Methane Synthesis from CO₂ and Ammonia over Supported Ni Catalyst via Isocyanate-Mediated Mechanism: Toward Carbon-Neutral Fuel Production

Abstract

Methane (CH₄) synthesis using carbon dioxide (CO₂) and green hydrogen (H₂) is a promising means of increasing CO₂ utilization for the development of a more sustainable society. To mitigate the difficulties associated with the handling of H₂, ammonia (NH₃) has been suggested as a suitable hydrogen carrier, meaning highly active catalysts for the one-step synthesis of carbon-neutral CH₄ from CO₂ and green NH₃ (CO₂ + NH₃ methanation) are needed. Here we show that supported Ni catalysts are potential candidates for CO₂ + NH₃ methanation, and this reaction proceeds via an isocyanate species (*NCO) as an intermediate, which differs from the mechanism of CO₂ + H₂ methanation. Operando DRIFTS analyses revealed that destabilization of *NCO over Ni and efficient hydrogen supply to this intermediate via NH₃ decomposition are important for achieving high CH₄ synthesis activity. A detailed investigation of the physicochemical properties of six different oxide-supported Ni catalysts indicated that these two factors are highly dependent on the basic properties of the oxide support, with the destabilization of *NCO and decomposition of NH₃ being promoted by supports with weak and strong basic strength, respectively. The implication is that catalysts with high CH₄ formation activity could be obtained by using mixed supports with two kinds of basicity. Thus, the present study provides design guidelines for the development of highly active catalysts for the one-step synthesis of CH₄ from CO₂ and NH₃.

1. Introduction

Anthropogenic emissions of carbon dioxide (CO₂), a major contributor to global warming and climate change, have markedly increased since the beginning of the 20th century. To preserve the global environment and facilitate the development of a more sustainable society, technologies for the reduction of CO₂ emissions are urgently needed. However, simply reducing the amount of CO₂ produced will likely not be sufficient, meaning technologies for the capture and utilization of CO₂ are also attracting attention.

Among the various CO₂ utilization reactions that have been proposed, much attention is currently being paid to CO₂ methanation, which is the production of carbon-neutral methane (CH₄) from the reaction of CO₂ and green hydrogen gas (H₂) (eq ). CH₄ is a valuable fuel for domestic use, industrial processes, and thermal power generation, and it can be easily transported and stored using the current pipeline infrastructure for natural gas. , However, currently, the construction of large-scale green H₂ production plants is limited to sites with sufficient renewable energy resources such as wind, solar radiation, and water sources. Under such circumstances, many regions lacking renewable energy resources would have to import green H₂, and H₂ is not easy to store and transport because it requires extremely low temperatures (−253 °C) for liquefaction. The most promising H₂ carrier currently is ammonia (NH₃) due to its high H₂ storage capacity per volume (121 kg_H2 m^–3) and ease of liquefaction under mild conditions (20 °C, 0.8 MPa). Therefore, with further technological advances, it is expected that a green H₂ supply network will be established that uses green NH₃ as a carrier, making NH₃ the starting point for CO₂ utilization in regions where renewable energy resources are scarce. In such systems, CH₄ would be produced via a two-step process of NH₃ decomposition to afford H₂ (NH₃ decomposition; eq ) followed by CO₂ hydrogenation to afford CH₄. However, such a two-step process would likely not be economical because of energy loss and the need for dedicated reaction facilities suitable for the two processes of NH₃ decomposition and methanation. Thus, if CH₄ could be produced by direct reaction of CO₂ and NH₃ (hereafter “CO₂ + NH₃ methanation”; eq ), the utilization of CO₂ would become more economical and efficient. Furthermore, using NH₃ as a hydrogen source would lower the reaction heat per CH₄ produced compared to CO₂ + H₂ methanation. This would help prevent the formation of hot spots, which are the main cause of catalyst deactivation. , Additionally, a simple process using an adiabatic reactor suitable for large-scale production could be designed. Because of these advantages, CO₂ + NH₃ methanation has the potential to contribute to the realization of carbon-neutral fuel production processes. Thus, the development of highly active catalysts for CO₂ + NH₃ methanation represents an important step toward increasing CO₂ utilization.

$$\begin{array}{l}{\mathrm{CO}}_2+4{\mathrm{H}}_2\to {\mathrm{CH}}_4+2{\mathrm{H}}_2\mathrm{O}\\ \mathrm{\Delta }{H}_{298\mathrm{K}}=-165\mathrm{kJ}/{\mathrm{mol}}_{\mathrm{CH}4}\end{array}$$ 1

$$\begin{array}{l}{\mathrm{NH}}_3\to 1/2{\mathrm{N}}_2+3/2{\mathrm{H}}_2\\ \mathrm{\Delta }{H}_{298\mathrm{K}}=+46\mathrm{kJ}/{\mathrm{mol}}_{\mathrm{NH}3}\end{array}$$ 2

$$\begin{array}{l}{\mathrm{CO}}_2+8/3{\mathrm{NH}}_3\to {\mathrm{CH}}_4+4/3{\mathrm{N}}_2+2{\mathrm{H}}_2\mathrm{O}\\ \mathrm{\Delta }{H}_{298\mathrm{K}}=-43\mathrm{kJ}/{\mathrm{mol}}_{\mathrm{CH}4}\end{array}$$ 3

Oxide-supported metal catalysts are widely recognized as promising heterogeneous catalysts for CO₂ + H₂ methanation. − A few previous studies have demonstrated that ruthenium (Ru), a noble metal, and nickel (Ni), a non-noble metal, are both highly active for both CO₂ + H₂ methanation and NH₃ decomposition. − Therefore, these could be considered as potential candidates for catalyzing CO₂ + NH₃ methanation. Indeed, Uddin et al. evaluated CO₂ + NH₃ methanation activity over Ru and Ni supported Al₂O₃ catalysts and reported that the Ru catalyst was significantly more active than the Ni catalyst. In addition, Saima et al. investigated physically mixing two different oxide-supported Ni catalysts, Ni/Al₂O₃ and Ni/CeO₂, and observed marked CH₄ formation from CO₂ and NH₃ over the mixed catalyst. Although these previous studies show that Ru and Ni are effective at catalyzing CO₂ + NH₃ methanation, our understanding of the reaction mechanisms and catalyst properties that affect catalytic activity remains incomplete. Furthermore, the activities of these reported catalysts remain insufficient.

Generally, the activity of a supported metal catalyst depends strongly on the physicochemical properties of the support material. For example, the acid/base property of the support is an important factor in determining catalytic activity because it directly modulates the adsorption and desorption of gas molecules and the electronic state of the active metal. , For CO₂ + H₂ methanation, supports with a large number of basic sites with moderate basicity are favorable for the activation of absorbed CO₂ species. , However, for NH₃ decomposition, it has been reported that support materials with strong basicity are most effective because they enhance the ability of the support to donate electrons to the active metal and mitigates H₂ poisoning. , Thus, for the development of effective catalysts for CO₂ + NH₃ methanation, it will be essential to understand the influence of the acid/base properties of different supports.

Here, we elucidated details of the reaction mechanism of CO₂ + NH₃ methanation over supported Ni catalysts based on operando DRIFTS analysis. The results obtained provide valuable insights into the surface state of the catalyst and the dynamic behavior of adsorption species under reaction conditions. In addition, we identified the physicochemical properties of the support that most affect catalyst activity. Supported Ru catalysts have shown excellent physicochemical properties and high activity per metal amount for CO₂ + NH₃ methanation. However, supported Ni catalysts are advantageous over supported Ru catalysts in terms of practicality because of the abundance of Ni deposits and therefore lower cost of catalyst manufacture. Six supported Ni catalysts using different oxides (Al₂O₃, MgO, SiO₂, CeO₂, ZrO₂, Y₂O₃) were prepared, and their activity for CO₂ + NH₃ methanation were studied. Operando DRIFTS analyses and analyses of the physicochemical properties of the catalysts were conducted to examine the reaction mechanism, revealing that CH₄ formation proceeds via an intermediate isocyanate species (*NCO), which is different from how the CO₂ + H₂ methanation reaction proceeds. Furthermore, the degree of stabilization of *NCO and the availability of hydrogen through NH₃ decomposition were both found to be crucial factors for CH₄ formation, with the basic properties of the oxide support being found to strongly affect these factors.

