*2.2. Catalyst Characterization*

The chemical composition of the catalysts was determined by an inductively coupled plasma optical emission spectrometer (ICP-OES). 0.05 g of the sample was dissolved in a mixture of nitric acid and hydrogen peroxide. The solid residue was filtrated and melted in sodium peroxide or sodium borate. The melt was heated in hydrochloric acid to extract the elements in the sample. The obtained liquids were mixed thoroughly and adjusted to a specific volume in a volumetric flask. The prepared liquid was subjected to ICP-OES (Shimadzu, Kyoto, Japan, ICPS-8100). Specific X-ray of λ = 413.765 nm was used for Ce, and those of λ = 280.270 and 383.231 nm were used for Mg. The ICP-OES detector was calibrated using reference liquids of the elements before the measurements. The composition of each catalyst was calculated from ICP-OES results assuming both Ce and Mg exist as oxides, CeO2 and MgO, in the catalyst.

The surface areas of the catalysts were determined using the Brunauer–Emmett–Teller (BET) method from the N2 adsorption isotherm at the temperature of liquid nitrogen (MicrotracBEL, Osaka, Japan, BELSORP MINI X). TEM images of the catalysts were taken as bright-field images with a transmission electron microscope (Thermo Fisher Scientific, MA, USA, Tecnai G2) with an accelerating voltage of 200 kV. Each sample was mounted on a carbon grid and measured on a single-tilt holder. The average particle sizes of the CeO2 were determined using the procedure described in the supporting information of our previous paper [44]. The projected areas of more than 100 nanoparticles on the TEM images were determined with image analysis software. Then, the sphere equivalent diameter was calculated for each particle. The arithmetic mean was calculated using an obtained frequency distribution of the diameter. The total surface area of the CeO2 particles was also estimated for each catalyst based on the frequency distribution. The morphology of the catalysts was also observed with a scanning transmission electron microscope (STEM). STEM images of the catalysts were obtained as bright-field images with a scanning electron microscope (Hitachi, Tokyo, Japan, SU9000) with an accelerating voltage of 30 kV. Energydispersive X-ray spectroscopy (EDS) mapping images were taken along with STEM images with an EDS detector (Oxford Instruments, Tokyo, Japan, X-max 100LE).

The redox properties of the catalysts were assessed by temperature-programmed reduction (TPR) and temperature-programmed oxidation (TPO) studies. TPR of the samples was carried out under a flow of 5% H2/Ar. All measurements were conducted using an

automated catalyst analyzer (MicrotracBEL, Osaka, Japan, BELCAT II) equipped with a thermal conductivity detector (TCD). For the TPR measurements, 0.05 g of the sample were placed on the bottom of a quartz tube and heated to 500 ◦C under an Ar flow to remove impurities. The sample was then quenched to 30 ◦C under a flow of Ar. TPR was conducted from 30 ◦C to 900 ◦C under a constant 30 cm3/min flow of 5% H2/Ar gas at a heating rate of 5 ◦C/min. The TPO measurements were conducted using a heating chamber (PIKE Technologies, WI, USA, DiffusIR) equipped with a quadrupole mass spectrometer (QMS) (Pfeiffer Vacuum, Aßlar, Germanny, OmniStar). CO2 is used as an oxidant gas. For each measurement, 0.0275 g of the sample were heated in a heating chamber to 550 ◦C under a flow of Ar and reduced for 20 min under a constant 50 cm3/min flow of 5% H2/Ar gas. The sample was then cooled to 100 ◦C under a flow of Ar. TPO was conducted at a heating rate of 20 ◦C/min from 100 to 740 ◦C under a constant 50 cm3/min flow of 5% CO2/Ar gas. The outlet gas was monitored by the QMS. The QMS signals corresponding to the m/z values of 2 (H2), 28 (CO), 40 (Ar), and 44 (CO2) were monitored throughout the reaction. The measurement cycle was ~1.0 s at a dwelling time of 0.2 s for each m/z. The gas composition was determined using the relative intensities of the signals to that of m/z 40 (Ar). The total gas concentration was normalized to 100% for each cycle. The QMS system was calibrated before the measurements. Standard gases were used to calibrate the m/z values of 2 (H2), 28 (CO), 40 (Ar), and 44 (CO2). The CO concentration was calculated by subtracting the fragmentation of CO2 and background. The fragmentation ratio of CO2 to m/z of 44 and 28 was determined using a flow of CO2/Ar.

