Enhanced Nitrate Production via Electrocatalytic Nitric Oxide Oxidation Reaction over MnO₂ with Different Crystal Facets

Abstract

The synthesis of nitrate (NO₃⁻) via electrocatalytic nitric oxide oxidation reaction (NOOR) is a green and efficient strategy for nitrogen fixation, which has great advantages over conventional nitrate synthesis. Notably, it also presents a promising solution for the remediation of NO pollutants. In this study, the structure–performance correlations of α-, β-, and δ-MnO₂ catalysts were investigated. These three polymorphs of MnO₂ exhibited disparate surface chemistries, NO adsorption capabilities, and NOOR catalytic activities. In a 1.0 M KOH electrolyte, α-MnO₂, characterized by its large-sized (2 × 2) one-dimensional tunnel structure, demonstrated the most outstanding NOOR catalytic performance, which achieved a remarkable NO₃⁻ yield of 665.2 mg·h⁻¹·mg_cat⁻¹ at a potential of 1.9 V, along with excellent stability and durability. Furthermore, a Zn-NO system was constructed, employing α-MnO₂ as the anode and a Zn plate as the cathode. This innovative setup integrated an energy storage system with NO electrochemical capture, yielding a NO₃⁻ production rate of 265.5 mg·h⁻¹·mg_cat⁻¹. The density functional theory calculations confirm the NO adsorption performance and catalytic activity of three different crystal forms of MnO₂.

1. Introduction

Nitric oxide (NO) is a significant air pollutant in the atmosphere, and it can react with O₂ to form NO₂ [1,2,3], causing serious environmental problems, such as acid rain, photochemical smog, and ozone holes, and posing health hazards to the population [3,4,5]. Therefore, treating atmospheric NO has become a key strategy to control air pollution [6]. The most popular selective catalytic reduction (SCR) technology and the newly developed photocatalytic technology still suffer from problems like severe conditions, high cost, and limited efficiency [7,8]. Therefore, room temperature electrocatalytic denitrification is considered a sustainable route, with the advantages of mild conditions, green and low carbon, and high NO removal efficiency, addressing both energy and the environment concerns.

Recently, nitrate synthesis from NO has also been regarded as an effective approach for NO removal, in addition to NH₃ synthesis through electrocatalytic NO reduction [8,9,10]. The traditional industrial nitrate synthesis process mainly generates NH₃ from N₂ and H₂ through the Haber–Bosch process and then obtains nitrate through the oxidation of NH₃ (Ostwald process), which requires high temperatures and pressures, increases energy consumption, and emits many greenhouse gases [11]. The electrocatalytic nitrate synthesis under mild conditions is green and sustainable, which could be an alternative approach to the traditional Haber–Bosch and Ostwald processes [12]. Compared with electrochemical nitrogen oxidation reaction (NOR), with limited efficiency and nitrate yield, electrocatalytic NO oxidation reaction (NOOR) is an easier approach to converting NO pollutant into high-value nitrate because NO has lower bond energy than the high dissociation energy of N≡N bonds [5,13]. In NOR, a Sr_0.9RuO₃ catalyst shows a NO₃⁻ yield of only 17.9 μmol·h⁻¹·mg_cat⁻¹ [14], while RuO₂@TiO₂ as a NOOR catalyst achieves a NO₃⁻ yield of 95.22 mg·cm⁻²·h⁻¹.

