabrication process was shown in Experimental Section). Comparing the XRD patterns of Cu-doped NF and pristine NF, the diffraction peaks exhibit a slight shift after galvanic displacement, confirming that Cu atoms successfully substitute Ni counterparts (FIG. 6). During synthesis, an obvious color change in NF from silver to black was observed after the annealing treatment (FIGS. 7a to 7d). Based on the scanning electron microscopy (SEM) characterizations, the CuO/NiO layer with an average thickness of ≈100 nm could be in situ formed in the nickel substrate (FIG. 1b and FIG. 1c). Meanwhile, the framework of NF could be well preserved to achieve sufficient electrochemically active surface area for enhanced NO₃⁻RR performance (FIG. 8). The morphology of CuO/NiO heterostructure was also characterized by the HRTEM and the corresponding selected area electron diffraction (SAED) (FIG. 1d and FIG. 1e). The well-defined CuO phase with an interplanar spacing of 0.232 nm for the (111) plane is closely connected with the (200) and (111) facets of NiO with the lattice fringes distances of 0.209 and 0.241 nm, respectively. The concrete phase will be indicated in subsequent XRD tests. Clear boundaries could as well be observed between CuO and NiO phases. In addition, the SAED pattern shows the diffraction rings with the random arrangement of bright spots, matching well with the (111) and (110) planes of CuO and the (111) and (220) planes of NiO. The element mapping presents the uniform distribution of Cu and Ni on the electrocatalyst surface, indicating the successful formation of the CuO/NiO tandem catalyst (FIGS. 1f-i).

To describe the formation of CuO/NiO heterostructures, a series of temperature-controlled catalysts were synthesized at 300, 400, 500, and 600° C. The cross-sectional SEM images depicted that the thicknesses of in situ formed layers increased along with the raised temperature (FIGS. 9a to 9b). Raman spectra of CuO/NiO heterostructures synthesized at different temperatures are displayed in FIG. 2a. CuO obtained at 300° C. and the bands indexed at 281, 330, and 605 cm⁻¹ were assigned to the Ag, Bg′, and Bg″ modes of CuO, which confirms the formation of single phase and highly crystalline CuO structures at the low temperature. The longitudinal optical (LO) of NiO (Ni^ll—O) at 510 cm⁻¹ could also be observed when increasing the temperature up to 400° C.^[31] As the calcination temperature further increases, Ni^lll—O vibrations appear at 475 and 550 cm⁻¹.^[32] XRD patterns were as well conducted to further reveal the synthetic process of CuO/NiO heterostructures (FIG. 2b). The sample synthesized at 300° C. showed several diffraction peaks at 35.42°, 38.71°, 44.51°, 51.85°, 76.37°, and 86.53°, corresponding to (002), (111), (40-2) lattice planes of CuO (JCPDS no. 48-1548) and (111), (200), (220) lattice planes of Ni (JCPDS no. 04-0850), respectively. Since the penetration depth of XRD is larger than the total thickness of the CuO/NiO layer, the diffracting peaks of the nickel foam substrate could also be observed. Apart from the diffraction patterns of CuO and Ni, the diffraction patterns of the sample synthesized at 400° C. demonstrated the formation of fcc-structured NiO as evidenced by the (111) and (200) planes, respectively, at 37.2° and 43.3° (JCPDS no. 71-1179).