The crystallographic structure of the catalysts was determined using in situ X-ray diffraction (XRD) measurements with an XRD instrument (Rigaku, Tokyo, Japn, RINT-TTR III) equipped with a Co X-ray source and an Fe filter. The scanning range was set from 2*θ* = 20 to 80◦ with a 0.02◦ step angle at a scanning rate of 40◦/min. The X-ray tube voltage and the current were 45 kV and 200 mA, respectively. The sample catalysts were crushed into powder and mounted on an infrared heating attachment (Rigaku, Tokyo, Japan, Reactor X). The infrared heating attachment was purged by a 5% H2/N2 flow of 150 cm3/min under the pressure of 0.1 Mpa. H2 was supplied from an H2 generator (Parker Hannifin, OH, USA, H2PEM-260) and diluted with N2 supplied from a gas cylinder (Tokyo Koatsu, Tokyo, Japan, >99.99995%). The sample was heated up to 900 ◦C at a heating rate of 5 ◦C/min under continuous gas flow. The scanning was performed every 10 ◦C from 200 to 900 ◦C. XRD measurements were also performed under ambient air with an XRD instrument (Rigaku, Tokyo, Japan, RINT-TTR III) equipped with a Cu X-ray source and an Ni filter. The scanning range was from 2*θ* = 20 to 70◦ with a 0.02◦ step angle at a scanning rate of 1◦/min at room temperature.

The oxidation state of Ce was measured on an X-ray photoelectron spectroscopy (XPS) analyzer (ULVAC-Phi, Kanagawa, Japan, Quantum-2000). The XPS analyzer was equipped with a monochromated Al X-ray source and a charge neutralizer. The pass energy and the recording step were controlled to 29.35 eV and 0.125 eV, respectively. The binding energy of an isolated u"' peak of Ce3d3/2 was adjusted to 916.70 eV to correct the peak shift derived from the charge-up of the catalysts [45,46]. C1s peak of adventitious carbon was not used for the correction due to the weak intensity of the peak and the overlap with the peaks of carbonates and Ce4s [47,48].

## *2.3. Dry Reforming Reaction*

The obtained catalysts of particle size 500 to 750 μm were subjected to the dry reforming reaction. All the test reactions were conducted at ambient pressure for 6 h at 800 ◦C using a tubular flow reactor The composition of the gas was fixed to 25% CH4, 25% CO2, and Ar as balance at a flow rate of 100 cm3/min. The amount of catalyst was 0.1 g for all the reactions, of which space velocity was 60,000 cm<sup>3</sup> hr−<sup>1</sup> g-cat−1. The inlet and outlet gas compositions were analyzed by a gas chromatograph (Shimadzu, Kyoto, Japan, GC-2014), employing Ar as the carrier gas. The absence of N2 eliminated peak overlap between CO

and the balance gas, which allowed the concentration of the product gases to be estimated correctly even at low CH4 conversions.