Although various precious metals and their related compounds have excellent NOOR performance, their drawbacks, such as limited reserves and high prices, greatly restrict their further development [15]. In contrast, abundant and inexpensive transition metal oxides exhibit excellent redox activity and have unfilled d-orbitals to gain and lose electrons. Therefore, they have critical applications in various fields, such as CO catalytic oxidation and methane-reforming oxidation [16], as well as NO oxidation [17]. The reported Co₃O₄/Ti is a suitable NOOR electrocatalyst with a maximum yield of 115.45 mg·cm⁻²·h⁻¹ [17]. Other transition metal oxides have also been investigated, showing a thermal catalytic NO oxidation activity order of Mn > Cr > Co > Cu, among which MnO_x has the most outstanding performance [13,18]. The catalytic activity of MnO_x heavily relies on its chemical composition and crystal structure. Wang et al. found that the NO oxidation performances followed the order of MnO₂ > Mn₂O₃ > Mn₃O₄, of which MnO₂ had the highest NO oxidation activity [19]. The MnO₂ crystalline phase is composed of octahedral [MnO₆] units, and the diversity of [MnO₆] connection modes leads to a relatively complex crystal structure of MnO₂ [20,21]. The different phase structures result in different gas adsorption, diffusion, and catalytic reactions, significantly affecting the catalytic activity of MnO₂ [22,23]. Liang et al. found that α-MnO₂ has the best catalytic oxidation effect for CO oxidation [24]. Yang et al. analyzed the key factors affecting the catalytic activity of α-, β-, γ-, and δ-MnO₂ for toluene oxidation [25]. Chen et al. also confirmed that MnO₂@CeO₂, with a high density of active oxygen vacancies, can exhibit significant thermal catalytic activity in NO oxidation reactions [26]. However, recent studies on NO oxidation by MnO₂ have focused on thermal catalysis, and investigations into electrocatalytic NO oxidation are still insufficient.

Herein, we prepared α-, β-, and δ-MnO₂ via a hydrothermal method and applied them as electrocatalysts for room temperature NOOR. Their electrocatalytic oxidation performance for NO was investigated in a three-electrode system. The effects of the quantity of low-valent Mn on the surface of various crystal structures, oxygen vacancies, and the NO adsorption properties on the catalytic activity of NOOR were systematically investigated by XPS, HRTEM, Raman characterization, and DFT calculations. Moreover, taking advantage of the outstanding catalytic ability of α-MnO₂, a Zn-NO reaction device that integrates an energy-storage system with NOOR was demonstrated. This device not only enables the electrochemical capture and resource utilization of NO exhaust gas but also can form a zinc-NO₃⁻ battery to release electrical energy. In this way, it transforms pollutants into value-added products, showing great application prospects.

2. Results and Discussion

The typical crystalline phases and morphology of as-prepared MnO₂ were investigated to probe the effect of its crystalline structure and electrochemical NOOR activity. The TXRD results (Figure 1a) confirm the successful synthesis of α-, β-, and δ-MnO₂ catalysts, according to JCPDS 44-0141 (cryptomelane), JCPDS 24-0735 (pyrolusite), and JCPDS 80-1098 (birnessite), respectively. The sharp and narrow peaks of α- and β-MnO₂ were attributed to their better crystallinity. In contrast, δ-MnO₂ exhibited broad bands and lower crystallinity. This phenomenon was also evident in the HRTEM results and was consistent with previous research findings [21,25,27,28]. The SEM image of α-MnO₂ (Figure 1b) presented a stacking-nanoneedle structure with a thickness of about 35 nm. Clear lattice stripes were observed in the HRTEM (Figure 1e), and the crystal plane spacing of 0.152 nm, 0.239 nm, and 0.181 nm corresponded to the α-MnO₂ (521), (211), and (411) planes, respectively. β-MnO₂ exhibited coarser and shorter size nanorods compared to α-MnO₂, and HRTEM (Figure 1f), its regular lattice of dotted stripes, and the precise and uniform 0.311 nm crystallographic spacing matched well with the (110) facet, suggesting a high degree of ordering of β-MnO₂. The δ-MnO₂ manifested as polymerized flower-like nanosheets composed of thin nanosheets with a thickness of 10 nm. The HRTEM images (Figure 1g) displayed crystalline facet spacings of 0.218 nm and 0.225 nm, corresponding to the (−112) and (201) facet spacings. When compared with α- and β-MnO₂, it was observed that δ-MnO₂ was poorly crystallized and less ordered. This finding was in perfect accordance with what the XRD pattern indicated. Moreover, the TEM images (Figure S1) of the three catalysts exhibited a morphology resembling that observed in the SEM images.