2. Experimental Section

2.1. Catalyst Preparation

Six oxide-supported Ni catalysts were prepared by an impregnation method after the oxide support was calcined. Ni­(NO₃)₂·6H₂O (Fujifilm Wako Pure Chemical Co., Ltd., Japan) was used as the nickel precursor. The following six oxides were used as the support materials: Al₂O₃ (AKP-G15; Sumitomo Chemical Co., Ltd., Japan), MgO (500A; Ube Material Industries, Ltd., Japan), SiO₂ (JRC-SIO-15; Fuji Silysia Chemical, Ltd., Japan), CeO₂ (JRC-CEO-6; Daiichi Kigenso Kagaku Kogyo Co., Ltd., Japan), ZrO₂ (JRC-ZRO-6; Daiichi Kigenso), and Y₂O₃. Y₂O₃ was prepared by a precipitation method using nitrate (Y­(NO₃)₃·nH₂O; Fujifilm Wako) and NH₃ water. The detailed methods are shown in the Supporting Information (Supporting Information S1.1). Prior to the impregnation, all oxide supports except the Y₂O₃ were calcined at 700 °C for 3 h in air. The support material was then added to an aqueous solution of nickel nitrate, and the mixture was stirred overnight at room temperature. The next day, the water was removed by evaporation, and the remaining slurry was dried overnight at 80 °C. The resulting powder was crushed and calcined at 500 °C for 5 h under flowing air. The Ni loading was fixed at 20 wt %. X-ray fluorescence analysis revealed that the actual Ni loading was comparable for each catalyst (see Table S1).

2.2. Catalytic Activity Test

The obtained catalyst powders were pelletized and grained to 250–500 μm. The CO₂ + NH₃ methanation activity over each catalyst was measured by using a conventional flow system with a tubular reactor under atmospheric pressure. After loading 150 mg of catalyst, the catalysts were prereduced in situ with pure H₂ (50 mL min^–1) for 1 h at 0.1 MPa and 400, 500, or 900 °C. The prereduction temperature was optimized for each catalyst, as the reducibility of Ni species is highly dependent on the support (see Section for details). After reduction, the catalyst was cooled to 300 °C under flowing Ar (50 mL min^–1) and then a mixture of CO₂ (5 mL min^–1) and NH₃ (13 mL min^–1) was passed over the catalyst (space velocity [SV] = 7200 mL h^–1 g_cat. ^–1). When CO₂ and NH₃ gas (and H₂O) are mixed below 150 °C, the formation of ammonium carbonate ((NH₄)₂CO₃), ammonium bicarbonate ((NH₄)­HCO₃), and ammonium carbamate (NH₂COONH₄) may cause blockages in gas pipes. To prevent the formation of such ammonium salts, a ribbon heater was used to maintain the temperature of the gas pipes at 150 °C. To remove unreacted ammonia by absorption in water, the effluent gas mixture was passed through a gas-washing bottle filled with pure water that was connected to the outlet of the reactor. A cold trap was also attached to the outlet of the gas washing bottle to remove the water vapor. The temperature of the catalyst was kept constant for 30 min, and the composition of the gas mixture passed through the gas-washing bottle and the cold trap was analyzed using an online gas chromatograph (Agilent 490 MicroGC, Agilent) equipped with a thermal conductivity detector and Molsieve 5A and PoraPLOT Q columns. This procedure was repeated at increments of 50 °C up to 700 °C. During the activity test, Ar (60 mL min^–1) was mixed into the gas mixture passed through the gas-washing bottle and the cold trap, and the total outlet gas flow rate was calculated using the flow rate of Ar as the internal standard. CH₄ yield, CO yield, and NH₃ conversion were calculated according to the following equations

$${\mathrm{CH}}_4\mathrm{yield}(\%)={\mathrm{CH}}_{4\mathrm{outlet}}/{\mathrm{CO}}_{2\mathrm{inlet}}\times 100$$

$$\mathrm{CO}\mathrm{yield}(\%)={\mathrm{CO}}_{\mathrm{outlet}}/{\mathrm{CO}}_{2\mathrm{inlet}}\times 100$$

$${\mathrm{NH}}_3\mathrm{conversion}(\%)={\mathrm{NH}}_{3\mathrm{consumed}}/{\mathrm{NH}}_{3\mathrm{inlet}}\times 100$$

where CO_{2 inlet} and NH_3 inlet are the flow rates of CO₂ and NH₃ at the reactor inlet (CO_{2 inlet}: 5 mL min^–1, NH_3inlet: 13 mL min^–1); CH_{4 outlet} and CO_outlet are the flow rate of CH₄ and CO at the outlet of the cold trap; and NH_{3 consumed} is the flow rate of NH₃ consumed in the reaction and was calculated based on the stoichiometric ratio of the outlet flow rate of N₂ (N_{2 outlet}) as NH_{3 consumed} = N_{2 outlet} × 2.

The experimental results were compared with thermodynamic equilibrium values calculated by the HSC Chemistry 6.1 commercial steady-state simulation package (Outotec Research Oy. Pori, Finland) based on the minimization of Gibbs free-energy.

For the stability test, CH_{4 outlet}, CO_outlet, and N_{2 outlet} were recorded versus time on stream at an SV of 7200 mL h^–1 g^–1 at 500 °C.

The activity of the catalysts for CO₂ + H₂ methanation and NH₃ decomposition were also measured using the same instrument setup. In each reaction, the amount of catalyst and the conditions for prereduction were the same as for the CO₂ + NH₃ methanation experiment. In the CO₂ + NH₃ methanation experiment, the gas composition when NH₃ was 100% decomposed was CO₂: 5 mL min^–1, H₂: 19.5 mL min^–1, and N₂: 6.5 mL min^–1. In the CO₂ + H₂ methanation experiment, the total gas flow rate was adjusted by adjusting the Ar flow rate such that the CO₂ and H₂ partial pressures were the same as those of the above gas composition (CO₂: 5 mL min^–1, H₂: 19.5 mL min^–1, Ar: 6.5 mL min^–1; SV = 12,400 mL h^–1 g^–1). In the NH₃ decomposition experiment, the total gas flow rate was adjusted by adjusting the He flow rate such that the NH₃ partial pressure was the same as that in the CO₂ + NH₃ methanation experiment (NH₃: 13 mL min^–1, He: 5 mL min^–1; SV = 7200 mL h^–1 g^–1). Detailed descriptions of the gas conditions for each reaction are provided in the Supporting Information (Supporting Information S1.2.).

2.3. Operando Diffuse Reflectance Infrared Fourier-Transform Spectroscopy (Operando DRIFTS)

Operando DRIFT spectroscopy was performed using an infrared spectrometer (FT/IR-4700 spectrometer; Jasco) equipped with a MCT detector. About 30 mg of finely ground catalyst or support was placed in a ceramic cup, and the cup was loaded inside the DRIFTS cell. Before measurement, the sample was reduced in situ under 100% H₂ stream (20 mL min^–1) for 1 h. The temperature of prereduction was the same for each catalyst as for the catalytic activity test. After prereduction, the temperature of the sample chamber was cooled to 100 °C under He flow (20 mL min^–1). The background spectrum was collected in the He atmosphere at 100 °C with a resolution of 4 cm^–1 and an average of 64 scans, and it was later subtracted from the spectrum acquired for the test sample. DRIFT spectra for the samples were acquired by using a CO₂ + NH₃ methanation (CO₂/NH₃/He = 1:2.6:6.4 mL min^–1) or CO₂ + H₂ methanation atmosphere (CO₂/H₂/He = 1:3.9:7.7 mL min^–1) and increasing the temperature of the sample chamber from 100 to 500 °C (ramping rate: 10 °C min^–1).

2.4. Catalyst Characterization

The specific surface areas of the supports and catalysts were determined by nitrogen adsorption at −196 °C by using a BELSORP-mini X instrument (MicrotracBEL, Japan) and the Brunauer–Emmett–Teller model.

X-ray diffraction analysis was performed by using a SmartLab X-ray diffractometer (Rigaku, Japan) equipped with a copper K-α radiation source. X-ray diffraction patterns were analyzed by using the PDXL2 software ver. 2.8.4.0 (Rigaku) and the International Centre for Diffraction Data, Crystallography Open, and AtomWork databases.

H₂ temperature-programed reduction (H₂-TPR) measurements were performed using a BELCAT-II apparatus (MicrotracBEL) to investigate the reducibility of Ni species for each catalyst. In a H₂-TPR measurement, 150 mg of calcined catalyst (before reduction) was loaded into the reactor and pretreated at 150 °C for 3 h. The catalyst was then reduced in 100% H₂ flow (50 mL min^–1) while the temperature was increased from room temperature to 1000 °C at a ramping rate of 10 °C min^–1. The H₂O profiles were monitored by a quadrupole mass spectrometer at m/z = 18.

CO-pulse chemisorption measurements were performed using a BELCAT-II apparatus to investigate the Ni surface area of each catalyst. In a CO-pulse chemisorption measurement, 150 mg of calcined catalyst (before reduction) was loaded into the reactor and prereduced in 100% H₂ flow (50 mL min^–1) for 1 h. The temperature of the prereduction was the same for each catalyst as for the activity test. Then, a mixture of 9.96% CO in He was injected in pulses 10 times. The Ni surface area (m_Ni ² g^–1) was calculated from the amount of CO adsorbed per catalyst mass (V _m : cm³ g^–1) according to the following equation:

$$\mathrm{Ni\; surface\; area}({{\mathrm{m}}^2}_{\mathrm{Ni}}{\mathrm{g}}^{-1})=V_{\mathrm{m}}\times (\mathrm{SF}/22414)\times 6.02\times {10}^{23}\times {\sigma }_{\mathrm{m}}\times {10}^{-18}$$

where SF (= 1) is the stoichiometric factor for CO chemisorption, and σ_m is the cross-sectional area of one Ni atom (= 0.0649 nm²).