The CH4 conversions were calculated using a procedure described in our previous report [49], assuming a two-step reaction shown in Equations (1) and (2). We used the reaction rates for Equations (1) and (2) as parameters to reproduce the outlet gas composition by the least-squares method. The calculated reaction rates for Equation (1) were used as the CH4 conversion rates. All the calculations were performed ignoring the carbon deposition on the catalysts and the C2 species in the outlet gas (<50 ppm). Therefore, the total flow of CH4 and CO2 in the inlet gas was equal to that of CO, CO2, and CH4 in the outlet gas during the calculation.

$$\text{CH}\_4 + \text{CO}\_2 \rightarrow 2\text{CO} + 2\text{H}\_2 \tag{1}$$

$$\text{H}\_2 + \text{CO}\_2 \rightarrow \text{CO} + \text{H}\_2\text{O} \tag{2}$$

### **3. Results and Discussion**

*3.1. Characterization of the Catalysts*

The content of CeO2 in the catalysts was calculated based on the results of ICP-OES. The obtained values were compared to the nominal values calculated from the amount of nitrates used. The nominal values and the experimental values are almost identical to each other (Table 1). This result suggests that all the catalysts were prepared at the nominal chemical composition as intended.


**Table 1.** Morphological properties of the prepared catalysts.

\* Contents of CeO2 and MgO in the catalysts were experimentally determined by ICP-OES. Sum of CeO2 and MgO was normalized to 100%. \*\* Arithmetic averages were calculated from TEM measurement results. Standard deviations were described after "±" for each average diameter. The diameter of pure CeO2 was estimated as an area-weighted average assuming the particles were spheres. \*\*\* Calculated assuming that CeO2 particles are hemisphere attaching their flat planes on MgO.

> The crystallographic structure of the prepared catalysts was assessed by XRD. No peaks other than CeO2 and MgO were observed across all of the XRD patterns (Figure 1). The diffraction patterns of CeO2 were weaker for CeO2/MgO catalysts than for that of pure CeO2 because the content of CeO2 in CeO2/MgO was less than 20% in the mass ratio (Table 1). Further, no clear peak shift was observed for peaks of CeO2 and MgO, suggesting that the formation of the solid solution is small. These results are in good accordance with previous studies [31,34]. The amount of CeO2 and MgO dissolving with each other was negligible even at 1500 ◦C. On the other hand, a clear difference was observed between the samples in the peak shape. The broader peaks were observed for the samples prepared in the dry atmosphere than those prepared in the ambient air, suggesting the finer MgO and CeO2 particles of the samples prepared in the dry atmosphere.

> The size of the CeO2 nanoparticles was also compared by TEM measurement (Figure 2). Small particles stuck on the large particles were assigned as CeO2 particles based on EDSmapping results (Figure S1). Figure 3 shows the frequency distribution of the CeO2 diameter. The results are summarized in Table 1. Small CeO2 clusters (<3 nm), as well as larger nanoparticles (>3 nm), were observed in all the CeO2/MgO catalysts as reported by Tinoco et al. [32]. The average diameter of CeO2 increased as the content of CeO2 in the

catalysts increased; 0.05-CeO2/MgO(air) and 0.05-CeO2/MgO(dry) contained larger CeO2 nanoparticles than 0.01-CeO2/MgO(air) and 0.01-CeO2/MgO(dry), respectively (Table 1). Meanwhile, smaller CeO2 nanoparticles were observed on the catalysts prepared under the dry atmosphere. The 0.01-CeO2/MgO(dry) sample contained 2.6-nm CeO2 on average, which is smaller than 5.3-nm CeO2 for 0.01-CeO2/MgO(air) (Table 1). The results of TEM were in good accordance with the results of XRD; the volume-weighted average of CeO2 nanoparticles was comparable for both TEM and XRD results (Table S2).

**Figure 1.** XRD patterns of (**a**) 0.01-CeO2/MgO and (**b**) 0.05-CeO2/MgO catalysts compared to that of pure CeO2. A Cu X-ray source and an Ni filter were used. MgO and CeO2 patterns were assigned based on the references [56,57].