The crystal structure of the crystalline phase of MnO₂ generally consists of [MnO₆] octahedral units, which form different tunneling or lamellar structures by sharing edges or corners [21,27]. α-MnO₂ and β-MnO₂ are the typical 1D tunneling structures of MnO₂, and the common 2D layered structure is mainly δ-MnO₂ (Figure S2). The structural and electronic properties of α-MnO₂, β-MnO₂, and δ-MnO₂ were further investigated using XPS and Raman spectroscopy. As shown in Figure 2a, the XPS surface scanning spectra of α- and δ-MnO₂ showed K elemental spikes, whereas β-MnO₂ did not, and this result was consistent with that in previous reports [21,25]. This suggests that the 1D tunnel or 2D layer structures of α-MnO₂ and δ-MnO₂ could accommodate the presence of K⁺, and K⁺ was vital in balancing the charges of α- and δ-MnO₂ in the crystal cells. In contrast, the β-MnO₂ tunneling structure was narrow and could not accommodate the presence of K⁺ and the crystal structure. In addition to this, α- and δ-MnO₂ had considerable lattice defects due to their intrinsic crystal defects or heterocations in tunnels, and a large number of [MnO₆] edges were exposed in the structure, which tended to form more Mn²⁺ and Mn³⁺ on their surface [29], while β-MnO₂ showed lower catalytic activity due to its regular [MnO₆] octahedral arrangement, fewer ectopic sites, and a low percentage of Mn²⁺ and Mn³⁺ on the surface. Figure 2b shows the XPS spectra of Mn 2p_3/2 for MnO₂, and the peaks at 642.7 eV, 641.7 eV, and 640.6 eV could be attributed to Mn⁴⁺, Mn³⁺, and Mn²⁺, respectively [25,30]. The content of low-valence Mn (Mn³⁺ and Mn²⁺) was in the following descending order: α-MnO₂ > δ-MnO₂ > β-MnO₂. The presence of Mn²⁺ and Mn³⁺ induced the generation of oxygen vacancies in the MnO₂ skeleton to maintain electrostatic neutrality. Thus, the content of low-valence Mn was regarded as a characteristic parameter for the catalytic activity of the MnO₂ surface [31,32]. Moreover, the bonds of Mn²⁺–O and Mn³⁺–O were weaker than that of Mn⁴⁺–O, and the presence of a large amount of low-valent manganese on the surface of α-MnO₂ led to the weakening of the Mn–O bonds, which allowed the oxygen atoms on its surface to be more easily released for participation in oxidation [25]. Recent studies had found that K⁺ in the unique tunneling structure of MnO₂ increased its electronic conductivity [33]. A low K/Mn atomic ratio was one of the reasons for the high Mn⁴⁺ ratio [21,34]. The XPS spectra of the O 1s of the samples are shown in Figure 2c, and the peaks at 533.4 eV, 531.7 eV, and 529.4 eV could be attributed to the surface hydroxyl oxygen (O_adsO-H), surface adsorbed oxygen (O_ads), and lattice oxygen (O_latt), respectively [25,29]. α-MnO₂ showed the most O_ads proportion, and O_ads reacted more in oxidation reactions due to its high shift electrophilicity than O_latt [26]. The Raman spectrum in Figure 2d demonstrated that the band shifts of α-, β-, and β-MnO₂ are at 630.8, 629.3, and 621.8 cm⁻¹. The band shift could be ascribed to the enhanced symmetric stretching vibration [31]. According to the electroneutrality principle, oxygen vacancy formation is related to Mn³⁺ and Mn²⁺, and oxygen-added substances are mainly located on surface oxygen vacancies [35,36]. Therefore, the above results suggest that α-MnO₂ might have higher activity in NO oxidation due to the many low-valent Mn cations and adsorbed oxygen.

