Electrochemical synthesis of nitric acid from air and ammonia through waste utilization

ABSTRACT Commercial nitric acid (HNO3) and ammonia (NH3) are mostly produced through the Ostwald process and the Haber-Bosch process, respectively. However, high energy demand and enormous greenhouse gas accompy these processes. The development of economical and green ways to synthesize HNO3 and NH3 is highly desirable for solving the global energy and environmental crisis. Here, we present two energy-efficient and environmentally friendly strategies to synthesize HNO3 and NH3 at distributed sources, including the electrocatalytic oxidation of N2 in air to HNO3 and the electrocatalytic reduction of residual \documentclass[12pt]{minimal} \usepackage{amsmath} \usepackage{wasysym} \usepackage{amsfonts} \usepackage{amssymb} \usepackage{amsbsy} \usepackage{upgreek} \usepackage{mathrsfs} \setlength{\oddsidemargin}{-69pt} \begin{document} }{}${\rm NO_{3}^{-}}$\end{document} contamination in water to NH3. The isotope-labeling studies combined with theoretical calculation reveal the reaction path of the two proposed strategies, confirming the origin of the electrochemical products. Importantly, the electrooxidation-generated \documentclass[12pt]{minimal} \usepackage{amsmath} \usepackage{wasysym} \usepackage{amsfonts} \usepackage{amssymb} \usepackage{amsbsy} \usepackage{upgreek} \usepackage{mathrsfs} \setlength{\oddsidemargin}{-69pt} \begin{document} }{}${\rm NO_{3}^{-}}$\end{document} ions may also serve as reactants for the electroreduction synthesis of NH3 in the future. Our work may open avenues for energy-efficient and green production of HNO3 and NH3 at distributed sources.


and NH 3 under benign conditions.
Herein, we present two electrochemical strategies. Strategy I is the electrocatalytic oxidation of N 2 to HNO 3 by using air as the nitrogen source. Strategy II is the electrochemical reduction preparation of NH 3(aq) from residual nitrate ion (NO − 3 ) contamination in water [25,26]: We found that N 2 is electro-oxidized into HNO 3 over platinum foil with ∼1.23% Faradaic efficiency at +2.19 V vs. RHE (the reversible hydrogen electrode) and the waste of NO − 3 is electro-reduced with approximately 33.6% NH 3(aq) selectivity at −0.65 V vs. RHE over Co 3 O 4 nanorod arrays. Our results demonstrate how the electrochemical methods of Strategy I and Strategy II produce HNO 3 and NH 3 at distributed sources. These findings provide a new avenue for producing the reactive nitrogen species in an 'economic' and 'clean' way, especially once the electrocatalysis reaction is driven by renewable energy [27].

RESULTS AND DISCUSSION
An H-type cell divided by a proton-exchange membrane was used for the electrocatalytic tests of Strategy I (Fig. 1a). To explore the catalytic behavior of N 2 electrooxidation (Strategy I), air was bubbled onto the anode. H 2 O in electrolyte (0.3 M K 2 SO 4 ) and N 2 in air combine with platinum foil as an electrocatalyst to form HNO 3 . We tested the linear sweep voltammetry curves of platinum foil in Ar-and air-saturated electrolyte under ambient conditions (Fig. 1b). All potentials in this work were recorded and converted to the RHE scale. As the potential moves above +2.13 V, the current density is distinguishably enhanced under air-saturated electrolyte, revealing that N 2 in air can be catalysed into oxidative products. The produced NO − 3 and NO − 2 are quantified based on the standard method [28] by using ultraviolet-visible (UV-Vis) spectrophotometry ( Supplementary Fig. S1, available as Supplementary Data at NSR online). Anion chromatography was also adopted to confirm the accuracy of UV-Vis spectrophotometry for detecting the yield of NO − 3 (Supplementary Table S1, available as