The difference in the size of the CeO2 nanoparticles was attributed to the difference in the BET area. As shown in Table 1, the BET area of MgO(dry) was approximately three times larger than that of MgO(air). MgO(dry) contained small cubic grains caused by the thermal decomposition of Mg(OH)2 (Figure S2) [42]. In contrast, larger octahedral grains were observed in MgO(air) instead of in cubic grains. This difference is ascribed to the water vapor contained in the ambient air. The water vapor accelerates the sintering of MgO of <5 nm in diameter [41,42,50]. Further, the stability of each facet of MgO depends on the humidity of the atmosphere; the {100} face is favored under the dry atmosphere, but the {111} face is under the humid atmosphere [38,51]. The morphological change of MgO due to the humidity caused the sintering of CeO2 nanoparticles during the calcination of the samples at 800 ◦C.

The precipitous drop in the BET surface areas of 0.01- and 0.05-CeO2/MgO(dry) compared to MgO(dry) was attributed to the condensed nitrate solutions formed during the drying process of the impregnation. During drying, the surface of MgO is covered by Mg(OH)2 since the acetone solution of Ce(NO3)3 contains water; Ce(NO3)3 6H2O was used as a precursor. Further, the acetone solution contains a small amount of water as an impurity [52]. Therefore, the equilibrium of Equation (3) is present during the drying process.

$$2\text{Ce}(\text{NO}\_3)\_3 + 3\text{Mg(OH)}\_2 \rightleftharpoons 2\text{Ce}(\text{OH})\_3 + 3\text{Mg(NO}\_3)\_2 \tag{3}$$

Mg(NO3)2 is soluble both in water and acetone [53,54]. Once the solution of Mg(NO3)2 formed, heterogeneous nucleation of Mg(NO3)2 and Ostwald ripening of MgO would proceed as reported for other oxides [55]. Therefore, the condensed nitrate solution would solve the MgO surface and increase its grain size, which reduced the surface area of MgO.

This result suggests that it is possible to increase the surface area by further optimizing the combination of Ce precursors and solvents.

**Figure 2.** TEM images of (**a**) 0.01-CeO2/MgO(air), (**b**) 0.05-CeO2/MgO(air), (**c**) 0.01-CeO2/MgO(dry), (**d**) 0.05-CeO2/MgO(dry), and (**e**) CeO2. STEM-EDS of the samples are shown in Figure S1.

**Figure 3.** Frequency distribution of the CeO2 diameter in (**a**) 0.01-CeO2/MgO, and (**b**) 0.05-CeO2/MgO.

## *3.2. Mobility of Oxygen*

The mobility of oxygen species on the catalysts was assessed by the TPR measurement (Figure 4). MgO had no peaks because it is irreducible under the measurement conditions (Figure S3). All the other catalysts showed peaks at ~500 and ~800 ◦C, but the relative intensity of these peaks varied depending on the catalyst. The former and latter peaks are ascribed to the surface capping oxygen and the bulk lattice oxygen of CeO2, respectively [1]. Notably, the intensity of the peaks at ~800 ◦C lowered as the CeO2 content decreased while the intensity of the peaks at ~500 ◦C remained comparable, suggesting that the low content of CeO2 led to finer CeO2 nanoparticles (Figure 4a,b). This result is in good agreemen<sup>t</sup> with the results of the TEM measurement (Table 1, Figures 2 and 3). Further, the intensity of the peaks at ~800 ◦C was lower for CeO2/MgO(dry) than for CeO2/MgO(air) for both contents of CeO2. These results sugges<sup>t</sup> that the surface capping oxygen was predominant for the finer CeO2 nanoparticles prepared under the dry atmosphere due to its larger surface area of CeO2 (Table 1). In particular, 0.01-CeO2/MgO(dry) had no peaks at ~800 ◦C, suggesting that both the surface capping oxygen and the lattice oxygen were removed at ~500 ◦C. This result is attributed to the fine CeO2 of d = 2.6 nm in 0.01-CeO2/MgO(dry) (Table 1). Such small CeO2 nanoparticles have high reducibility since their unit cell expanded [13,20]. The lattice oxygen of such CeO2 nanoparticles would diffuse ~1 nm from the center of the CeO2 nanoparticles to the surface under the reducing atmosphere.