CO₂ temperature-programed desorption (CO₂-TPD) measurements were performed using a BELCAT-II apparatus to investigate the amount and strength of CO₂ adsorption on each catalyst. In a CO₂-TPD measurement, 100 mg of calcined catalyst (before reduction) was loaded into the reactor and prereduced in 100% H₂ flow (50 mL min^–1) for 1 h. The temperature of the prereduction was the same for each catalyst as for the activity test. Then, the sample was exposed to a CO₂/He (CO₂: 4.97%) gas mixture (50 mL min^–1) at 40 °C for 120 min. The physiosorbed CO₂ was removed by He purging at 40 °C for 120 min. Then, the temperature was increased to 1000 °C at a ramping rate of 10 °C min^–1 under He flow (30 mL min^–1). The desorbed CO₂ was monitored by a quadrupole mass spectrometer at m/z = 44.

X-ray absorption fine structure (XAFS) analyses were performed on each supported Ni catalyst after reduction treatment at the BL14B2 beamline at SPring-8 (Hyogo, Japan) with approval from the Japan Synchrotron Radiation Research Institute. An Si(111) double-crystal monochromator was used for the monochromation of X-rays. The reduced catalysts were ground well with a calculated amount of boron nitride powder (Fujifilm Wako) and pelletized into self-supporting disks (ϕ = 7 mm). To analyze the structure of the catalyst after the durability test, another XAFS measurement was performed at the BL11S2 beamline at Aichi SR (Aichi, Japan). To prevent the catalyst from being exposed to air, the inlet and outlet of the reactor were both sealed after the durability test. Then, the reactor was placed inside and unsealed in an Ar-filled glovebox. In the glovebox, a disk (ϕ = 7 mm, diluted with boron nitride) was prepared. In the XAFS measurements, Ni K-edge spectra were recorded at room temperature in transmission mode. The XAFS spectra were analyzed using ATHENA software and the Ni K-edge X-ray absorption near edge structure (XANES) spectrum was obtained to elucidate the electronic state of Ni.

High-angle annular dark-field scanning transmission electron microscopy (HAADF-STEM) images of the supported Ni catalysts were obtained using a JEM-2100F HK electron microscope (JEOL, Japan) operating at 200 kV. The samples were dispersed in ethanol at room temperature, and the dispersion was dropped onto a carbon-coated copper grid and vacuum-dried for 24 h at room temperature.

3. Results and Discussion

3.1. Influence of Support Material on CO₂ + NH₃ Methanation

The activity of each supported Ni catalyst for CO₂ + NH₃ methanation was evaluated. In this experiment, the effluent gas from the reactor was passed through a gas-washing bottle filled with pure water to remove unreacted NH₃. However, a part of the unreacted CO₂ was also captured by the water in the bottle because the dissolved NH₃ makes the water alkaline. This makes estimation of the CO₂ conversion based on the composition of the effluent gas difficult. For example, the CO₂ conversion afforded by the Ni/Al₂O₃ catalyst calculated from the CO₂ concentration at the outlet of the gas-washing bottle was approximately 50% at 300 °C (Figure S1), despite the reaction’s hardly progressing because CH₄ and CO production were not observed (Refer to Figure a,b). Therefore, hereafter the activity is discussed in terms of yields of CH₄ and CO.

1.

Temperature dependence of (a) CH₄ yield, (b) CO yield, (c) NH₃ conversion over six oxide-supported Ni catalysts. CO₂: 5 mL min^–1, NH₃: 13 mL min^–1 at 0.1 MPa. Equilibrium (1) is the theoretical CH₄ yield, calculated assuming five products: CH₄, CO, N₂, H₂, and H₂O. Equilibrium (2) is the theoretical CO yield calculated assuming only four products: CO, N₂, H₂, and H₂O. (d) Relationship between CH₄ yield and NH₃ conversion at 500 °C for each Ni catalyst.

The relationship between CH₄ yield and reaction temperature is shown in Figure a. Equilibrium values were calculated assuming the following five products: CH₄, CO, N₂, H₂, and H₂O (denoted as Equilibrium (1)). At reaction temperatures below 400 °C, the CH₄ yield was negligible over each catalyst. However, marked increases of CH₄ yield were observed from 450 °C over Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃, and from 500 °C over Ni/CeO₂, Ni/MgO, and Ni/ZrO₂, with the former three catalysts affording higher CH₄ yields (approximately 40% at 550 °C) than the latter three. Over Ni/Al₂O₃, a CH₄ yield of around 40% was obtained at 500 °C even when the SV was reduced to 3600 mL h^–1 g^–1 (Figure S2a). Although a direct comparison of the activities of the present catalysts with those of previously reported catalysts , is challenging because of differences in the amounts of metal loaded and reaction conditions, the CH₄ yield obtained over our Ni/Al₂O₃ catalyst was higher than that of two previously reported catalysts (Table S2). The CH₄ yields of Ni/MgO and Ni/CeO₂ at 550 °C were lower than those of Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃. The activity of Ni/ZrO₂ was the lowest among the tested catalysts, with a CH₄ yield of less than 20% at 550 °C.

When the catalysts were used for CO₂ + H₂ methanation (Figure S3a), the rate of CH₄ formation increased at temperatures between 250 and 300 °C for all samples. Ni/CeO₂ and Ni/Y₂O₃ exhibited the highest activities, with CH₄ yields exceeding 70% at 250 °C. In contrast, the CH₄ yields afforded by Ni/Al₂O₃ and Ni/SiO₂ were less than 30% at the same temperature. These results show that CO₂ + NH₃ methanation requires higher temperatures than does CO₂ + H₂ methanation for CH₄ production and that the best support for CH₄ formation differs between the two types of methanation.

CO formation during CO₂ + NH₃ methanation was examined for each of the catalysts (Figure b). CO was formed over each of the catalysts at temperatures above 350 °C. At temperatures below 450 °C, the production of CO was considered to follow an equilibrium calculated assuming only four products: CO, N₂, H₂, and H₂O (excluding CH₄) (denoted Equilibrium (2)), based on the earlier finding that no CH₄ was formed below 450 or 500 °C for any of the catalysts (Figure a). Equilibrium (1) was again calculated assuming five products, as in Figure a. From 350 to 400 °C, the CO yield increased similarly for each catalyst. The fact that the CO yield exceeded Equilibrium (1) at these temperatures suggested that the reaction proceeded according to Equilibrium (2) at low temperatures. At temperatures above 550 or 600 °C, the CO yields for the catalysts were approximately equal to Equilibrium (1). Thus, the reaction equilibrium was shifted from Equilibrium (2) to Equilibrium (1) by increasing the reaction temperature to more than 550 or 600 °C. In the case of CO₂ + H₂ methanation, essentially no CO was formed at temperatures below 400 °C (Figure S3b), whereas during CO₂ + NH₃ methanation, CO was formed over all of the catalysts, even at temperatures below 400 °C (Figure b). These results indicate that CO formation is predominant at low temperatures during CO₂ + NH₃ methanation, and that CO₂ hydrogenation proceeds via a different mechanism from that of CO₂ + H₂ methanation. Figure S2b shows the SV dependence of CO yield over Ni/Al₂O₃. With decreasing SV, CO yield was decreased. In contrast, the CH₄ yield increased (Figure S2a). This suggests that increasing the CH₄ formation rate can suppress the CO formation rate. Also, it was considered that CO could be an intermediate in the formation of CH₄, or a reaction pathway could have been formed in which CO is preferentially produced at low temperatures.

NH₃ conversion during CO₂ + NH₃ methanation was examined for each of the catalysts (Figure c). The calculation conditions for Equilibrium (1) were the same as those for the analyses shown in Figures a,b. Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃ afforded the highest NH₃ conversions irrespective of temperature. This trend is consistent with that observed for CH₄ yield (Figure a). The relationship between CH₄ yield and NH₃ conversion was investigated at 500 °C (Figure d). At 500 °C, marked CH₄ formation was observed for each catalyst and the yield remained much lower than the equilibrium value. It is therefore possible to compare the reaction rates under kinetic control at this temperature. A linear correlation between CH₄ yield and NH₃ conversion was observed. Figure S2c shows the SV dependence of NH₃ conversion over Ni/Al₂O₃. The NH₃ conversion increased with decreasing SV, while the CH₄ yield increased (Figure S2a), which is consistent with the trend shown in Figure d.