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Supplementary Data at NSR online). The effect of the anode potentials on the yields of oxidative products and the corresponding Faradaic efficiencies were investigated (Fig. 1c). Although the gap in the current density between Ar and air condition keeps enlarging with the increase in potential, the highest Faradaic efficiency of 1.23% for NO − 3 was obtained at +2.19 V. The yields of NO − 3 and NO − 2 at +2.19 V achieve 0.06 and 0.0004 μmol h −1 cm −2 , respectively, and remained almost unchanged with the increase in potential ( Fig. 1c and Supplementary Fig. S2, available as Supplementary Data at NSR online). So, the optimal operating potential for N 2 electrooxidation over platinum electrocatalyst in this work was +2.19 V. Furthermore, some control experiments, including Ar-saturated electrolyte with a potentiostatic test (+2.19 V) and air-saturated electrolyte without external potential, were performed. Undetected oxidative products in both cases further confirmed the electrocatalytic oxidation of N 2 in air to NO − 3 and NO − 2 as designed in Strategy I (Supplementary Table S2, available as Supplementary Data at NSR online). For practical application, the durability of the catalyst is crucial. After consecutive recycling tests, the catalytic performance showed no obvious decline, demonstrating the high stability of platinum foil toward the electrooxidation of N 2 (Fig. 1d). As shown in Fig. 1e, Strategy I goes through multiple processes. We calculated the free energy of N 2 electrooxidation over the Pt (200) plane based on the X-ray diffraction (XRD) result ( Supplementary Fig. S3, available as Supplementary Data at NSR online). First, the N 2 molecule was chemically absorbed by the platinum to form N 2 * with a total energy change of −0.097 eV, indicating that this reaction can take place spontaneously. N 2 * reacted with OH − to produce N 2 OH * and then N 2 OH * was dehydrogenated to N 2 O * or continued to react with OH − to produce N 2 O 2 H 2 * with a larger reaction energy. Both N 2 O * and N 2 O 2 H 2 * will evolve into NO * with the intermediate of N 2 O 2 H * and NOH * . Note that NO * could be directly desorbed from the catalyst surface and then oxidized into HNO 3 and HNO 2 in solution (Equation (2)). Meanwhile, NO * can be oxidized to NO 2 * with the intermediate of NO 2 H * . With further increase in potential, NO 2 * desorbed from the catalyst surface. Finally, NO 2 was transformed to HNO 3 in solution through Equation (3). Based on these results, we can deduce the reaction path from N 2 in air to HNO 3 via the as-proposed Strategy I: (3) To demonstrate Strategy II, involving electroreducing residual NO − 3 contamination in water, KNO 3 was added to a cathode cell and reduced into NH 3(aq) (Fig. 2a). Presently, the main sources of nitrate for the as-proposed Strategy II are residual contamination in water, including industrial wastewater, domestic sewage, animal waste and nitrogen fertilizers. In the future, if the efficiency of Strategy I can be further improved, the electrooxidation generation of NO − 3 may also serve as a reactant for Strategy II. We constructed Co 3 O 4 nanorod arrays supported on Ti mesh as a model electrocatalyst ( Supplementary Fig. S4, available as Supplementary Data at NSR online). The linear sweep voltammetry curves of the Co 3 O 4 electrode in 0.1 M K 2 SO 4 electrolyte with and without KNO 3 were performed under ambient conditions (Fig. 2b). The current density was obviously enhanced in the presence of KNO 3 , indicating that NO − 3 in solution can be catalysed into reductive products. The yields of NH 3(aq) [28] and NO − 2 were quantified based on UV-Vis spectrophotometers ( Supplementary Fig. S5, available as Supplementary Data at NSR online). The cation chromatography method was also adopted to confirm the accuracy of UV-Vis spectrophotometry for detecting the yield of NH 3(aq) (Supplementary Table S3, available as Supplementary Data at NSR online). Following the potentiostatic test conducted at −0.65 V vs. RHE, which corresponded to 100 mA cm −2 of current density (Fig. 2b), the selectivity of different reductive products is displayed in Fig. 2c  These results demonstrate the high durability of the Co 3 O 4 electrode for NO − 3 electroreduction. Note that the electroreduction of NO − 3 was reported in previous work, but they focused on the degradation of residual NO − 3 in water into environmentally friendly products [32][33][34]. We herein propose utilization of the waste of NO − 3 , provided by environmental contaminations, to produce the high value-added NH 3(aq) via electroreduction.
To confirm the origin of the NO − 3 generated from N 2 electrooxidation, we designed an isotopiclabeling study using 15 N 2 (>99 atom% 15 N) and 14 N 2 (with the natural abundance of 0.36 atom% 15 N) as the feeding gas for Strategy I [35]. As seen in Fig. 3a, the sample using 14 N 2 as the feeding gas produced 0.44% 15 NO − 3 , while the isotopic-labeled sample showed 17.40% abundance of 15 NO − 3much higher than the natural abundance of 15 N. The concentration difference of 15 N for the isotopiclabeled sample between the 15 N 2 reactant and the 15 NO − 3 product may arise from the leaking and/or residue of air in the reactor. These results clearly confirm that the N element of the generated NO − 3 via Strategy I came from N 2 . We also performed an isotopic-labeling study using K 15 NO 3 (20.3 atom% 15 N) and K 14 NO 3 (0.36 atom% 15 N) as the reac-tants to explore the origin of the NH 3(aq) . It can be seen that the isotopic-labeled sample exhibited 18.98% 15 NH 3(aq) but the sample using K 14 NO 3 without an isotopic label as a reference showed only 0.36% 15 NH 3(aq) (Fig. 3b). These results demonstrate that the N element of the formed NH 3(aq) via Strategy II originated from the NO − 3 species.