In addition to the intensity of the reduction peak at ~500 ◦C, the shape of the peak also changed depending on the preparation condition; a small shoulder appeared at ~350 ◦C for the samples prepared under dry conditions. This low-temperature peak was attributed to the size-dependent orientation of CeO2 nanoparticles on the MgO surface. Reportedly, the crystallographic orientation of CeO2 nanoparticles varies depending on the size of CeO2; CeO2 of 1–3 nm in diameter faces its {111} surface to the {111} surface of MgO, but CeO2 of 10–20 nm in diameter faces its {100} surface to the {111} surface of MgO [32]. These studies sugges<sup>t</sup> that preferential faceting of CeO2 nanoparticles depends on the size of CeO2. Further, the reducibility of the CeO2 surface strongly depends on their faceting [23]; the reduction peak position shifted from ~500 ◦C for conventional CeO2 to ~300 ◦C for nanorod CeO2. Therefore, we considered that the reducibility of the surface capping oxygen was modulated by the difference in the faceting of CeO2 nanoparticles caused by the size variation.

**Figure 4.** H2-TPR profiles of (**a**) 0.01-CeO2/MgO and (**b**) 0.05-CeO2/MgO catalysts compared to that of pure CeO2.

Redox properties of the catalysts were further assessed by TPO using CO2 as an oxidant. The catalysts were reduced in a 5%H2/Ar flow at 550 ◦C before the TPO measurements. As shown in Figure 5, the formation of CO starts at ~350 ◦C on all the catalysts. Further, the peak intensity is ~1.5 times higher on 0.05-CeO2/MgO(dry) compared to those on 0.05-CeO2/MgO(air) and pure CeO2. These results are consistent with the H2-TPR (Figure 4b); the reduction peak at ~500 ◦C on 0.05-CeO2/MgO(dry) is larger than those of the other catalysts. CO2 refilled the oxygen vacancies on the surface that formed during the reduction in an H2 flow. This result also demonstrated the improved oxygen mobility of 0.05-CeO2/MgO(dry) than those of 0.05-CeO2/MgO(air) and pure CeO2. Further, the results shown in Figures 4 and 5 sugges<sup>t</sup> that CeO2/MgO(dry) works as a catalyst for chemical looping combustion (CLC) [6–9]. The CLC proceeds as a two-step reaction; the catalyst is subjected to reducing conditions and then re-oxidized by CO2 or H2O to produce CO or H2, which is identical to the reaction in Figure 5.

**Figure 5.** CO2-TPO profiles of the reduced catalysts. The catalysts were heated at 20 ◦C/min in a 5% CO2/Ar flow after the reduction at 550 ◦C in a 5% H2/Ar flow.

XPS measurement was also performed on 0.01-CeO2/MgO(air) and 0.01-CeO2/MgO(dry) to clarify the difference of CeO2 nanoparticles caused by the preparation condition. As shown in Figure 6, a distinct difference was observed in the spectra. There were ten peaks derived from Ce4+ and Ce3+. The peaks annotated by v, v", v"', u, u", and u"' were attributed to Ce4+. Meanwhile, v0, v', u0, and u' were due to Ce3+. The binding energies of each peak are described in the caption of Figure 6. Intense peaks of Ce3+ denoted by v' and u' were observed for 0.01-CeO2/MgO(dry). Further, the relative intensity of the Ce3d3/2 peak denoted by u"' was lower for 0.01-CeO2/MgO(dry) than for 0.01-CeO2/MgO(air), suggesting the presence of Ce3+ in 0.01-CeO2/MgO(dry) [45,46]. Reportedly, the relative intensity of u"' is not linearly correlated to the content of Ce3+, but such a clear difference in the intensity of u"' suggests that more than 30% of Ce in 0.01-CeO2/MgO(dry) was Ce3+ [46,58]. This difference in the Ce3+ content is ascribed to the size of CeO2. Several studies reported that CeO2 nanoparticles smaller than 3 nm in diameter were relaxed in their crystal structure [13,19]; such small CeO2 nanoparticles were vulnerable to in situ reductions during the exposure to X-ray under vacuum. As shown in Figure 3, 0.01- CeO2/MgO(dry) contained a lot of CeO2 nanoparticles smaller than 3 nm, which would be reduced during the XPS measurement. Such high reducibility of 0.01-CeO2/MgO(dry) matched well with the results of H2-TPR (Figure 4). 0.01-CeO2/MgO(dry) was reduced at a lower temperature than 0.01-CeO2/MgO(air). Here, it is noteworthy that the effect of precursors on the valence of Ce was minor because both samples were prepared using the same precursor [13].