The electrochemical properties of different crystalline phases of α-, β-, and δ-MnO₂ were investigated by electrochemical tests. The electrochemically active surface area (ECSA) indicated the effective area of the electrode involved in the catalytic reaction, which was calculated and fitted by the CV curves of the catalysts at different scanning rates (Figure S3). As shown in Figure 3a, the slopes of the fitted curves for α-MnO₂ were more significant than those for β-MnO₂ and δ-MnO₂, which suggests that α-MnO₂ had a larger ECSA and a more substantial central activity. Regarding the electrochemical impedance spectroscopy (EIS) test, the Nyquist plots of α-, β-, and δ-MnO₂ are shown in Figure 3b. The curve radius of α-MnO₂ was the smallest, indicating that it had a lower charge-transfer impedance and faster electron-transfer rate, which were favorable for the electrocatalytic performance of α-MnO₂. The electrocatalytic performances of different crystalline phases of MnO₂ for NO oxidation was tested in a gas-tight electrolytic cell. Figure 3c presents the LSV curves of α-, β-, and δ-MnO₂ for the NO oxidation reaction in a 1 M KOH solution. These curves clearly demonstrate that within the potential range of 1.0~2.4 V (vs. RHE), α-MnO₂ exhibited a higher current density, greater surface oxidation activity, and more effective engagement in the reaction. When comparing the LSV curves obtained in an Ar-saturated 1 M KOH electrolyte, Figure 3d revealed a substantial increase in the current density of α-MnO₂ in a NO atmosphere. This significant enhancement strongly implies that the NOOR reaction took place on the surface of the α-MnO₂ electrode. A comparison of the LSV curves of bare carbon paper (CP) in a NO atmosphere (Figure 3e) showed that the current density of the bare CP at the same potential was significantly lower than that of the α-MnO₂/CP electrode. This finding strongly suggests that the vast majority of the NOOR likely occurred on the surface of the α-MnO₂ catalyst rather than on the surface of the CP. To evaluate and compare the electrocatalytic performance of the three catalysts for NOOR, the α-, β-, and δ-MnO₂ electrodes were subjected to the NOOR reaction tests. These tests were carried out within the potential range of 1.6~2.0 V, with a constant NO gas flow rate maintained at 20 mL·min⁻¹. The chronoamperometry curves for the reactions are presented in Figure 3f and Figure S4. In these curves, it can be observed that the current densities of all three samples rose as the potential increased. As anticipated, at the same potential, the α-MnO₂ electrode exhibited the highest current density, while the β-MnO₂ electrode showed the lowest. This observation is in line with the trend depicted in the LSV plot of Figure 3c.

The yields of the product NO₃⁻ and the by-product NO₂⁻ were essential indicators for evaluating the NOOR performance of the three catalysts to determine the optimum operating potential. The calibrated concentration–absorption peak curves of standard NaNO₃ and NaNO₂ solutions at different concentrations and the corresponding standard curves of NO₃⁻ and NO₂⁻ are shown in Figures S5 and S6. Figure 4a,b illustrates the yields of NO₃⁻ and NO₂⁻ for α-, β-, and δ-MnO₂ catalysts at various potentials. Regarding the product NO₃⁻, the yields of β- and δ-MnO₂ initially increased and then decreased as the potential rose within a specific range, reaching their maximum values at 1.9 V. In contrast, the NO₃⁻ yield of α-MnO₂ showed a positive correlation with the applied potential, increasing steadily as the voltage varied from 1.6 V to 2.0 V. Evidently, at the same potential, the NO₃⁻ yield of α-MnO₂ surpassed those of β- and δ-MnO₂. Especially at 1.9 V and 2.0 V, the yields of α-MnO₂ were 665.2 and 703.3 mg·h⁻¹·mg_cat⁻¹, respectively, significantly higher than those of the other two catalysts. Regarding the by-product NO₂⁻, the yields of all three catalysts rose as the voltage increased. The yields of α-MnO₂ were slightly higher than those of β- and δ-MnO₂. Specifically, the yield of α-MnO₂ reached 103.5 mg·h⁻¹·mg_cat⁻¹ at 1.9 V and 128.6 mg·h⁻¹·mg_cat⁻¹ at 2.0 V. At 1.9 V and 2.0 V, the NO₃⁻ yields of α-MnO₂ showed little difference. However, there was a more significant disparity in the yields of the by-product. Taking into account the reaction economy and electrocatalytic activity, α-MnO₂ emerged as the optimal electrocatalyst for the NOOR, and 1.9 V was identified as the optimal reaction potential. At the optimal potential, the bare CP in the control experimental group exhibited only a trace amount of NO₃⁻ output. Moreover, the yield of α-MnO₂/CP was negligible under OCP in an Ar atmosphere (Figure 4c). It was hypothesized that the trace NO₃⁻ yield of the electrode in the Ar atmosphere can be attributed to the N₂ dissolved in the electrolyte. The NO₃⁻ yields of bare CP were also measured at different potentials ranging from 1.6 V to 2.0 V (Figure S8), and the obtained yields were significantly lower than those of α-MnO₂. Therefore, it can be concluded that the product nitrate is primarily generated via the NOOR taking place on the surface of the α-MnO₂ catalyst. Compared to previously reported NOOR electrocatalytic materials, α-MnO₂ exhibited the highest NO₃⁻ yield (Table S1).