In this experiment, the NH₃ conversion was calculated based on the amount of N₂ detected at the outlet (N_2 outlet). The formation of N-containing compounds other than N₂ (e.g., HCN and NO _x ) was investigated using a mass spectrometer. Figure S4 shows the composition of the outlet gas (upstream of the gas-washing bottle) during the CO₂ + NH₃ methanation activity test using Ni/Al₂O₃. The m/z = 28 (N₂ or CO) signal was detected over the entire test temperature range. However, m/z = 27 (HCN) and m/z = 30 (NO₂, NO) signals were hardly detected. Therefore, the amount of N-containing products other than N₂ was almost negligible, and we concluded that the N-balance of the reaction can be adequately discussed based on the outlet flow rate of N₂. Also, since no carbon-containing compounds other than CH₄ and CO were detected, the total amount of CH₄ and CO produced was assumed to equal the amount of CO₂ consumed during the reaction. CO₂ conversion and CH₄ selectivity were calculated based on the following equations and are shown in Figure S5a,b, respectively.

$${\mathrm{CO}}_2\mathrm{conversion}(\%)=(({\mathrm{CH}}_{4\mathrm{outlet}}+{\mathrm{CO}}_{\mathrm{outlet}})/{\mathrm{CO}}_{2\mathrm{inlet}})\times 100$$

$${\mathrm{CH}}_4\mathrm{selectivity}(\%)=({\mathrm{CH}}_{4\mathrm{outlet}}/({\mathrm{CH}}_{4\mathrm{outlet}}+{\mathrm{CO}}_{\mathrm{outlet}}))\times 100$$

The catalysts could be classified into two groups based on their CO₂ conversion behavior with respect to reaction temperature, with one group comprising Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃ and the other Ni/MgO, Ni/CeO₂, and Ni/ZrO₂ (Figure S5a). The former three exhibited higher CO₂ conversion activity than the latter three, which is supported by their superior CH₄ formation rates (see Figure a). Of the latter three, Ni/MgO and Ni/CeO₂ showed higher CH₄ yields than Ni/ZrO₂ (see Figure a), but their CO₂ conversions were almost identical. This was attributed to the CO yield’s of Ni/ZrO₂ being higher than that of Ni/MgO and Ni/CeO₂ at 500 and 550 °C (see Figure b). Indeed, the CH₄ selectivity of Ni/ZrO₂ was lower than that of Ni/MgO and Ni/CeO₂ (Figure S5b).

In the CO₂ + NH₃ methanation test, it was confirmed that the CH₄ yield, CO yield, and NH₃ conversion could be measured with good reproducibility when using the Ni/Al₂O₃ catalyst (Figure S6).

The intrinsic NH₃ decomposition activity (i.e., the ammonia decomposition activity in an atmosphere where CO₂ is not present; NH₃: 13 mL min^–1, He: 5 mL min^–1) of each catalyst was also evaluated (Figure S7). In the absence of CO₂, the change of NH₃ conversion versus reaction temperature was different among Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃. Ni/Y₂O₃ exhibited a much higher NH₃ conversion than the other catalysts at low temperatures (350–550 °C). Ni/Al₂O₃ and Ni/SiO₂ showed lower NH₃ conversions compared with Ni/Y₂O₃. Thus, CO₂ + NH₃ methanation activity was correlated with NH₃ conversion, and the catalysts that exhibit high NH₃ conversion differ depending on presence or absence of CO₂.

3.2. Reaction Mechanism of CO₂ + NH₃ Methanation

The results obtained so far suggested that the optimal catalyst properties differed between CO₂ + NH₃ methanation and CO₂ + H₂ methanation. Furthermore, the NH₃ conversion activity of the catalysts differed depending on the presence or absence of CO₂. These results suggest that CO₂ + NH₃ methanation is not a simple combination of NH₃ decomposition and CO₂ + H₂ methanation. Therefore, we investigated the reaction mechanism of CO₂ + NH₃ methanation by means of an Operando DRIFTS analysis.

To help understand the differences between CO₂ + H₂ methanation and CO₂ + NH₃ methanation, the reaction mechanism of the former was investigated. DRIFT spectra were obtained for the Ni/Al₂O₃ catalyst under CO₂ and H₂ flow (CO₂: 1.0 mL min^–1, H₂: 3.9 mL min^–1, and He: 7.7 mL min^–1). During the measurements, the temperature of the sample chamber was increased at a ramping rate of 10 °C min^–1 from 100 to 500 °C. The spectra obtained at different reaction temperatures are shown in Figure . At temperatures of 100 and 200 °C, in addition to the band assigned to gas-phase CO₂ at 2300–2400 cm^–1, bands assigned to bicarbonate (HCO₃*) were observed at 1650, 1431, and 1226 cm^–1. , This species is a well-known adsorbed CO₂ species on oxides. When the temperature was raised to 300 °C, the intensity of the HCO₃* bands decreased, while bands assigned to formate species (HCOO*) appeared at 1590, 1390, and 1374 cm^–1. , Furthermore, bands assigned to carbonyl species on Ni (CO*) and gas-phase CH₄ (3015 cm^–1) were also observed. The CO* species exhibited multiple adsorption configurations, with the higher wavenumber (2003 cm^–1) corresponding to linearly adsorbed species, and the lower wavenumbers (1914 and 1866 cm^–1) corresponding to bridged adsorbed species. These changes in the DIRFT spectra with increasing temperature suggest that HCOO* is formed by hydrogenation of HCO₃*, and CO* is formed by hydrogenation of HCOO*, at the Ni–support interface. At 400 and 500 °C, the HCOO* bands decreased, and the intensity of the CH₄ gas bands increased. Several researchers have investigated the reaction mechanism of CO₂ + H₂ methanation over supported catalysts by DRIFTS analysis and have proposed that CH₄ is produced through hydrogenation of CO* on the active metal. , To verify whether CO* is an intermediate for CH₄ in CO₂ + H₂ methanation, the change of the IR band of CO* was observed under H₂ flow. To ensure formation of CO* on Ni/Al₂O₃, CO₂ and H₂ gases (CO₂: 1.0 mL min^–1, H₂: 3.9 mL min^–1, He: 7.7 mL min^–1) were passed over the catalyst at 300 °C for 20 min. Then, the sample chamber was purged with He gas (20 mL min^–1) while cooling to 100 °C. After that, the temperature was increased from 100 to 300 °C with feeding of H₂ and He gas (H₂: 3.9 mL min^–1, He: 8.7 mL min^–1). During the heating, the transient response of the DRIFT spectrum (Figure S8a) was recorded, and the outlet gas composition was observed with a mass spectrometer to investigate the formation behavior of m/z = 15 (CH₄) (Figure S8b). In Figure S8a, the IR band of HCO₃* decreased, and the band of HCOO* increased between 100 and 150 °C. At this time, the MS signal of m/z = 15 barely increased (Figure S8b), indicating that HCO₃* is not an intermediate of CH₄. From 150 to 200 °C, the IR band of CO* decreased. In this temperature range, an increase in the m/z = 15 signal was observed. The HCOO* band also increased at 150–200 °C. Since this experiment was conducted under a flow of H₂, HCOO* is thought to have been produced by hydrogenation of adsorbed CO₂ rather than by oxidation of CO* by an oxidizing agent such as byproduct H₂O. From these results, it was considered that CH₄ is formed from CO* similar to that reported in previous studies , (Scheme ).

2.

DRIFT spectra for the Ni/Al₂O₃ catalyst under a CO₂ + H₂ methanation atmosphere (CO₂: 1.0 mL min^–1, H₂: 3.9 mL min^–1, He: 7.7 mL min^–1).

1. Proposed Mechanism for CH₄ Formation during CO₂ + H₂ Methanation over Oxide-Supported Ni Catalysts.

Figure shows the DRIFT spectra obtained under CO₂ and NH₃ flow (CO₂: 1.0 mL min^–1, NH₃: 2.6 mL min^–1, He: 6.4 mL min^–1) for the Ni/Al₂O₃ catalyst. The reaction temperatures and ramping rates were the same as in the experiment shown in Figure . At each temperature, in addition to the absorption band of CO₂ gas (∼2350 cm^–1), bands attributable to NH₃ gas were observed at 1625 and 3334 cm^–1. At 100 °C, bands attributable to HCOO* and carbonate (CO₃*) were observed at 1580 and 1434 cm^–1, respectively. When the reaction temperature was increased to 200 °C, new bands attributable to isocyanate species (*NCO) appeared at 2251 and 2178 cm^–1. , The formation of *NCO suggests the formation of isocyanic acid (HNCO) (eq ) on the catalyst, followed by the dissociation of its N–H bond. It has been reported that ammonium carbamate (NH₂COONH₄) and urea (NH₂CONH₂) are formed from CO₂ and NH₃ and that urea is readily converted to isocyanic acid by thermal decomposition. ,

3.

DRIFT spectra for the Ni/Al₂O₃ catalyst under a CO₂ + NH₃ methanation atmosphere (CO₂: 1.0 mL min^–1, NH₃: 2.6 mL min^–1, He: 6.4 mL min^–1).