**Figure 6.** XPS spectra of 0.01-CeO2/MgO(air) and 0.01-CeO2/MgO(dry). The peak positions of v0, v, v', v", v"', u0, u, u', u", and u"' are 880.60, 882.60, 885.45, 888.85, 898.40, 898.90, 901.05, 904.05, 907.45, and 916.70 eV, respectively. The binding energies of all the peaks are identical to the values of the reference [46].

#### *3.3. In Situ XRD Study under the Reducing Atmosphere*

In situ XRD was performed to visualize the structural variation of CeO2 nanoparticles during the reduction. Instead of 0.01-CeO2/MgO, 0.05-CeO2/MgO(air) and 0.05- CeO2/MgO(dry) were used to increase the peak intensity to perform further analyses. The measurements were performed under the same condition as that of H2-TPR, but an N2 balance was used as a balance instead of Ar. CeO2 was also subjected to the same measurement as a reference. The results are shown in Figure 7. Only the peaks of CeO2 and MgO were observed in the patterns since the phase change from CeO2 to Ce2O3 is

slow. The absence of Ce2O3 was consistent with the previous studies [1,59,60]. As shown in the patterns of CeO2/MgO (Figure 7b,c), the peaks of MgO linearly shifted to lower 2*θ* as the temperature increased, suggesting a continuous thermal expansion of MgO crystals. The peaks of CeO2 also shifted to the same direction; however, the peaks showed the inflection points at ~650 ◦C (Figure 7b,c). This implies that structural relaxation was caused by the removal of lattice oxygen of the CeO2 nanoparticles. Additionally, CeO2 in both CeO2/MgO(air) and CeO2/MgO(dry) underwent structural change at lower temperatures than pure CeO2 (Figure 7a–c). The inflection point appeared at ~800 ◦C for pure CeO2 (Figure 7a). This result suggests that CeO2 in CeO2/MgO was reduced faster than pure CeO2, which is in good agreemen<sup>t</sup> with H2-TPR (Figure 4).

**Figure 7.** In situ XRD patterns of (**a**) CeO2, (**b**) 0.05-CeO2/MgO(air), and (**c**) 0.05-CeO2/MgO(dry). A Co X-ray source and an Fe filter were used for the measurements. White arrows indicate the inflection points observed on the (220) peaks of CeO2. The unassigned weak peaks were attributed to the diffraction of Co K-*β*.

Another difference between the CeO2/MgO and the pure CeO2 was observed in the shape of the diffraction patterns of CeO2 (Figure 7a–c). Notably, the splitting of diffraction patterns was observed for pure CeO2 around 2*θ* = ~55◦ and ~65◦ at ~750 ◦C (Figure 6a and Figure S4). This result suggests that the reduction proceeded heterogeneously in CeO2 grains. As shown in Figure 4, H2-TPR of pure CeO2 detected a long tailing to a high temperature at >800 ◦C. Such tailing means that the reduction of pure CeO2 was limited by the diffusion of oxygen in CeO2 [61–63]. Such diffusion-controlled reduction would render the outer side of CeO2 particles more reduced than the inner part of them. In contrast, such splitting was not observed for CeO2/MgO catalysts. The reduction of CeO2 nanoparticles in CeO2/MgO proceeded more homogeneously than pure CeO2 due to the small size of CeO2 nanoparticles.