Repeatability and stability are also important for assessing the NOOR performance of catalysts. The NOOR reaction was conducted at the α-MnO₂ electrode for six cycles, each for 30 min, using LSV and chronoamperometry curves under optimal voltage and identical experimental conditions to evaluate the reproducibility and cyclic stability of α-MnO₂ electrode. As shown in Figure S9, in the six-cycle experiments, the LSV and chronoamperometry curves of α-MnO₂ overlapped, and the yields (Figure 4d) were nearly identical. This suggests that the α-MnO₂ electrode demonstrated excellent stability and repeatability within a certain range. Furthermore, the long-term stability of the α-MnO₂ electrode was tested for 12 h in a fixed volume electrolyte, with a constant NO intake maintained at 20 mL·min⁻¹. Figure 4e showed a decreasing trend in the current density at the electrode during 12 h of continuous electrolysis, with a decrease of 18% before and after the reaction. It is hypothesized that part of this decrease can be attributed to the increase in the electrolyte mass-transfer resistance caused by the gradual rise in the product concentration in the electrolyte, which tends toward saturation. To eliminate this interference, the performance of the α-MnO₂ electrode in the fresh electrolyte was tested both before and after the 12 h electrolysis. As depicted in the inset of Figure 4e, compared with the unreacted electrode, the electrode after 12 h of prolonged electrolysis did not show a significant decline in the LSV curve. The NO₃⁻ yields of the electrodes before and after electrolysis were examined separately to evaluate their electrocatalytic NOOR performance changes. Figure 4f showed that the NO₃⁻ yield of the α-MnO₂ electrode decreased by about 6.1% after 12 h of continuous use, and the catalytic performance of the electrode decreased slightly, but not significantly. In conclusion, the α-MnO₂ electrode exhibited excellent long-term reaction stability.