When DRIFT spectra were obtained under the same gas conditions using an Al₂O₃ support alone (i.e., without Ni loading), the *NCO band was observed at 2245 cm^–1 only at temperatures above 200 °C (Figure S9). Therefore, the two bands that appeared in the spectrum for Ni/Al₂O₃ were attributed to *NCO adsorbing on different sites, with the higher wavenumber band corresponding to *NCO adsorbed on Al₂O₃ and the lower wavenumber band corresponding to *NCO adsorbed on Ni. At 300 °C, the intensity of the bands of *NCO on Al₂O₃ and Ni increased, suggesting that the formation of HNCO was promoted. In contrast, the intensities of the HCOO* and CO₃* bands were decreased. It is possible that these CO₂-adsorbed species were thermally decomposed or consumed to form *NCO. At 400 °C, the intensity of the band attributed to *NCO-adsorbed Ni decreased. When the temperature was raised to 500 °C, the bands of the CO₂ adsorption species almost disappeared, and the intensity of the *NCO bands markedly decreased. At this temperature, a band attributable to gas-phase CH₄ (3015 cm^–1) was observed. Note that, in contrast to the spectrum observed under the CO₂ + H₂ atmosphere, bands assignable to CO* on Ni were not observed at any temperature. The data in Figure b show that CO was formed at temperatures above 350 °C during CO₂ + NH₃ methanation. However, CO* was not observed in the DRIFTS measurements, presumably because the CO adsorption sites on Ni were occupied by *NCO and NH₃. Therefore, in CO₂ + NH₃ methanation, CH₄ must be formed from adsorbed species other than CO*, suggesting a different reaction mechanism compared with that of CO₂ + H₂ methanation. Furthermore, the absence of CO* suggests that CO poisoning is unlikely to occur in CO₂ + NH₃ methanation.

To investigate interference from other nitrogen-containing species, DRIFT spectra were acquired for the Ni/Al₂O₃ catalyst under NH₃ gas flow (NH₃ 2.6, He 7.4 mL min^–1) (Figure S10). The fact that almost no IR bands other than those of gas-phase NH₃ (3334, 1625 cm^–1) were observed showed that the bands observed in Figure were not affected by interference from other nitrogen-containing species such as *NH₂ and *NH.

The rates of formation of each reaction product (CH₄, CO, N₂, H₂) with respect to reaction temperature in the activity tests using Ni/Al₂O₃ are shown in Figure (inset shows results at a low temperature of 300–400 °C). The formation rates of N₂ and CO began to increase from 350 °C. In the DRIFTS analysis, the intensity of the *NCO bands decreased between 300 and 400 °C, suggesting that N₂ and CO were formed by thermal decomposition of *NCO (eq ). In addition, the rate of formation of H₂ increased in the same temperature range, suggesting that H₂ formation occurred from H atoms dissociated from HNCO. In this reaction, the ratio of NH₃/CO₂ ([8/3]/1) in the reactant gases was higher than the stoichiometric ratio of the HNCO formation reaction (1/1) (eq ). Therefore, it is possible that excess NH₃, which is not involved in the formation of HNCO, was adsorbed on Ni together with *NCO, and that H₂ formation also occurred by NH₃ decomposition. At higher temperatures, the rates of formation of N₂ and H₂ continued to increase, but the rate of CO formation remained low at temperatures above 450 °C where formation of CH₄ occurred (Figure ). The increase in H₂ partial pressure may have resulted in the hydrogenation of *NCO to CH₄ (eq ) and thus suppressed the CO formation associated with the thermal decomposition of *NCO.

$${\mathrm{CO}}_2+{\mathrm{NH}}_3\to \mathrm{HNCO}+{\mathrm{H}}_2\mathrm{O}$$ 4

$${}^{*}\mathrm{NCO}\to 0.5{\mathrm{N}}_2+\mathrm{CO}$$ 5

$${}^{*}\mathrm{NCO}+3{\mathrm{H}}_2\to {\mathrm{CH}}_4+0.5{\mathrm{N}}_2+{\mathrm{H}}_2\mathrm{O}$$ 6

4.

Formation rate of the indicated products over the Ni/Al₂O₃ catalyst as a function of reaction temperature.

To confirm the validity of the proposed mechanism, the change of the DRIFTS bands attributed to *NCO (an intermediate of CH₄) and gas-phase CH₄ were investigated under H₂ flow. The experimental conditions were similar to those in Figure S8. First, to ensure the formation of *NCO on Ni/Al₂O₃, CO₂, and NH₃ gases (CO₂: 1.0 mL min^–1, NH₃: 2.6 mL min^–1, He: 6.4 mL min^–1) were passed over the catalyst at 300 °C for 20 min. Then, the supply of the reactant gases was stopped, and the sample chamber was purged with He gas (20 mL min^–1) while cooling to 100 °C. Once cooled, the temperature was increased from 100 to 400 °C with feeding of H₂ and He gas (H₂: 3.9 mL min^–1, He: 8.7 mL min^–1), and the transient response of the DRIFT spectrum was observed (Figure ). At 100 °C, bands attributed to *NCO adsorbed on Al₂O₃ and Ni were observed at 2233 and 2192 cm^–1, respectively. Bands attributed to HCOO* were observed at 1607 and 1587 cm^–1 and to CO₃* at 1455 cm^–1. At 200 °C, the intensity of the *NCO band was markedly decreased, and it almost disappeared at 300 °C. A band attributed to gas-phase CH₄ was observed at temperatures above 200 °C. The intensities of the CO₃* and HCOO* bands also decreased above 300 °C, indicating that among the adsorbed species, *NCO is more easily hydrogenated than CO₂. Taken together, these findings imply that *NCO is an intermediate of CH₄ in CO₂ + NH₃ methanation over Ni/Al₂O₃ (Scheme ). In addition, it was thought that N₂ was formed when CH₄ was formed from *NCO, as shown in eq (i.e., C–N bond cleavage occurs). However, IR-based analyses are unable to detect nitrogen gas because it is infrared-inactive. In addition, the sample quantity used for DRIFT apparatus (30 mg) was too small to observe desorption species. We therefore performed the following experiment to confirm N₂ formation using a fixed-bed flow reactor system (BELCAT-II), which could be filled with a larger amount of sample (100 mg) than the DRIFT apparatus. First, the Ni/Al₂O₃ catalyst was subjected to a reduction treatment (H₂ 33 mL min^–1/500 °C/1 h). Then, a gas mixture of CO₂, NH₃, and He (CO₂ 0.36%, NH₃ 0.93%, total 70 mL min^–1) was passed over the catalyst at 300 °C for 30 min to form *NCO. The reactant gases were then removed by switching to He gas (30 mL min^–1). After decreasing the temperature to 100 °C under He purging, the temperature was raised under flowing 5% H₂/He gas mixture (30 mL min^–1), and the desorbed species were observed with a mass spectrometer. Figure S11a shows the behavior of the m/z = 28 signal as a function of reaction temperature. The signal peak was observed at approximately 300 °C. This temperature was close to that at which the decrease of the *NCO band occurred in Figure . It should be noted that the m/z = 28 signal reflects not only N₂ but also CO. Figure shows that when CO₂ and NH₃ gases are passed over the Ni/Al₂O₃ catalyst, CO₂ adsorbed species such as CO*, HCOO*, and HCO₃* are formed in addition to *NCO. It is possible that CO formed from these CO₂ adsorbed species; however, when the pretreatment was performed with CO₂ and H₂ gas instead of CO₂ and NH₃ gas, no m/z = 28 signal was observed (Figure S11b), indicating that CO gas was not produced from the CO₂ adsorbed species and that the m/z = 28 signal in Figure S11a was derived only from N₂. Therefore, we concluded that CH₄ was formed from *NCO, and N₂ was formed by C–N bond cleavage.

5.

Changes of the DRIFTS bands attributed to *NCO in response to temperature (100–400 °C) under H₂ flow (H₂: 3.9 mL min^–1, He: 8.7 mL min^–1) over Ni/Al₂O₃ catalyst. Before measurement, the catalyst was contacted with CO₂ and NH₃ gas (CO₂: 1 mL min^–1, NH₃: 2.6 mL min^–1, He: 6.4 mL min^–1) at 300 °C and then purged with He gas.

2. Proposed Mechanism for CH₄ Formation during CO₂ + NH₃ Methanation over the Oxide-Supported Ni Catalysts.

Formation of *NCO was also confirmed under CO₂ and NH₃ flow over catalysts other than Ni/Al₂O₃ (i.e., Ni/SiO₂ and Ni/CeO₂; see Figure ). The influence of the oxide support on the stability and activity of *NCO is discussed in Section . The behavior of the rate of formation of the reaction products with respect to reaction temperature was similar for all of the catalysts; that is, the rates of formation of H₂, N₂, and CH₄ increased with increasing reaction temperature, while at high temperatures that of CO remained low and that of CH₄ increased (Figure S12). These results suggest that CO is produced over the catalysts by thermal decomposition of *NCO at low temperatures and *NCO is hydrogenated to CH₄ at high temperatures (Scheme ). From these results, it can be inferred that CH₄ is also formed from *NCO over the other catalysts. The material balances at low and high temperatures are shown in Scheme S1, with CH₄ produced by hydrogenation of *NCO at high temperatures.