CONCLUSIONS
In summary, we present two energy-efficient and environmentally friendly strategies to prepare HNO 3 and NH 3 at distributed sources. Strategy I is the electrocatalytic oxidation of N 2 in air to HNO 3 and platinum foil is adopted as the model catalyst, showing a generation rate of 0.06 μmol h −1 cm −2 for HNO 3 at +2.19 V. Strategy II is the electrocatalytic reduction of residual NO − 3 contamination in water to NH 3(aq) and Co 3 O 4 nanorod arrays supported on a Ti mesh, exhibiting a selectivity of 33.6% for NH 3(aq) at −0.65 V. Combined with the theoretical calculation, the reaction path from N 2 in air to HNO 3 via the as-proposed Strategy I was deduced. The isotopelabeling studies using 15 N 2 and K 15 NO 3 confirmed the origin of the electrochemical products. Moreover, the platinum electrode and Co 3 O 4 electrode showed high durability for the electrocatalytic oxidation of N 2 and the electrocatalytic reduction of NO − 3 , respectively. In the future, NO − 3 from Strategy I may also serve as a reactant for the electroreduction synthesis of NH 3 . Our results presented here provide new avenues for energy-efficient and green production of HNO 3 and NH 3 at distributed sources.

reduction via Strategy II
In a typical process, 2 mmol Co(NO 3 ) 2 ·6H 2 O, 10 mmol urea and 8 mmol NH 4 F were dissolved in 36 mL distilled water under stirring for 5 min. The aqueous solution was moved to a 50-mL Teflon-lined autoclave and then a piece of Ti mesh (1 × 3 cm 2 ) was immersed in the above solution. The autoclave was sealed and heated at 120 • C for 9 h, followed by cooling down to ambient temperature. The sample was washed using distilled water and ethanol six times and then dried in a vacuum oven overnight. The dried sample was annealed at 300 • C for 2 h in air to acquire the final product of Co 3 O 4 nanorod arrays supported on a Ti mesh.

Characterization
The scanning electron microscopy images were acquired from a Hitachi S-4800 scanning electron microscope. Transmission electron microscopy and high-resolution transmission electron microscopy images were taken using a JEOL-2100F system. The XRD was measured using a Bruker D8 Focus Diffraction System with a Cu Kα source (λ = 0.154178 nm). X-ray photoelectron spectrum analysis was recorded via a PHI 5000 Versaprobe system using monochromatic Al Kα radiation. All binding energies were revised according to the C 1-s peak at 284.8 eV. The ultraviolet-visible (UV-Vis) absorbance spectra were measured on a Beijing Purkinje General T6 new century spectrophotometer. Anion chromatography was performed on an ICS-1100, Thermo. Cation chromatography was conducted on an ICS-900, Thermo. The concentration of 15 N isotope labeling was established by isotopic mass spectrometry (MAT-271). The pH values of the electrolytes were determined using a pH-meter (LE438 pH electrode, Mettler Toledo, USA).