Inspired by the remarkable advancements in Zn-air batteries, we developed a Zn-NO reaction device, building upon previous studies [37,38]. This device harnesses the outstanding electrocatalytic capabilities of α-MnO₂ for the NOOR, ingeniously integrating an energy storage system with the electrochemical capture of NO. By using NO as an energy carrier, it can effectively store renewable energy. During the discharge of the Zn-nitrate battery, NO is simultaneously converted into a high-value product NH₃. In the operation process, NO was first directed through the Zn-NO reaction device to undergo the NOOR. As illustrated in Figure 5a, this device was separated by an anion-exchange membrane and employed α-MnO₂ as the anode, Zn plate as the cathode, and 1 M KOH as the electrolyte. In the charge process, NO conducted a continuous oxidation reaction at the α-MnO₂ electrode: NO + 4OH⁻ → NO₃⁻ + 3e⁻ + 2H₂O. Additionally, the cathode reaction was 2H⁺ + 2e⁻ → H₂. The produced NO₃⁻ can be further converted into NH₄OH by the self-powered Zn-NO₃⁻ battery. In the discharge process, the Zn anode was consumed to produce electricity and discharge products (Figure 5b) through the following reactions: Zn + 2H⁺ → Zn²⁺ + H₂ + 2e⁻ and NO₃⁻ + 7H₂O + 8e⁻ → NH₄OH + 9OH⁻. The electrical energy released by this process could light up the light-emitting diode (Figure S10). Therefore, the Zn-NO device offered a two-fold advantage. On one hand, it effectively achieved the resource recovery and environmentally-friendly utilization of NO. On the other hand, it can function as a Zn-NO₃⁻ battery to generate and release electrical energy. Through this dual-function mechanism, it successfully accomplished the conversion of pollutants into valuable substances, thus demonstrating considerable application potential. The electrochemical performance of the Zn-NO cell was investigated. The OCV of the Zn-NO device was 1.4 V (vs. Zn), as shown in Figure 5c. The Zn-NO device was charged using multi-current steps at current densities of 0.0003, 0.005, 0.025, 0.07, and 0.12 mA· cm⁻², which corresponded to the current densities in the three-electrode system at 1.6 V, 1.7 V, 1.8 V, 1.9 V, and 2.0 V (vs. RHE). As shown in Figure 5d, the device voltage remained stable throughout the reaction, reaching 1.84 V, 2.41 V, 3.07 V, 4.02 V, and 5.00 V (vs. Zn), respectively. The reaction achieved an optimum NO₃⁻ yield of 265.5 mg·h⁻¹·mg_cat⁻¹ at 0.07 mA·cm⁻², which was consistent with the optimal response current density in the three-electrode system. The α-MnO₂-based Zn-nitrate battery with strong stability can power multiple light-emitting diodes (as shown in Figure S10), indicating its promising application prospects.

DFT calculations were carried out to explore the microstructure of these three MnO₂. Previous studies had demonstrated a general correlation between the work functions of metal oxide cations in oxidized and reduced states. Specifically, cations in the oxidized state exhibited higher work functions than those in the reduced state [39]. Based on the characteristic XRD peaks of the three MnO₂ in Figure 1a, the work functions of the (211) crystallographic plane of α-MnO₂, the (101) plane of β-MnO₂, and the (−111) plane of δ-MnO₂ were calculated respectively. As illustrated in Figure 6, the work function of α-MnO₂ (7.086 eV) was greater than that of β-MnO₂ (4.968 eV) and δ-MnO₂ (6.026 eV), indicating that the α-MnO₂ had a stronger oxidizing property. The surface energies of the catalyst crystals and the chemisorption energies of the adsorbates were closely associated with the degree of antibonding filling in the surface–adsorbate interaction. Specifically, the higher the degree of antibonding filling, the weaker the interaction between the surface and the adsorbate [40]. Therefore, the bonding strength of NO molecules at different adsorption sites in the three-phase MnO₂ can be quantified by the crystal-orbital Hamiltonian population (COHP). To this end, the integrated crystal-orbital Hamiltonian population (ICOHP) and the density of states (DOS) of Mn were calculated for the adsorption sites of α-, β-, and δ-MnO₂. Figure 6d–f show the contribution of different adsorption sites of the three MnO₂ to bonding (ICOHP > 0). In comparison with β- and δ-MnO₂, α-MnO₂ exhibited a more significant increase in the bonding strength between the substrate and the complex (as shown in Figure 6d). This indicates that there was a stronger interaction between NO and α-MnO₂. The DOS calculation results (Figure S12) were consistent with this. After NO adsorption, the center of the d-band of α-MnO₂ shifted positively, indicating that the catalyst surface had stronger adsorption for NO. This enhanced adsorption improved the catalytic ability of the catalyst [41]. Adsorption energy was used to indicate the strength of adsorption of an intermediate to a surface, with negative or positive adsorption energy indicating that the adsorption energy was energetically favorable or unfavorable to adsorption to the surface [42,43]. To further investigate the adsorption performance of different crystalline phases of MnO₂ towards NO, the adsorption energies of NO adsorbed on three adsorption sites were calculated by taking the (001) crystalline surface of three phases of MnO₂, as shown in

Table 1. Adsorption energy (eV) of different adsorption sites on the (001) plane of α-, β-, and δ-MnO₂.