9.

Operando DRIFTS analysis of *NCO adsorption to (a) Ni/SiO₂ catalyst and (b) Ni/CeO₂ catalyst. The band attributed to *NCO adsorption was monitored while the temperature was increased from 100 to 500 °C under He flow. Before the measurement, the catalyst was contacted with CO₂ and NH₃ gas (CO₂: 1 mL min^–1, NH₃: 2.6 mL min^–1, He: 6.4 mL min^–1) at 300 °C and then purged with He gas.

3.3. Effects of the Physicochemical Properties of the Catalysts

In Section , it was suggested that CH₄ is produced by different reaction mechanisms during CO₂ + NH₃ methanation compared with CO₂ + H₂ methanation and that *NCO acts as an intermediate. In the following sections, we discuss the catalyst properties that affect CO₂ + NH₃ methanation activity based on the proposed reaction mechanism.

The physicochemical properties of each catalyst were compared. First, the temperatures applied in the prereduction process with H₂ are described. In the present study, the reduction temperature was set to prevent sintering of the Ni particles while promoting the reduction of Ni species to increase activity. Therefore, the prereduction temperatures were changed for each catalyst based on the reduction behavior of the Ni species, as estimated by H₂-TPR (Figure S13). In the H₂-TPR profile of each catalyst (m/z = 18), several peaks attributed to H₂O (i.e., reduction) were observed. The heterogeneity of NiO particle size on the support and differences in the interaction between Ni species and the support are considered responsible for the appearance of multiple reduction peaks. For Ni/Al₂O₃, Ni/SiO₂, Ni/ZrO₂, and Ni/Y₂O₃, the reduction peaks all appeared at temperatures lower than 500 °C, indicating that the Ni species were reduced by H₂ treatment at low temperature. For Ni/CeO₂, two of three observed reduction peaks were observed at temperatures below 350 °C, indicating that most Ni species on CeO₂ are easily reduced. The ease of reduction of Ni species on CeO₂ supports has been reported in previous studies and is attributed to the weak interaction between Ni²⁺ and CeO₂. , The third reduction peak in the profile of Ni/CeO₂ was observed at 680 °C and was attributed to reduction of Ce⁴⁺ to Ce³⁺. For Ni/MgO, two peaks were observed, one at below 500 °C and one at above 800 °C, with the second peak indicating that the Ni species contained in Ni/MgO were very difficult to reduce. These Ni species are thought to originate from the formation of NiO-MgO solid. Based on these results, the H₂ prereduction temperature was set at 400 °C for Ni/CeO₂, 900 °C for Ni/MgO, and 500 °C for the other catalysts.

X-ray diffraction measurements of the catalysts reduced at the aforementioned temperatures are shown in Figure S14. Diffraction peaks attributed to the support oxides and Ni metals were observed for each sample, indicating the reduction of Ni species. The Ni K-edge X-ray absorption near edge structure (XANES) spectrum of each supported Ni catalyst after reduction is shown in Figure . The shapes of the spectra of Ni/SiO₂, Ni/CeO₂, Ni/Y₂O₃, and Ni/ZrO₂ were comparable to that of Ni foil, indicating that the Ni particles in these catalysts existed in the metallic state. The spectrum of Ni/Al₂O₃ was similar to that of Ni foil, but the pre-edge and white-line intensities were slightly closer to those of NiO compared with the spectra of the above four catalysts, suggesting the presence of a small amount of oxidized Ni even after reduction. From the H₂-TPR results for Ni/Al₂O₃ (Figure S13), some of the Ni species was reduced at temperatures above 500 °C, suggesting that some of the Ni species were strongly bound to the support as Ni–Al–O. , These Ni species are likely difficult to reduce during the reduction treatment. The shape of the spectrum of Ni/MgO was similar to that of NiO, indicating that most of the Ni species in this catalyst were not reduced. The strong interaction of Ni species with MgO likely induced formation of NiO-MgO solid, which prevented the reduction of Ni species.

6.

Ni K-edge XANES spectra of the six oxide-supported Ni catalysts along with those of Ni and NiO as references.

The specific surface area (SSA: m² g^–1), Ni crystallite size (nm), and Ni surface area (m_Ni ² g^–1) of each catalyst are shown in Table . The Ni surface area corresponds to the number of Ni sites on the catalyst surface. The Ni crystallite sizes of Ni/Al₂O₃ and Ni/Y₂O₃ were smaller than 10 nm, and their Ni surface areas were much larger than those of the other samples (Ni/Al₂O₃: 4.5 m_Ni ² g^–1, Ni/Y₂O₃: 3.4 m_Ni ² g^–1). The large number of Ni sites in Ni/Al₂O₃ can be attributed to moderate interaction between Ni²⁺ and the support based on the formation of Ni–Al–O and the large SSA of the support (180 m² g^–1). In contrast, Ni was highly dispersed on the Y₂O₃ support, even though the SSA of Y₂O₃ support (30 m² g^–1) was significantly smaller than that of Al₂O₃ support. A similar phenomenon was also reported by Okura et al. and was attributed to the microstructure of Y₂O₃ generated by treatment with the nitrate solution during catalyst preparation. In the present study, the SSA of the Y₂O₃ support after Ni loading (calcined: 54 m² g^–1) was larger than that before (30 m² g^–1), suggesting that the Y₂O₃ support surface was refined by contact with the nitrate solution during the preparation process, as previously reported. Thus, the characteristic microstructure of the Y₂O₃ support may have resulted in the large Ni surface area of Ni/Y₂O₃ compared with that of Ni/Al₂O₃. Even though the Ni crystallite size of Ni/MgO was smaller than that of Ni/Al₂O₃ and Ni/Y₂O₃, the Ni surface area of Ni/MgO was the smallest among the catalysts. This is because most of the Ni species in Ni/MgO were not reduced after the H₂ treatment, as indicated by the XANES analysis (Figure ). The SSA of the SiO₂ support was the largest (330 m² g^–1), which resulted in a large Ni surface area following that of Ni/Al₂O₃ and Ni/Y₂O₃. The Ni crystallite sizes of Ni/CeO₂ and Ni/ZrO₂ were larger than 20 nm, and thus Ni surface areas of these catalysts were smaller than those of Ni/Al₂O₃, Ni/SiO₂ and Ni/Y₂O₃. In Ni/CeO₂, the interaction between the Ni species and the support was weak, suggesting that sintering of Ni particles occurred despite the low reduction temperature (400 °C). In Ni/ZrO₂, the small SSA of the support (19 m² g^–1) likely facilitated aggregation of the Ni particles. The Ni surface areas of Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃, which exhibited high CO₂ + NH₃ methanation activity, were large among the catalysts, suggesting that the number of Ni sites contributes to CO₂ + NH₃ methanation activity. This result is attributed to the fact that the Ni surface is the main active site for CO₂ + NH₃ methanation (i.e., the formation of *NCO followed by its hydrogenation to CH₄ and NH₃ decomposition proceed on the Ni surface). Ni/Al₂O₃, Ni/SiO₂, and Ni/Y₂O₃ have high CO₂ + NH₃ methanation activity (Figure a), and Ni exists mainly in the metallic state in these catalysts (Figure ). This lends further support to the idea that metallic Ni functions as the main active site for this reaction. However, there are differences in the Ni surface areas of the above three catalysts, in particular the Ni surface area of Ni/SiO₂ (2.1 m_Ni ² g^–1) was about half that of Ni/Al₂O₃ (4.5 m_Ni ² g^–1). Therefore, CO₂ + NH₃ methanation activity cannot be solely dependent on the number of Ni sites.

1. Physicochemical Properties of the Six Oxide-Supported Ni Catalysts.

		specific surface area (m2 g–1)
catalyst	reduction temperature (°C)	support	calcined	after reduction	Ni crystallite size (nm)	Ni surface area (mNi 2 g–1)
Ni/Al2O3	500	180	139	129	9.0	4.5
Ni/MgO	900	28	73	42	7.5	0.7
Ni/SiO2	500	330	213	225	13	2.1
Ni/CeO2	400	73	49	57	29	1.8
Ni/ZrO2	500	19	15	17	23	1.5
Ni/Y2O3	500	30	54	57	8.9	3.4

3.4. Role of NH₃ Decomposition in the Overall Reaction Mechanism

As shown in Figure d, CH₄ yield was correlated with NH₃ conversion during CO₂ + NH₃ methanation. The indication is that NH₃ decomposition activity is an important factor for determining the activity of oxide-supported Ni catalysts in this reaction. Therefore, the influence of the surface properties of the support on NH₃ decomposition activity and their role in the overall reaction mechanism were examined in more detail.