To shed light on the electron distribution at the interface of CuO/NiO heterostructures, the band structures, chemical states, and electronic interactions were thoroughly investigated. First, the ultraviolet photoelectron spectroscopy (UPS) measurements were performed to ascertain the work functions (Φ) and the valence band maximum (E_V) of CuO, NiO, and CuO/NiO heterostructures (Figure S5, Supporting Information). The work function of CuO was calculated to be 4.41 eV, which is larger than that of NiO (4.16 eV), as presented in FIGS. 10a to 10c. Hence, the Fermi level (E_F) versus the vacuum level of CuO was calculated to be −4.41 eV, below that of NiO (−4.16 eV). Next, the E_V position of CuO and NiO are found to be located at −4.58 eV and −4.61 eV, accordingly (inset in FIG. 10a and FIG. 10b. Then, the conduction band energy (E_C) positions of CuO and NiO are determined to be −2.71 eV and −0.66 eV. Thus, the energy band structures of the CuO—NiO p-p heterojunctions could be compiled and depicted in FIG. 2c. Because the position of the Fermi level of NiO is higher than that of CuO, electrons will flow from NiO to CuO until a Fermi equilibrium state. The electron flow leads to the opposite space-charge region at the CuO/NiO heterostructure interface and further generates the built-in electric field. The strongly coupled heterointerfaces with built-in electric fields can facilitate interfacial charge migration and accelerate NO₃⁻RR reaction kinetics. The element valence states of CuO/NiO heterostructures were further elucidated by X-ray photoelectron spectroscopy (XPS) (FIG. 2d and FIG. 2e). The Cu 2p3/2, Cu 2p1/2 peaks and the related satellite peaks of CuO and CuO/NiO could be ascribed to CuO species. It is worthwhile to point out that those characteristic Cu²⁺ peaks of the CuO/NiO composites shift to low binding energy compared with those of CuO. The core-level Ni 2p spectrum in NiO and CuO/NiO exhibit the characteristic Ni 2p3/2, Ni 2p1/2 peaks and the related satellite peaks, indicating the presence of Ni²⁺ with small amounts of Ni³⁺. The Ni 2p peaks of CuO/NiO composites present a slight positive shift compared with those of bare NiO components. Those slight shifts provide important evidence for successfully constructing the electronic coupled interface between CuO and NiO, which leads to charge transfer from NiO to CuO. This electron movement tendency is consistent with the UPS analysis mentioned above. To obtain further insights, density functional theory (DFT) calculations were employed to analyze the electronic distribution of CuO/NiO heterostructures. In this case, the charge density map of CuO/NiO heterostructure was calculated (FIG. 2f). The yellow and indigo-blue regions represent the accumulation and depletion of charge density, respectively. Apparently, the interface of CuO/NiO gathers with large electron clouds, which implies the formation of a built-in electric field at the interface, leading to the continuous electron flow between two active phases. Thus, the interface region of the CuO/NiO heterostructure will be the optimal adsorbed sites for intermediate species. The electronic modulation of the CuO/NiO is also reflected by the changed d-band density of states (d-DOS). Based on FIG. 2g, we can see that the d-band center of CuO in the CuO/NiO heterostructure (−2.28 eV) shows a downshift compared with that in the pure-phased CuO (−2.21 eV). In contrast, the d-band center of the NiO-phase of the CuO/NiO heterostructure is about −1.11 eV, which is much closer to the Femi level compared to that of the pure-phased NiO (−1.21 eV) (FIG. 2h). This altered DOS of the CuO-phase and NiO-phase of the CuO/NiO heterostructure may attribute to a various hybridization between the d orbitals of CuO/NiO and 2p orbitals of the absorbed NO₃RR intermediate species, thus exhibiting a considerable difference in the covalent interaction among them compared to that of CuO and NiO. These characterizations and calculations, taken together, suggest that the interface of CuO/NiO heterostructures has the potential to improve the NO₃⁻RR performance.