Catalysts	Top Site	Bridge Site	Hollow Site
α-MnO2	−10.417	−1.1455	−0.8948
β-MnO2	0.8184	7.9489	11.1268
δ-MnO2	0.8494	0.8494	13.1298

. The NO adsorption energy of the three MnO₂ followed the order α-MnO₂ < δ-MnO₂ < β-MnO₂. α-MnO₂ exhibited the lowest adsorption energy, and the adsorption and capture of NO were most likely at its adsorption site, which greatly enhanced the catalytic activity. From these results, it can be inferred that under the same reaction conditions, the differences in the oxidation capacities of the three crystalline phases of MnO₂ towards NO may be associated with their distinct oxidation activities and surface adsorption energies.

3. Experiments

3.1. Synthesis Methods

We used a simple one-step hydrothermal method to synthesize α-, β- and δ-MnO₂, respectively [25,44]. The typical procedures are illustrated in Figure 7, in which three MnO₂ catalysts with different crystal structures were prepared by adjusting the molar ratios of KMnO₄ and MnSO₄·H₂O (n_KMnO4:n_MnSO4·H2O = 2.5, 0.5, and 5.5) and then maintaining them at different reaction temperatures for 12h. The obtained catalysts were calcined in an air atmosphere at 360 °C for 2 h with a heating rate of 5 °C·min⁻¹. Subsequently, α-, β-, and δ-MnO₂ electrodes were fabricated on a carbon paper substrate. All electrochemical tests were carried out at room temperature under a NO gas flow rate of 20 mL·min⁻¹. For specific details, please refer to the Supplementary Information (SI).

3.2. Material Characterization

The morphologies were characterized by scanning electron microscopy (SEM, Hitachi S-4800, Hitachi High-Tech Corporation, Tokyo, Japan) and high-resolution transmission electron microscopy (HRTEM, Jeol 2100F, JEOL Ltd., Tokyo, Japan). X-ray diffraction (XRD) patterns were collected by PANalytical X’pert PRO (Almelo, Netherlands) with Cu Kα radiation to determine the crystal phase of the catalyst samples. X-ray photoelectron spectroscopy (XPS) from the Thermo Scientific ESCALAB 250 Xi (Walthan, MA, USA) system analyzed the surface chemical state. Raman spectroscopy was performed on a Finder Vista Laser micro-Raman Spectroscopy (Zolix Instruments CO., Ltd., Wuhan, China). Absorption spectra were obtained using a UV–visible spectrophotometer (UV-1800, Shimadzu Corporation, Tokyo, Japan).

3.3. Electrochemical Measurement

The electrochemical performances of the α-, β-, and δ-MnO₂ electrodes were tested using an electrochemical workstation (CS300, Corrtest Instruments, Wuhan, China). Electrochemical tests such as linear scanning voltammetry (LSV), cyclic voltammetry (CV), chronoamperometry (CA), electrochemical impedance spectroscopy (EIS), and electrochemically active surface area (ECSA) of the catalysts were accomplished under ambient conditions at room temperature and atmospheric pressure. The NOOR experiments were performed in a 1.0 M KOH electrolyte using a three-electrode H-cell system consisting of a piece of the MnO₂ electrode, a platinum rod, and a Hg/HgO electrode as the working, counter, and reference electrodes, respectively. The α-, β-, and δ-MnO₂ catalysts were loaded on carbon paper (CP) with an area of 1 × 1.5 cm². All potentials were converted to the RHE:

$${\mathrm{E}}_{\mathrm{RHE}}={\mathrm{E}}_{\mathrm{Hg}/\mathrm{HgO}}+0.098 \mathrm{V}+0.059\times \mathrm{pH}$$