First, we focused on intrinsic NH₃ decomposition activity (i.e., NH₃ decomposition activity in an atmosphere where CO₂ is not present). It is reported that the basic properties of the support affect NH₃ decomposition activity. , We therefore performed CO₂-TPD measurements to investigate the basic properties of each catalyst (Figure ). For Ni/CeO₂ and Ni/Y₂O₃, CO₂ desorption was observed at temperatures above 600 °C, indicating that these catalysts have sites with high basic strength. In NH₃ decomposition over oxide-supported Ni catalysts, the associative desorption of adsorbed N is considered the rate-determining step. , The CO₂-TPD profiles also showed that Ni/CeO₂ and Ni/Y₂O₃ have sites that exhibit superior electron-donating ability because of their high basic strength. Thus, it is likely that the superior electron-donating ability of these catalysts promoted the associative desorption of adsorbed N, resulting in the high NH₃ decomposition activity shown in Figure S7. In addition, Ni/Y₂O₃ was shown to have a larger Ni surface area (Table ). Therefore, it is likely that Ni/Y₂O₃ exhibited a much higher NH₃ decomposition activity compared with that of the other catalysts due to both its high electron-donating ability and large number of active sites. For Ni/Al₂O₃, Ni/SiO₂, and Ni/ZrO₂, CO₂ desorption was observed below 160 °C. The NH₃-decomposition activity of these catalysts was inferior to that of Ni/CeO₂ and Ni/Y₂O₃, presumably because of the low basic strength. For Ni/MgO, CO₂ desorption was observed at around 200 °C. The indication was that this catalyst has sites with higher basic strength than those of Ni/Al₂O₃, Ni/SiO₂, and Ni/ZrO₂. However, the Ni surface area of Ni/MgO was the smallest among the catalysts. The suggestion was that the low number of active sites was the reason for the low NH₃-decomposition activity of Ni/MgO (Figure S7).

7.

CO₂ temperature-programed desorption profiles of the six oxide-supported Ni catalysts.

Next, the NH₃ conversion at 500 °C afforded by each catalyst was compared in the presence or absence of CO₂ (Figure ). For Ni/Al₂O₃ and Ni/SiO₂, NH₃ conversion in the presence of CO₂ was higher than that without CO₂; whereas for the other catalysts, the opposite was observed.

8.

NH₃ conversion of the six oxide-supported Ni catalysts at 500 °C in the presence (NH₃ conv. with CO₂) or absence of CO₂ (NH₃ conv. without CO₂).

NH₃ conversion as a function of reaction temperature was also compared in the presence or absence of CO₂ (Figure S15). For Ni/Al₂O₃ and Ni/SiO₂, NH₃ conversion in the presence of CO₂ exceeded that without CO₂ at 500 and 450 °C, respectively, with CH₄ yield also increasing. In CO₂ + NH₃ methanation, NH₃ does not simply decompose, some of it also reacts with CO₂ to form *NCO (an intermediate of CH₄, see Scheme S1b). Also, from eq , N₂ is produced when *NCO is hydrogenated to CH₄. Such N₂ formation contributes to increasing NH₃ conversion (see calculation of NH₃ conversion in Section .). It should thus be noted that in CO₂ + NH₃ methanation, NH₃ conversion reflects not only the NH₃ decomposition rate but also the conversion rate of *NCO to CH₄. Taking this into consideration, the results shown in Figure S15 suggest that the rate of formation of N₂ associated with hydrogenation of *NCO to CH₄ is faster than the rate of formation of N₂ associated with NH₃ decomposition at 500 °C over Ni/Al₂O₃ and Ni/SiO₂. In other words, these two catalysts showed superior reactivity of *NCO with H₂. In contrast, NH₃ conversion in the presence of CO₂ over the other catalysts did not exceed that without CO₂ even at the high temperatures at which the CH₄ yield increased. In particular, NH₃ conversion in the presence of CO₂ was much lower than that without CO₂ over Ni/CeO₂ and Ni/Y₂O₃ at any temperature. The indication was that the reactivity of *NCO with H₂ is low over Ni/CeO₂ and Ni/Y₂O₃. Alternatively, it could be that NH₃-decomposition activity over Ni/CeO₂ and Ni/Y₂O₃ is suppressed in the presence of CO₂. Thus, Ni/Al₂O₃ and Ni/SiO₂ showed high CO₂ + NH₃ methanation activity due to the superior reactivity of *NCO, an intermediate of CH₄, rather than their intrinsic NH₃-decomposition activity. In contrast, for Ni/CeO₂ and Ni/Y₂O₃, either the reactivity of *NCO was low or their intrinsic NH₃ decomposition activity was lowered (see Section for a more detailed discussion of these topics). However, Ni/Y₂O₃ showed much higher NH₃-decomposition activity than the other catalysts (Figure S7), and despite the low reactivity of *NCO and the lowered NH₃-decomposition activity, Ni/Y₂O₃ exhibited CO₂ + NH₃ methanation activity comparable to that of Ni/Al₂O₃ and Ni/SiO₂ because of its superior ability to supply H₂ to *NCO based on its significantly high intrinsic NH₃-decomposition activity.

3.5. Stability and Activity of *NCO over the Catalysts

Finally, we discuss the properties that affect the reactivity of *NCO with H₂. CO₂-TPD analysis indicated the possibility that the reactivity of *NCO over Ni/CeO₂ and Ni/Y₂O₃, which have sites with high basic strength, is low (Figure ). In contrast, the reactivity of *NCO over Ni/Al₂O₃ and Ni/SiO₂, which have sites with low basic strength, is presumably high. Therefore, the basic strength of the catalyst may affect not only its intrinsic NH₃ decomposition activity but also the reactivity of *NCO. The reactivity of *NCO is assumed to correlate with the strength of its adsorption on the catalyst. Therefore, the strength of adsorption of *NCO was investigated by DRIFTS analysis for two catalysts with different basic strengths: Ni/CeO₂, which has high basic strength, and Ni/SiO₂, which has low basic strength. The reactant gas (CO₂: 1.0 mL min^–1, NH₃: 2.6 mL min^–1, He: 6.4 mL min^–1) was first contacted with each catalyst at 300 °C to induce the formation of *NCO. Then, the supply of reactant gas was stopped, He gas was supplied, and the temperature was cooled to 100 °C. Finally, the transient response of the DRIFT spectrum was observed while heating the sample chamber from 100 to 500 °C under He gas flow (20 mL min^–1). The strength of adsorption of *NCO was assessed based on its decomposition temperature (Figure ). At 100 °C, bands assigned to *NCO adsorbed on Ni were observed for both catalysts. For Ni/SiO₂, the *NCO (band at 2205 cm^–1) was almost fully decomposed at 300 °C. In contrast, for Ni/CeO₂, the *NCO clearly remained on the Ni (2174 cm^–1) at 300 °C, was slightly decomposed at 400 °C, and had almost disappeared at 500 °C. Thus, *NCO was adsorbed much more strongly on Ni/CeO₂ than it was on Ni/SiO₂. The strength of *NCO adsorption was also examined over Ni/Y₂O₃, which had high basic strength similar to that of Ni/CeO₂ (Figure S16). Over Ni/Y₂O₃, a small amount of adsorbed *NCO (2192 cm^–1) was observed even at 500 °C. Together, these results suggest that the basic strength of the catalyst affects the strength of adsorption of *NCO, and that the higher the basic strength, the higher the strength of adsorption of *NCO. The difference in the strength of adsorption of *NCO on the different catalysts may be due to the fact that this species is located near the Ni–support interface, where it is more susceptible to interactions with the support. The wavenumber of the band attributed to *NCO adsorption on Ni varied with the support (Ni/SiO₂: 2205 cm^–1, Ni/CeO₂: 2174 cm^–1, Ni/Y₂O₃: 2192 cm^–1), which supports the idea that the adsorption of *NCO is influenced by support interactions. It is possible that the use of supports with high basic strength, such as CeO₂ and Y₂O₃, produces sites at the Ni–support interface that strongly stabilize *NCO, leading to a decrease of the reactivity of *NCO. In addition, the strong adsorption of *NCO at the Ni–support interface leads to a lowering of the electron-donating ability of the support, because many electrons are allocated to adsorption of *NCO. The result is that it is difficult for associative desorption of adsorbed N during NH₃ decomposition to proceed, and the NH₃-decomposition activity is lowered. This may be an additional factor inhibiting the CO₂ + NH₃ methanation activity of catalysts with high basic strength, such as Ni/CeO₂ and Ni/Y₂O₃. Thus, it is likely that catalysts with low basic strength, such as Ni/Al₂O₃ and Ni/SiO₂, are advantageous because they increase the reactivity of *NCO with H₂ and maintain the ability of the support to donate electrons to the active metal. We compared the turnover frequency (TOF) of the catalysts calculated from the amount of surface Ni estimated via CO-pulse measurements and the rate of formation of CH₄ at 500 °C (Figure S17). The fact that Ni/SiO₂ showed the highest TOF suggested that the reactivity of *NCO on Ni/SiO₂ was very high.