Electrocatalytic NO₃⁻— to —NH₃ conversion proceeds at reduction potentials, at which CuO/NiO catalysts suffer from potential-dependent phase reconstruction, leading to the formation of multi-intermediate phase. In situ Raman spectroscopy was utilized to monitor the self-reconstruction behavior of the pre-synthesized CuO/NiO catalysts during this process (FIG. 3a). NaNO₃ was added for the nitrate reduction test, and Na₂SO₄ was used as an electrolyte. Raman spectroscopy of the CuO/NiO catalyst at different applied potentials (FIG. 3b and FIG. 3c) show the initial characteristic peaks at 298 cm⁻¹, associated with CuO phases, persist at as low as −0.2 V vs RHE. Remarkably, at −0.15 V vs RHE, CuO phases gradually disappeared, and two sets of characteristic peaks of Cu₂O emerged at 148 and 218 cm⁻¹. Further decreasing the potential, the Raman signals of Cu₂O get weakened until they completely disappear, indicating a totally reduced Cu₂O phase. Meanwhile, the broad Raman peaks of NiO appear at a wide range of around 510 cm⁻¹ corresponding to the LO. To further clarify the surface phase compositions during the whole NO₃⁻RR, in situ Raman spectroscopy was applied to monitor the phase transformation under potentiostatic operation for 2 h at −0.2 V (FIG. 11a). The quickly attenuated CuO peak and the enhanced peak intensity of Cu₂O provide strong evidence of the partial reduction of CuO to Cu₂O during the first 10 min. With the prolonged reaction time, the peak of CuO disappears, while the Cu₂O peaks are maintained during the 2 h electrolysis, which impedes the initial CuO phase totally transferred to the stable Cu₂O phase as the reaction progressed. The broad Raman peaks of NiO also remain unchanged. The coupling of Cu₂O and NiO could be confirmed as the main active catalyst for NO₃⁻— to —NH₃ conversion. For the phase transformation of the pre-synthesized CuO electrocatalyst, the Raman peak associated with Cu—O in the CuO catalyst quickly disappears, and the two sets of peaks assigned to the Cu₂O phase emerged at 148 and 218 cm⁻¹ disappear at −0.15 V which is much higher than the active Cu₂O phase transformation potential in CuO/NiO (FIG. 3d). For the pre-synthesized NiO catalysts, the Ni—O mode in NiO is also rapidly attenuated with reducing potentials (FIG. 3e). Additionally, the same but less notable Raman signals of Cu₂O and NiO were detected in the absence of NO₃⁻(FIG. 3f), partly owing to the NO₃⁻RR delaying the electrochemical phase transformation. The XPS further confirmed the stabilization of CuO/NiO surface phase compositions. Cu and Ni valence state variations along depth etching were further investigated (FIG. 12a to 12c). It is to be noted that, as the etching depth increases, the peaks of Cu²⁺ and Ni³⁺ disappear from 0 to 20 nm, while those of Cu⁺ and Ni²⁺ still persist in 100 nm, implying a small amount of Cu²⁺ and Ni³⁺ majorly located at the subsurface region. According to those results, we