The variation of unit cell parameter a0 of CeO2 was calculated assuming that the structural change of CeO2 nanoparticles was isotropic. The calculated values of a0 were normalized by the values at 200 ◦C to see the variation depending on the temperature. As shown in Figure 8, the unit cell parameter a0 increased at a lower temperature for 0.05- CeO2/MgO(dry) than for 0.05-CeO2/MgO(air). The first derivative of the unit cell parameter held a peak at 640 ◦C and 680 ◦C for 0.05-CeO2/MgO(dry) and 0.05-CeO2/MgO(air), respectively (Figure S5). This difference was ascribed to the smaller size of CeO2 in 0.05-CeO2/MgO(dry). As confirmed by H2-TPR (Figure 4), 0.05-CeO2/MgO(dry) was reduced at a lower temperature than 0.05-CeO2/MgO(air) due to the higher dispersion

of CeO2 (Figure 3). The faster removal of the oxygen from the lattice led to the structural change at a lower temperature. In addition, structural relaxation proceeds more easily as the size of CeO2 decreases [13,17,19]. The results of in situ XRD also demonstrated that 0.05-CeO2/MgO(dry) was reduced at lower temperatures than 0.05-CeO2/MgO(air).

**Figure 8.** Unit cell parameter a0 of CeO2 nanoparticles of 0.05-CeO2/MgO(air) and 0.05- CeO2/MgO(dry) calculated from in situ XRD measurement. The vertical axis is normalized by the value of a0 at 200 ◦C.

## *3.4. Dry Reforming Reaction*

As shown in previous sections, CeO2 nanoparticles in 0.05-CeO2/MgO(dry) were smaller than those in 0.05-CeO2/MgO(air). Such small CeO2 nanoparticles led to the high mobility of the lattice oxygen. In this section, the difference of the catalysts was demonstrated via a dry reforming reaction, in which the catalysts were exposed to dry reductive gas at high temperatures. Figure 9 shows the average reaction rate of CH4 throughout the reaction. The time course profile of the CH4 conversion rate is also shown in Figure S6. All the catalysts remained white or yellow after the reaction, suggesting the absence of carbon deposition. Both 0.05-CeO2/MgO(air) and 0.05-CeO2/MgO(dry) outperformed pure CeO2 despite their low content of CeO2, i.e., 18.3 wt%. Further, a higher conversion was attained over 0.05-CeO2/MgO(dry) than over 0.05-CeO2/MgO(air). These differences were ascribed to the dispersion of CeO2 nanoparticles. The average diameter of pure CeO2 was 61.6 nm while those of 0.05-CeO2/MgO(air) and 0.05-CeO2/MgO(dry) were 6.9 nm and 4.5 nm, respectively (Table 1). The finer CeO2 nanoparticles led to the higher surface area of CeO2, resulting in the higher catalytic activity. In addition, 0.05- CeO2/MgO(air) exhibited higher activity than pure CeO2 although the area of CeO2 in 0.05-CeO2/MgO(air), 12.5 m2/g, was smaller than that of pure CeO2, 13.5 m2/g (Table 1). Our previous study also suggests that MgO promotes CO2 supply to the CeO2 surface due to the strong basicity of the MgO [36]. Such interaction would also contribute to the higher catalytic activity of CeO2/MgO.