Electrochemical measurements
Electrochemical measurements were conducted using an electrochemical workstation (CHI 660D, Chenhua, Shanghai). A typical H-type electrolytic cell divided by a proton-exchange membrane (Nafion 117) was used. Except for special instructions, all potentials were recorded against the RHE. The potentials against the saturated calomel electrode (SCE) were translated to those against the RHE using the following equation: E (vs. RHE) = E (vs. SCE) + 0.2415 + 0.059 × pH. All the polarization curves were the steady lines after many cycles and the current density was normalized to the geometric surface area.
For the electrooxidation of N 2 to HNO 3 via Strategy I, Pt plates (1 × 1 cm 2 ) were used as both the working electrode and the counter electrode, the reference electrode was SCE and 0.3 M K 2 SO 4 solution (70 mL) was adopted as the electrolyte. Air (99.99% purity) was bubbled into the anodic compartment with a flow rate of 10 mL min −1 in the whole electrochemical process. The linear sweep voltammetry was performed at a rate of 10 mV s −1 and the potentiostatic test was tested at a different anodic voltage for 20 h with the electrolyte agitated at a stirring rate of ∼350 rpm. An absorption flask containing 5 mL K 2 SO 4 solution (0.3 M) was connected to the gas outlet of an anodic half cell to avoid the loss of products due to air bubbling. After the electrooxidation measurements, the components of the mixed solutions in the anodic compartment and absorption flask were both analysed. For comparison, the electrooxidation measurements were also measured with all the testing conditions consistent with aforementioned N 2 oxidation except that air was replaced by Ar or the external potential was removed.
The electroreduction of NO − 3 to NH 3 via Strategy II was carried out in a three-electrode configuration with as-prepared Co 3 O 4 (1 × 1 cm 2 ) electrode, SCE and platinum foil as working electrode, reference electrode and counter electrode, respectively; 0.1 M K 2 SO 4 solution (80 mL) was used as the electrolyte and evenly distributed to the cathode and anode compartment. KNO 3 (100 g L −1 ) was added into the cathode compartment as a reactant. Prior to the NO − 3 electroreduction test, the cathode electrolyte was purged with Ar (99.99% purity) for 30 min. The linear sweep voltammetry was performed at a rate of 20 mV s −1 and the potentiostatic test was conducted at −0.65 V for 3 h at a stirring rate of ∼350 rpm. For comparison, the electroreduction measurements were also conducted with all the testing conditions consistent with the aforementioned NO − 3 reduction except that the Co 3 O 4 cathode was replaced by the Ti mesh or the Co 3 O 4 cathode was immersed in electrolyte (0.1 M K 2 SO 4 ) without the addition of KNO 3 .

Ion-concentration detection methods
The electrolytes pre and post test were first diluted to appropriate concentration and then tested using a UV-Vis spectrophotometer to quantify the concentration. The concentrations of nitrate-N, nitrite-N and ammonia-N were estimated by UV-Vis spectrophotometry according to the standard method. The specific approaches are as follows.

Determination of nitrate-N
First, a certain amount of electrolyte was taken out of the electrolytic cell and diluted to 5 mL to the detection range. Then, 0.1 mL 1 M HCl and 0.01 mL 0.8 wt% sulfamic acid solution were added to the aforementioned solution. The absorption spectrum was tested using an ultraviolet-visible spectrophotometer and the absorption intensities at wavelengths of 220 and 275 nm were recorded. The final absorbance value was calculated using the equation: A = A 220nm -2A 275nm . The concentration-absorbance curve was made using a series of standard potassium nitrate solutions and the potassium nitrate crystal was dried at 105-110 • C for 2 h in advance.

Determination of nitrite-N
A mixture of p-aminobenzenesulfonamide (4 g), N-(1-Naphthyl)ethylenediamine dihydrochloride (0.2 g), ultrapure water (50 mL) and phosphoric acid (10 mL, ρ = 1.70 g/mL) was used as a color reagent. A certain amount of electrolyte was taken from the electrolytic cell and diluted to 5 mL to the detection range. Next, 0.1 mL color reagent was added into the aforementioned 5-mL solution and mixed to uniformity, and the absorption intensity at a wavelength of 540 nm was recorded after sitting for 20 min. The concentration-absorbance curve was calibrated using a series of standard sodium nitrite solutions.

Determination of ammonia-N
Ammonia-N was determined using Nessler's reagent as the color reagent. First, a certain amount of electrolyte was taken from the electrolytic cell and diluted to 5 mL to the detection range. Next, 0.1 mL potassium sodium tartrate solution (ρ = 500 g/L) was added and mixed thoroughly, then 0.1 mL Nessler's reagent was put into the solution. The absorption intensity at a wavelength of 420 nm was recorded after sitting for 20 min. The concentration-absorbance curve was made using a series of standard ammonium chloride solutions and the ammonium chloride crystal was dried at 105 • C for 2 h in advance.
The All the results obtained using UV-Vis spectrophotometry were compared with those of ion chromatography (Supplementary Tables S1 and S3, available as Supplementary Data at NSR online). The pre-treatment of the cation chromatography measurement of ammonia-N after the electroreduction test was as follows. First, the electrolyte was put into a round flask and then the pH was adjusted to ∼8 by the addition of NaOH. The solution was then heated to ∼100 • C and went through condensation until the 90% electrolyte was distilled. The distilled solution was collected using 35 mL ultrapure water and used for the cation chromatography measurement.