deduce that the emergence of Cu²⁺ and Ni³⁺ at the topmost sample surface can be attributed to the undesirable oxidation in ambient air. Thus, Cu₂O and NiO are the remaining phases after 2 h electrolysis. Moreover, SEM images of CuO/NiO catalysts after 2 h electrochemical reduction suggest that the whole morphology is maintained, while the surface shows relatively sharp vertices and edges (FIG. 3g). XRD patterns were as well conducted to further reveal the transformation and stabilization of the Cu₂O/NiO heterostructure electrocatalyst (FIG. 13). The HRTEM image provides the lattice spacings of 0.213, 0.241, and 0.209 nm, which can be indexed to the (200) plane of Cu₂O and (111) and (200) plane of NiO (FIG. 4h and FIG. 4i). To check out the possible changes of Cu⁺, in situ Raman spectroscopy was utilized to monitor the self-reconstruction behavior of CuO/NiO catalysts during 10 cycles of electrolysis (2 h per cycle). Impressively, as shown in FIGS. 11a to 11j, the peaks of Cu⁺ shows no obvious change during the reaction of 10 cycles of NO₃—RR, confirming its excellent stability. Thus, based on the above characterization results, in situ electrochemical self-reconstruction phenomenon of CuO/NiO heterostructure was confirmed, and Cu₂O/NiO could be the main active phase for NO₃⁻— to —NH₃ conversion.

In order to unveil the contributions of as-synthesized heterostructures towards the superior NO₃⁻RR characteristics, an H-type electrolytic cell was constructed to measure and compare the performance of heterostructures in 200 ppm NaNO₃—N and 0.5 M Na₂SO₄ electrolyte (FIG. 4a). Referring to FIG. 4a, the counter electrode or platinum foil is designated as 1, the working electrode or the CuO/NiO@NF is designated as 2, the reference electrode or Ag/AgCl electrode is designated as 3. Both the added NO₃ and generated NH₃ are located at the cathode. Unless otherwise specified, all potentials were corrected versus the reversible hydrogen electrode. The original linear sweep voltammetry (LSV) curves were conducted to assess the intrinsic activity of the as-synthesized CuO/NiO, CuO, and NiO. The LSV curves in Na₂SO₄ electrolytes with and without NaNO₃ were investigated. The obviously enhanced current density for all the catalysts was attributed by the reduction of NO₃⁻ ions (FIGS. 14a to 14c). Interestingly, two distinct reduction peaks (peak R1 and R2) were observed in the curves of CuO/NiO and CuO. According to the previous study, the peak 1 (P₁) was allocated to the reduction of adsorbed NO₃⁻(*NO₃—) to adsorbed NO₂⁻(*NO₂—) corresponding to a two-electron transfer reaction (equation 1), while the peak 2 (P₂) was assigned to the *NO₂⁻ reduction into *NH₃ (equation 2) corresponding to a six-electron transfer reaction.^[50,51] After the second reduction peak (P₂), the current density seems to increase dramatically with the decrease of potential, which can be attributed to the hydrogen evolution reaction (HER).