Figure S18 shows the results of a durability test for the Ni/Al₂O₃ catalyst at 500 °C for 60 h. Over 60 h, the CH₄ yield and NH₃ conversion gradually decreased, while the CO yield increased, indicating degradation of the catalyst. To examine the cause of this catalyst degradation, the oxidation state of Ni before and after the durability test was investigated by XAFS measurement. The Ni K-edge XANES spectrum of the Ni/Al₂O₃ catalyst before and after the durability test is shown in Figure S19a. After the durability test, the spectrum approached that of Ni foil, indicating that the reduction of Ni had progressed. This is thought to be because the reducing gas (NH₃, H₂, CO, etc.) was able to pass over the catalyst for a long period of time at the same temperature as the reduction treatment (H₂ 50 mL min^–1, 500 °C, 1 h). Therefore, it can be said that the degradation of the catalyst is not due to the oxidation of Ni. Figure S19b shows the Fourier transforms of the k ³-weighted EXAFS oscillations before and after the stability test. After the durability test, the magnitude of the Fourier transform peak assigned to Ni–Ni bond increased, suggesting that Ni sintering occurred. Based on these results, we carried out STEM observations of the catalyst before and after the durability test (Figure S20). The particle size distributions were obtained by analyzing the sizes of 100 particles. Before the durability test, very small particles of less than 5 nm were observed in addition to Ni particles of 10 nm or larger (Figure S20a). After the durability test, the proportion of particles measuring less than 5 nm had decreased, while the proportion of particles measuring between 5 and 10 nm had increased (Figure S20b). After the test, the median diameter of the Ni particles had increased from 7.0 to 9.7 nm. The indication is that sintering of the nickel particles had occurred, resulting in catalyst deactivation. Carbon (C) deposition is another factor that can contribute to catalyst degradation in CO₂ + H₂ methanation. To investigate whether C deposition occurs in CO₂ + NH₃ methanation, elemental analysis of Ni/Al₂O₃ after reduction and durability testing was performed. An automatic elemental analyzer (2400IICHNS/O type, PerkinElmer, Japan) was used for the measurement. The elemental profiles (C, H and N) of the catalyst after reduction and durability testing are shown in Table S3. After the durability test, the C, H, and N content in the catalyst were slightly increased compared with after reduction. We presume that the increase in the C content was not due to carbon deposition (e.g., coke or other solid carbon compounds), but rather to the effects of remaining adsorbed species such as *NCO and HCOO*. The implication is that preventing Ni from sintering is the most important factor for improving catalyst durability.

In summary, the present results indicate that the reactivity of *NCO with H₂ and the rate of formation of H₂ (corresponding to intrinsic NH₃ decomposition activity) are both important for CO₂ + NH₃ methanation. Also, the basic strength of the catalyst likely affects these two activity control factors, with supports with low basic strength being most effective for the former and supports with high basic strength being most effective for the latter. Among the present catalysts, Ni/Al₂O₃ and Ni/SiO₂ exhibited high CO₂ + NH₃ methanation activity because they afforded a superior reactivity of *NCO with H₂. This is presumably because the basic strength of the Al₂O₃ and SiO₂ supports is low, and no sites for strong adsorption of *NCO are formed at the Ni–support interface. In contrast, the Y₂O₃ and CeO₂ supports had high basic strength, which created sites for strong adsorption of *NCO at the Ni–support interface after Ni loading; as a result, the reactivity of *NCO with H₂ was decreased over these catalysts. It is also possible that the strong adsorption of *NCO to the basic sites formed at the Ni–support interface reduced the electron-donating ability of the support, which is important for promoting NH₃ decomposition. However, Ni/Y₂O₃ had significantly higher intrinsic NH₃-decomposition activity than the other catalysts due to its high basic strength and large Ni surface area. As a result, despite the low reactivity of *NCO and the lowered electron-donating ability, Ni/Y₂O₃ showed CO₂ + NH₃ methanation activity comparable to that of Ni/Al₂O₃ and Ni/SiO₂. Based on the above results, we conclude that a support with moderate basic strength would most likely provide a suitable balance between NH₃ decomposition activity and*NCO reactivity (Scheme ). Therefore, optimizing the basic strength of the catalyst will be important for obtaining catalysts with high CO₂ + NH₃ methanation activity.

3. Relationship between Activity Controllable Factors (NH₃ Decomposition and *NCO Destabilization) and Support Basicity for the Oxide-Supported Ni Catalysts .

^a RDS: rate-determining step.

4. Conclusion

To obtain design guidelines for the development of highly active catalysts for CO₂ + NH₃ methanation, we investigated the activity and physicochemical properties of six oxide-supported Ni catalysts and examined the reaction mechanism by operando DRIFTS analysis. Operando DRIFTS measurements showed that *NCO was formed on the catalysts at low temperatures under an atmosphere of CO₂ and NH₃. A CO species adsorbed on Ni (CO*) was observed under an atmosphere of CO₂ + H₂, but this species did not form under the atmosphere of CO₂ + NH₃ because *NCO occupied the Ni surface. During CO₂ + H₂ methanation, CH₄ was formed with CO* as an intermediate, whereas during CO₂ + NH₃ methanation, CH₄ was formed with *NCO as an intermediate. These results suggest that the reaction mechanism for CH₄ formation in CO₂ + NH₃ methanation is different from that in CO₂ + H₂ methanation. The reactivity of *NCO with H₂ over the catalysts is an important factor underlying CO₂ + NH₃ methanation activity, and it differed depending on the type of support. *NCO was strongly adsorbed on Ni/CeO₂ and Ni/Y₂O₃, which both had high basic strength. Therefore, the ease of *NCO hydrogenation to CH₄ did not depend solely on the ability to form H₂. The fact that NH₃ decomposition activity was much higher over Ni/Y₂O₃ than over the other catalysts resulted in a faster rate of CH₄ formation. Ni/Al₂O₃ and Ni/SiO₂ had lower NH₃ decomposition activity than Ni/Y₂O₃; however, because of the weak adsorption of *NCO on these catalysts, they exhibited superior reactivity of *NCO with H₂, resulting in high CO₂ + NH₃ methanation activity. The present results also suggested that NH₃ decomposition activity and *NCO reactivity are affected by the basic strength of the support. Therefore, to develop more active catalysts, it will be necessary to control the basicity of the support by combining oxides with two kinds of basic strength (strong and weak) and to design materials that afford both high NH₃ decomposition activity and high *NCO reactivity. One approach for the development of oxide-supported Ni catalysts with more efficient CH₄ production from CO₂ and NH₃ is the use of morphology control to increase the metal dispersion on the support with optimized basic properties. In short, our present findings provide not only design guidelines for the development of highly active catalysts for CO₂ + NH₃ methanation but insights into the CO₂ conversion reaction mechanism involving carbon, hydrogen, and nitrogen.

Although careful verification of total carbon footprint and economic feasibility will be needed in the future, CO₂ + NH₃ methanation has great potential to contribute to the realization of carbon-neutral fuel production processes. Currently, NH₃ is synthesized by an energy-intensive process, and CH₄ is a low-value compound compared to NH₃. Therefore, to achieve effective CH₄ production, it will be necessary to develop related technologies so that green H₂ and green NH₃ can be produced in large quantities and at an economically competitive cost. Separation of CH₄ from gas mixtures containing N₂ at high concentration is also an important issue. Further improvement of catalyst activity and synthesis process will contribute to achieving these goals. The present findings regarding catalyst design are expected to facilitate realization of a practical CO₂ capture and utilization process, which is a potentially important step toward realization of a more sustainable and carbon-neutral society.

The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/jacsau.5c01058.

• Support preparation; gas conditions used in the activity test; XRF measurement; CO₂ conversion estimated from unreacted CO₂; dependence of CH₄, CO yields and NH₃ conversion on SV; table for summary of CO₂ + NH₃ methanation activity over various oxide-supported catalysts; the activity of CO₂ + H₂ methanation; analysis of N-containing compounds; reproducibility of CO₂ + NH₃ methanation activity test; the activity of NH₃ decomposition; transient response of CO₂ adsorbed species under H₂ flow; DRIFT spectrum for support under a CO₂ + NH₃ methanation atmosphere; investigation of N₂ formation from catalyst surface after contact with CO₂ and NH₃; reaction product (CH₄, CO, N₂, H₂) formation rates; material balance of CO₂ + NH₃ methanation; H₂ temperature-programed reduction profiles; X-ray diffraction patterns of the catalysts; NH₃ conversions vs reaction temperature in the presence or absence of CO₂; DRIFTS analysis of *NCO adsorption on Ni/Y₂O₃; TOFs for CO₂ + NH₃ methanation; stability test and the characterizations of the catalysts after the stability test (PDF)

K.S. and K.N. designed and coordinated this project. Y.U. conducted all experiments and analyzed the data. Y.U., K.N., and K.S. cowrote the manuscript.

The authors declare no competing financial interest.