Calculation of the yield, selectivity and Faradaic efficiency
For the N 2 electrooxidation experiments via Strategy I, the yield of NO − 3 and NO − 2 was calculated using Equation (4) and Equation (5), respectively:

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and the total charge passed through the electrode using Equation (6): where c NO3 − is the concentration of NO x − , V is the volume of the electrolyte in the anode compartment, M NO3 − is the molar mass of NO − 3 , M NO2 − is the molar mass of NO − 2 , t is the electrolysis time, S is the geometric area of the Pt plate, F is the Faradaic constant (96 485 C mol −1 ) and Q is the total charge passing the electrode. (Note that NO and NO 2 gas consumed two and four electrons, respectively. We prudently chose NO gas as the product to calculate the Faradaic efficiency.) For the NO − 3 electroreduction experiments via Strategy II, the yield of NH 3(aq) was calculated using Equation (7): The conversion of NO − 3 was obtained from Equation (8): The selectivity of product was acquired by Equation (9): where c NH3 is the concentration of NH 3(aq) , c NO3 − is the concentration difference of NO − 3 before and after electrolysis, c 0 is the initial concentration of NO − 3 , and c is the concentration of products, including NH 3(aq) , NO − 2 .

N isotope-labeling experiment
The isotope-labeling reactants of 15  For isotope labeling the N 2 electrooxidation, we carried out the batch experiments using 15 N 2 as the feeding gas for five successive times and collected all the electrolytes together after electrolysis. In order to increase the concentration of 15 NO − 3 for isotopic mass spectrometry, the collected electrolytes were first alkalified to pH ∼7 by adding 1 M KOH solution and then concentrated by 10 times via distilling at 70 • C. For comparison, 14 N 2 with naturalabundance 15 N (0.36 atom%) was used as a feeding gas to replace 15 N 2 with other conditions being consistent.
For isotope labeling the NO − 3 electroreduction, the pH of final electrolyte was adjusted to ∼3 using 1 M HCl solution before isotopic mass spectrometry. For comparison, K 14 NO 3 with naturalabundance 15 N (0.36 atom%) was used as the reactant to replace K 15 NO 3 with the other conditions being consistent.

Theoretical simulation
Although the XRD pattern of platinum after the electrooxidation process (not shown here) showed no obvious new species, the high oxidation state of PtOH or PtO x might exist in an amorphous state. Considering the complexity of the electrochemical oxidation process, herein, we only adopted Pt as a model for simulation. The theoretical calculations were conducted using density functional theory with the Perdew-Burke-Ernzerbof form of generalized gradient approximation functional [36]. The plane wave energy was cut off at 400 eV. The Vienna ab initio simulation package was used [37,38]. The Fermi scheme was used for electron occupancy with an energy smearing of 0.1 eV. The first Brillouin zone was adopted in the Monkhorst−Pack grid [39]. The 3 × 3 × 1 k-point mesh was taken for the surface calculation. The energy of 1.0 × 10 −6 eV atom −1 and force of 0.01 eVÅ −1 were set as the convergence criterion for geometry optimization. The spin polarization was considered in all calculations. To accurately describe the van der Waals (vdW) interaction, the non-local van der Waals density functional (vdW-DF) was employed in our work [40,41].
The Pt (200) surface was obtained by cutting the Pt bulk along the {200} direction based on the XRD result ( Supplementary Fig. S3, available as Supplementary Data at NSR online). The thickness of the surface slab was chosen to be a three-layer slab. In all structural optimization calculations, the bottom atoms were frozen, while the other atoms were allowed to relax. A vacuum layer as large as 15Å was used along the c direction normal to the surface to avoid periodic interactions. For the nitrogen oxidation reaction, the reaction energies of the elementary reactions were employed to estimate the activity of the catalyst.
The free-energy changes of the nitrogen oxidation were calculated to show the reaction trend. The change in free energy ( G) of the per reaction step