*NO₃⁻+H₂O+2e⁻→*NO₂⁻+2OH⁻  (1)

*NO₂⁻+5H₂O+6e⁻→*NH₃+70H⁻  (2)

Further comparing the LSV curves of CuO/NiO, CuO, and NiO (FIG. 4b), CuO/NiO, CuO appear to share a similar behavior toward NO₃⁻RR at the initial stage of the reaction, indicating the important role of CuO in P₁ step. In the P2 region, CuO/NiO displays a positive potential shift compared to CuO. Moreover, CuO/NiO yields the maximum current density of −33.5 mA cm⁻² at a potential of −0.2 V vs. RHE, which is 1.5 times of that of CuO and 7.1 times of that of NiO. CuO/NiO possesses a relatively larger current density than CuO and NiO, indicating the higher intrinsic activity of CuO/NiO. However, as the applied potential becomes negative, lower than −0.2 V and −0.45 V, respectively, the current density of CuO/NiO and CuO decrease. The reason is that the surfaces are blocked by the adsorbed reduced intermediates and products until the surface is renewed. However, the current density of CuO/NiO exhibits a slight and steady decrease compared to that of CuO, proving the combined NiO phases promptly eliminate the influence of toxic adsorbates. Moreover, the CuO/NiO exhibits a larger diffusion-limited maximum total current density (j_total) of −33.5 mA cm⁻² than that of CuO. Thus, there is a synergy between the Ni-based and Cu-based phases in CuO/NiO for the highest NO₃⁻RR activity, especially ranging from −0.15 V to −0.4 V. This way, we chose −0.2 V as the operation voltage for further characterization of the NO₃⁻RR products. A series of controlled-potential measurements were carried out to obtain FE of NH₃, in turn achieving the partial current densities of NH₃ (FIG. 15a to 15b. Benefiting from the great FE of NH₃ of CuO/NiO, it could reach a partial current density of NH₃ of 60.61 mA cm⁻² at −0.6 V much higher than that of CuO and NiO. To accurately derive NO₃⁻—NH₃ conversion activities of catalysts, we normalize the partial current densities of NH₃ by the electrochemical active surface area (ECSA). ECSA is measured using the equation of ECSA=C_dl/C_ref., where C_dl is the double-layer capacity, and C_ref is assumed as 40 μF cm⁻² as a moderate value for the specific capacitance of a flat surface. The ECSA normalized partial current densities of NH₃ are shown in FIGS. 16a to 16e, which suggests the high intrinsic NO₃RR performance of CuO/NiO. With the prolonged electrolysis time, the concentration of NO₃⁻—N kept decreasing from 200 ppm (−0.2 V, pH 7.20 at 0 min) to 21.1 ppm (−0.2 V, pH 13.50 at 120 min), whereas the concentration of NH₄⁺-N continuously increasing to 158.3 ppm indicating a nitrate conversion efficiency of 89.5%, a NH₃ selectivity of 88.5% and a high NH₃ yield rate of 2.1 mol m⁻¹ m⁻² (FIG. 4c and FIGS. 17a to 17f). In particular, the concentration of NO₂⁻ products remains low during the electrolysis, which further confirms the high selectivity to NH₃. As depicted in FIG. 21, the catalytic current obviously decreased during the first 1 hour of electrolysis and then showed slight degradation with the prolonged reaction time. The special chronoamperometric curve (i-t) should be attributed to the slowdown of the reaction rate. In the NO₃⁻RR process, the concentration of NO₃⁻ keeps decreasing with the accumulated NH₃ in the electrolyte (FIG. 4c), which impedes this catalytic reaction by constraining the kinetic mass transfer of NO₃⁻, resulting in the worse catalytic activity with the prolonging of electrolysis time. As depicted in FIG. 25, the NO₃⁻—N removal rate slows down with the increased electrolysis time. The higher removal rate in the first hour means the quick reduction of nitrate. In that way, the catalytic current exhibits a sharp degradation because the current is related to the nitrate concentration. From 0 V to −0.5 V vs. RHE, the yield rate of NH₃ gradually increases while the Faradaic efficiency curve shows a volcanic shape with a maximum value of 95.6% at −0.2 V (FIG. 4d). Here, the slight deterioration of Faradaic efficiency at more negative potentials is caused by the competing HER. The Faradaic efficiency and ammonia yield show no obvious descend tendency after ten cycles of 2 h electrocatalytic NO₃⁻RR reaction, confirming its excellent stability (FIG. 4e). To trace the source of the NH₃, isotope labeling experiments were conducted by adding ¹⁵NO₃⁻as the N source. 1H nuclear magnetic resonance (NMR) spectroscopy exhibited doublet patterns, and the coupling constant of 73 Hz corresponded to the signal of ¹⁵NH₄⁺. There is no signal of ¹⁴NH₄⁺ confirming strong evidence that the obtained NH₃ originated from the electrolysis of ¹⁵NO₃⁻ instead of any other ammonia pollution (FIG. 22a). Furthermore, adding ¹⁴NO₃⁻/¹⁵NO₃⁻in a ratio of 1 to 1, the unchanged ratio of ¹⁴NH_4+/¹⁵NH₄⁺ after the electrocatalytic reduction once again identifies added NO₃⁻is the only source of nitrogen during the reduction process here (FIG. 4f). The concentration of NH₃ calculated by the ¹H NMR and colorimetric methods are nearly equal, evidently proving the accuracy of the two quantitative methods (FIG. 4g and FIG. 22a). The achieved NO₃⁻RR activity over CuO/NiO is also superior to most reported catalysts (FIG. 4h).