Before the start of the experiment, Ar was passed at a flow rate of 15 mL·min⁻¹ to expel the air from the electrolytic cell and to avoid the influence of N₂ on the accuracy of the experiment. During the NOOR experiment, NO (10 vol%) gas flowed uniformly into the electrolytic cell at a 20 mL·min⁻¹ flow rate. Magnetic stirring at 500 rpm was used to assist in the complete dissolution of NO and to ensure the accuracy of the results during sampling and testing. The LSV was scanned at a rate of 10 mV·s⁻¹, and CV measured the ECSA of the catalyst in the non-Faraday region at scanning rates of 10, 20, 40, 60, 80, and 100 mV·s⁻¹, respectively. The EIS was tested over the frequency range of 1000 kHz–0.1 Hz.

3.4. Catalyst Activity Evaluation

The NOOR activity of the electrocatalysts was assessed by calculating the product NO₃⁻ and the by-product NO₂⁻ yield rate (mg·h⁻¹·mg_cat⁻¹). The yield rate was computed using the following equations:

$${{\mathrm{NO}}_3}^{-} \mathrm{yield}=\left(c_{{\mathrm{NO}}_3^{-}}\times V\right)/\left(t\times m_{\mathrm{cat}}\right)$$

$${{\mathrm{NO}}_2}^{-} \mathrm{yield}=\left(c_{{\mathrm{NO}}_2^{-}}\times V\right)/\left(t\times m_{\mathrm{cat}}\right)$$

where c (mg·mL⁻¹) is the measured concentration of NO₃⁻ and NO₂⁻, determined by UV spectrophotometry, V (mL) is the electrolyte volume, t (h) is the electrolysis time, and m_cat is the mass of the catalyst.

3.5. Computational Details

The electronic and microstructural analysis of MnO₂ in three different crystalline phases was performed using CASTEP and Dmol3 of the Materials Studio. Based on the XRD of the three MnO₂, the work function was calculated to explore their electron release capabilities. A differential charge density analysis was performed to investigate the electronic density of the (001) crystal plane of the three MnO₂ phases. The adsorption capacity of NO at the top site, bridge site, and hollow site exposed to the (001) plane of three MnO₂ was investigated, and the N=O bonding strengths and the changes in the density of states and d-band centers of Mn atoms before and after NO adsorption were investigated. The adsorption sites in the (001) surface of the three MnO₂ are shown in Figure 8. The following formula calculates the adsorption energy.

$$E_{\mathrm{ad}}=E_{\mathrm{total}}-E_{\mathrm{NO}}-E^{*}$$

where E_ad, E_total, E_NO, and E* represent the adsorption energy, the total energy of the catalyst and NO in the model, the energy of the NO molecule, and the catalyst’s energy, respectively. The truncation energy used in the calculations is 641.7 eV, and the k-point is set in 3 × 3 × 2.

4. Conclusions

In this paper, the different catalytic properties of three different crystalline MnO₂ for NOOR were investigated. Electrochemical tests showed that α-, β-, and δ-MnO₂ with disparate structures exhibited different surface chemistries and NOOR activities. In the 1.0 M KOH electrolyte, α-MnO₂ possessed the best NOOR catalytic ability, with a NO₃⁻ yield of 665.2 mg·h⁻¹·mg_cat⁻¹ at a potential of 1.9 V, and has good stability and durability. In the Zn-NO device assembled with α-MnO₂ as the anode, the optimum NO₃⁻ yield was 265.5 mg·h⁻¹·mg_cat⁻¹ by combining the energy storage system with NOOR. In addition, the adsorption behavior of NO adsorption by three MnO₂ was investigated using DFT calculations, and the adsorption energies, work functions, and densities of states were calculated for the different crystalline surfaces and adsorption structures. The study of the adsorption activity yielded that the catalytic activity of MnO₂ for NOOR showed α-MnO₂ > β-MnO₂ > δ-MnO₂, which was basically consistent with the experimental results. It is hoped that these results can provide some research ideas for the utilization of MnO₂ for NO electrochemical capture and pollutant resource utilization.