The selectivity of the reaction is also illustrated in Figure S7 as the time course profile of the outlet gas composition. All the reaction proceeds under the reaction condition of a low conversion rate of CH4, ~4%. Due to the low conversion rate, the selectivity of H2 against CO was as low as 0.1–0.2 (Figure S7). Notably, the average value of H2/CO throughout the 6-h reaction was slightly smaller for pure CeO2, 0.11, than 0.05-CeO2/MgO(air), 0.18, and 0.05-CeO2/MgO(dry), 0.17. The origin of this difference is uncertain, but we attributed it to the CH4 conversion rate and reducibility of CeO2. As shown in Figure 9, the CH4 conversion rate over the pure CeO2 is lower than those over 0.05-CeO2/MgO. This means the total supply of H2 is less over CeO2 than over 0.05-CeO2/MgO, causing the difference

in the H2/CO ratio. Another possible origin of the difference is the reducibility of CeO2. H2- TPR demonstrated that reduction of pure CeO2 did not complete even at 800 ◦C (Figure 4). This result suggests that the pure CeO2 continued to be reduced by the product H2 during the reaction. Such a reaction would lead to a lower H2/CO ratio due to the consumption of H2.

**Figure 9.** Average reaction rates of CH4 of 6-h dry reforming reaction over CeO2 and CeO2/MgO.

The temporal variation of the CH4 conversion rate also depended on the catalysts. The pure CeO2 was deactivated rapidly within the initial two hours but then gradually (Figure S6). On the contrary, such an initial rapid drop of the CH4 conversion rate was not observed for 0.05-CeO2/MgO(air) and 0.05-CeO2/MgO(dry). This difference in the initial behavior was attributed to the stoichiometry of the CeO2 surface influenced by the diffusion of oxygen in the CeO2 particles. As shown in Figure 4, the reduction of pure CeO2 at 800 ◦C was diffusion controlled. The reduction did not complete even at 900 ◦C. This result suggests that it takes a long time to reach the equilibrium in the oxygen content of the pure CeO2 during the reaction. Contrary, the reduction of 0.05-CeO2/MgO(air) and 0.05-CeO2/MgO(dry) was almost completed at 800 ◦C in H2-TPR (Figure 4), suggesting that it takes a short time to reach the equilibrium in the oxygen content of CeO2/MgO during the reaction. Therefore, there would be fewer oxygen vacancies on the surface of pure CeO2 than on CeO2 nanoparticles in CeO2/MgO. Such fewer oxygen vacancies of the pure CeO2 would improve the initial activity because previous studies reported that the C-H activation energy on CeO2 decreased as the number of oxygen vacancies decreased [64,65].

Notably, 0.05-CeO2/MgO(air) exhibited steady catalytic activity throughout the 6-h reaction, but 0.05-CeO2/MgO(dry) was gradually deactivated (Figure S6). The gradual deactivation of 0.05-CeO2/MgO(dry) was attributed to the gradual sintering of CeO2 nanoparticles caused by the moisture content in the gas. Dry reforming produces H2O along with H2 and CO. Although the concentration of water in the product gas was lower than that of ambient air (Table S3), 0.05-CeO2/MgO(dry) was exposed to the water vapor at 800 ◦C during the reaction, causing the structural change of 0.05-CeO2/MgO(dry). On the contrary, 0.05-CeO2/MgO(air) was already heated at 800 ◦C in the ambient air containing ~3 vol% of H2O before the reaction (Table S3). Therefore, the byproduct H2O in the product gas affected 0.05-CeO2/MgO(dry) more than 0.05-CeO2/MgO(air).

As discussed above, 0.05-CeO2/MgO(dry) gradually deactivated during the dry reforming reaction while 0.05-CeO2/MgO(air) exhibited stable activity. This difference in the behavior of 0.05-CeO2/MgO(air) and 0.05-CeO2/MgO(dry) suggests that the preparation of catalysts under the dry atmosphere is effective only for the reactions in the absence of H2O. Solar-thermochemical reaction is an example of such reactions [7]. The solarthermochemical reaction proceeds as a two-step reaction; the catalyst is subjected to high temperatures to be reduced thermodynamically and then re-oxidized by CO2 at low temperatures to produce CO. The CeO2/MgO(dry) would be suitable for the reaction since it demonstrated high stability and reducibility under dry conditions (Figures 4 and 5).