To further illustrate the catalytic activity of NO₃⁻RR on CuO/NiO and Cu₂O/NiO, the rate-determining step (RDS) energy barriers on those two catalysts are calculated in FIG. 5a. The results indicate that the RDS energy barrier of nitrate conversion on CuO/NiO is higher than that on Cu₂O/NiO, proving a more favorable nitrate reduction reaction on Cu₂O/NiO. Therefore, the above calculation evidences that Cu₂O/NiO is the active phase for NO₃⁻RR. DFT calculations were performed to reveal the electronic structures of Cu₂O/NiO heterostructures and the origin of the enhanced NO₃⁻RR performance. The interface model bridging the Cu₂O (200) and NiO (200) surface was built according to the observed crystalline lattice planes in HRTEM images of the catalysts after electrolysis (FIG. 3h). As the charge density difference for Cu₂O/NiO depicted in FIG. 5b, abundant charge (the yellow area) accumulated at the interface between Cu₂O and NiO indicates that most of the charge transfer in Cu₂O/NiO occurs at the interface. That is, electron-coupled interactions between Cu₂O and NiO phases similar to the above CuO/NiO heterostructure can change the electron distribution at the interface and optimize the electronic structure of Cu₂O/NiO, thus enhancing the NO₃⁻RR catalytic activity. Additionally, the charge density difference between before and after adsorption of NO₃⁻on the interface of Cu₂O/NiO (*Cu₂O/NiO), on the surface of Cu₂O phase (*Cu₂O) and NiO phase (*NiO) of Cu₂O/NiO were evaluated (FIGS. 5c-e), respectively. The calculation results exhibited more obvious charge transfer from the oxygen atoms of NO₃⁻to Cu sites and Ni sites at the interface of Cu₂O/NiO, providing powerful clues to a stronger interaction between *NO₃ and the interface region of Cu₂O/NiO. Furthermore, Gibbs free energy of NO₃⁻RR process was calculated by a series of deoxidation steps (*NO₃⁻→*NO₂ →*NO→*N) and several hydrogenation steps (*N→*NH→*NH₂ →*NH₃) based on previous study.^[55] As depicted in FIG. 5f and FIGS. 23a to 23b, the rate-determining step (RDS) of NO₃⁻RR for *Cu₂O/NiO and *NiO was the reduction of NO* to N* with the differences of Gibbs free energies (ΔG) of 0.16 and 0.35 eV, respectively. However, the RDS of NO₃⁻RR should be hydrogenation of *N to *NH for *Cu₂O with AG of 0.30 eV (FIG. 5e), thus the RDS of NO₃⁻RR for *Cu₂O/NiO is much easier to occur. Similar to the results of charge density difference diagrams, ΔG of the adsorption of *NO₃ on the interface of Cu₂O/NiO (ΔG=−2.04 eV) is higher than the one of *NiO (ΔG=−1.45 eV) and *Cu₂O (ΔG=−1.72 eV), which displays that the interface of Cu₂O/NiO is the optimal active sites for the absorption of NO₃⁻. The LSR in the multi-step reactions means the binding of reactants and products are scaled linearly; in other words, strong adsorption of reaction intermediates will likely lead to the strong adsorption of products, therefore impeding the reaction considerably. The LSR says that in an 8-electron transfer NO₃⁻RR reaction, the adsorption energies of reactants NO₃⁻ and products NH₃ are linearly linked. That is, strong adsorption of NO₃⁻ resulted in strong NH₃ adsorption, thus decelerating NH₃ release and lead to surface poisoning. Whereas too weak adsorption results in a large energy barrier of NO₃⁻ adsorption and thus slows down the reaction. Thus, the PDOS (FIG. 5g) profiles for NO₃⁻ adsorption and NH₃ desorption were studied. The energy difference (ΔE) between 3d band center (ε_d) of the Cu site and Ni site at the interface of Cu₂O/NiO (*Cu₂O/NiO) and NO₃⁻2p band center (ε_p) was calculated to be 2.93 eV. As a comparison, ΔE (4.27 eV) between ε_d of *Cu₂O/NiO and ε_p of NH₃ is much larger than that of absorbed NO₃⁻. This finding indicates that Cu₂O/NiO possesses a strong *Cu₂O/NiO 3d-NO₃⁻2p orbital hybridization and a weak orbital hybridization between *Cu₂O/NiO 3d-NH₃ 2p. Based on previous studies, the strong orbital hybridization between *Cu₂O/NiO and NO₃⁻ indicates a higher Cu—O and Ni—O covalency, and the weak orbital hybridization means NH₃ product can easily release from the active sites. Thus, the appropriate orbital hybridization with NO₃⁻ and NH₃ endow Cu₂O/NiO with superior NO₃⁻RR performance. To further understand the synergistic role of Cu₂O and NiO for NO₃⁻RR, the orbital hybridization of NO₃⁻ and NH₃ with Cu₂O and NiO were compared (FIGS. 24a to 24b). The calculated results showed that the lower ΔE between Ed of Cu₂O and ε_p of NO₃⁻ and the larger ΔE for NiO 3d-NH₃ 2p orbital hybridization, suggesting Cu₂O considerably enhances NO₃⁻ adsorption, while NiO could promote the desorption of NH₃. This calculation results match well with the LSV experiment results (FIG. 4b). Therefore, taking together the computational and experimental results, the reconstructed active phase Cu₂O/NiO modulates the local electronic structures and possesses the tandem effect of Cu₂O and NiO, thus leading to the efficient nitrate to ammonia electrosynthesis at low overpotentials.

CONCLUSION

The above has described the synthesis and stabilization of the intermediate phase of transition metals oxide heterostructures via electrochemical redox activation-induced phase reconstruction strategy for efficient ammonia electrosynthesis. The experimental and theoretical calculations reveal that the intermediate coupling phases induce electron redistribution and dramatically lower the overpotential in the rate-determining step. Moreover, the calculations of orbital hybridization reveal that NO₃⁻ adsorption is considerably enhanced by Cu₂O, while NiO could promote the desorption of NH₃ and thus promptly renew the surface. Those two cooperative catalytic sites enable high efficiency and selectivity for NH₃ electrosynthesis. When the applied potential is −0.2 V, the reconstructed and stabilized Cu₂O/NiO catalysts exhibit an excellent NH₃ FE of 95.6%, a super-high NH₃ selectivity of 88.5%, and the ammonia-evolving rate of 2.1 mol m⁻¹ m⁻² outperforming most of the reported catalysts at the same electrocatalytic conditions. Importantly, this study provides a systematic strategy to analyze and stabilize the active intermediate phase of CuO/NiO catalysts for efficient ammonia electrosynthesis, contributing to an in-depth understanding of the accurate electrocatalytic mechanisms.

It should be understood that certain features of the invention, which are, for clarity, described in the content of separate embodiments, may be provided in combination in a single embodiment. Conversely, various features of the invention which are, for brevity, described in the content of a single embodiment, may be provided separately or in any appropriate sub-combinations. It is to be noted that certain features of the embodiments are illustrated by way of non-limiting examples.

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Claims

1. A method of preparing a catalyst for use in synthesizing ammonia from nitrate, comprising the steps of: providing a nickel foam material,removing an oxide layer on the nickel foam material,subjecting the nickel foam material to cooper sulphate (Cu2SO4) solution for a predetermined period of time for forming a copper (Cu) dopped nickel foam electrode, andsubjecting the Cu dopped nickel foam electrode to an annealing treatment at a predetermined temperature for a predetermined period of annealing time in atmospheric air thus forming a catalyst of copper oxide-nickel oxide nickel foam (CuO/NiO@NF) with a heterostructure.

2. A method as claimed in claim 1, wherein the step of removing the nickel oxide (NiOX) layer on the nickel foam material is achieved by sonicating the nickel foam material in an acid solution for a predetermined period of time.

3. A method as claimed in claim 2 wherein the acid solution is hydrochloric acid (HCl) solution, sulfuric acid (H2SO4) solution, or nitric acid (HNO3) solution, the concentration of the acid is 0.5-1 M, and the period of time of the sonicating is 5-10 minutes.

4. A method as claimed in claim 2, comprising a step, after the step of removing the NiOX, of removal of residuals from the acid solution, wherein the residual removal is achieved by rinsing with acetone or water.

5. A method as claimed in claim 1, wherein the Cu2SO4 solution has a concentration of 40-60 mM but preferably 50 mM and the period of time is 30-40 minutes but preferably 30 minutes.

6. A method as claimed in claim 1, wherein the predetermined temperature of the annealing temperature ranges from 300° C. to 600° C., or 300° C., 400° C., 500° C. or 600° C., and the predetermined annealing time is 30-60 mins.

7. A method as claimed in claim 6, wherein the predetermined annealing temperature is 400° C.

8. A method of preparing an electrolytic cell for use in synthesizing of ammonia, comprising a step of: providing the catalyst CuO/NiO@NF made according to a method as claimed in claim 1, a platinum foil, and a saturated Ag/AgCl electrode, wherein the catalyst CuO/NiO@NF, the platinum foil, and the saturated Ag/AgCl electrode act as a working electrode, a counter electrode, and a reference electrode, respectively in the electrolytic cell.

9. A method as claimed in claim 8, wherein the catalyst CuO/NiO@NF is in the formation of a sheet and has a dimension of 1×1 cm.

10. A method as claimed in claim 8, comprising a step of providing a sodium sulfate (Na2SO4) solution as a catholyte and an anolyte of the electrolytic cell for use in a cathode chamber and an anode chamber of the electrolytic cell.

11. A method as claimed in claim 10, wherein the concentration of the (Na2SO4) solution is 0.5-1 M.

12. A method as claimed in claim 10, comprising a step of providing a proton exchange membrane between the cathode chamber and an anode chamber of the electrolytic cell.

13. A method as claimed in claim 12, wherein the proton exchange membrane is a Nafion® 117 membrane.

14. A method of synthesis of ammonia from nitrate, comprising the steps of: providing an electrolytic cell made according to claim 8,prior to using the electrolytic cell, purging all solutions used in preparing the electrolytic cell preparation and the electrode preparation including the Na2SO4 solution with pure nitrogen (N2) gas for a predetermined period of purging time.

15. A method as claimed claim 14, wherein the purging time is 10 mins.

16. A method as claimed in claim 15, comprising a step of adding a source of nitrate (NO3) to the cathode chamber and allowing reaction to take place for generating ammonia (NH3).

17. A method as claimed in claim 16, comprising a step of collecting the generated ammonia (NH3) from the cathode chamber.