Catalytic-CO2-Desorption Studies of BZA-AEP Mixed Absorbent by the Lewis Acid Catalyst CeO2-γ-Al2O3

Traditional organic amines exhibit inferior desorption performance and high regeneration energy consumption. The implementation of solid acid catalysts presents an efficacious approach to mitigate regeneration energy consumption. Thus, investigating high-performance solid acid catalysts holds paramount importance for the advancement and implementation of carbon capture technology. This study synthesized two Lewis acid catalysts via an ultrasonic-assisted precipitation method. A comparative analysis of the catalytic desorption properties was conducted, encompassing these two Lewis acid catalysts and three precursor catalysts. The results demonstrated that the CeO2-γ-Al2O3 catalyst demonstrated superior catalytic desorption performance. Within the desorption temperature range of 90 to 110 °C, the average desorption rate of BZA-AEP catalyzed by the CeO2-γ-Al2O3 catalyst was 87 to 354% greater compared to the desorption rate in the absence of the catalyst, and the desorption temperature can be reduced by approximately 10 °C. A comprehensive analysis of the catalytic desorption mechanism of the CeO2-γ-Al2O3 catalyst was conducted, and indicated that the synergistic effect of CeO2-γ-Al2O3 conferred a potent catalytic influence throughout the entire desorption process, spanning from the rich solution to the lean solution.


Introduction
Elevated emissions of greenhouse gases are contributing to severe climate change. The international community has reached a consensus on the necessity of controlling greenhouse gas emissions. According to the IPCC's 1.5 • C global warming report [1], human activities are estimated to have induced approximately 1.0 • C of global warming above pre-industrial levels, with a possible range of 0.8 • C to 1.2 • C. Should the current rate of warming persist, global warming may reach 1.5 • C between 2030 and 2052. Carbon capture and storage (CCS) represents a crucial means of reducing carbon dioxide emissions in the future, with the most promising application being CO 2 capture by organic amines [2]. However, traditional organic amine absorbents exhibit the drawback of high regeneration energy consumption. Introducing catalysts can facilitate carbamate decomposition and CO 2 desorption at reduced temperatures [3].
Idem et al. [4] first reported the employment of solid acid catalysts in the CO 2 desorption process involving amine-rich solutions. The researchers demonstrated that H-ZSM-5 and γ-Al 2 O 3 , two prevalent industrial solid acid catalysts, possess the ability to enhance the desorption performance of mono-ethanolamine (MEA) solution. Shi et al. [5] employed single and mixed amines (MEA, MEA-MDEA (N-methyl-diethanolamine), MEA-DEAB (4-(diethylamino)-2-butanol)) to compare the catalytic desorption effects of γ-Al 2 O 3 and H-ZSM-5 at 90-95 • C. They discovered that the catalytic efficacy of H-ZSM-5 surpassed that of γ-Al 2 O 3 . The addition of MDEA or DEAB (as a tertiary amine) to MEA provides enabling it to maintain catalytic activity even under harsh operating conditions [11][12][13][14]. Furthermore, in various other domains, metal-supported catalysts have demonstrated excellent catalytic performance [15][16][17]. Consequently, the utilization of a metal oxidesupported γ-Al 2 O 3 catalyst holds great potential for facilitating carbon dioxide desorption and reducing regeneration energy consumption.
Mao et al. [18] investigated the absorption and desorption characteristics of the BZA-AEP mixed amine absorbent, which demonstrated remarkable performance. In comparison to mono-ethanolamine (MEA), the BZA-AEP absorbent exhibited a 48% increase in average CO 2 absorption rate, a 120% enhancement in CO 2 desorption capacity, and a 161% rise in average CO 2 desorption rate. However, the regeneration efficiency of BZA-AEP was approximately 55%. To further augment its desorption efficacy, two Lewis acid catalysts were synthesized in this study, and the catalytic desorption properties of these catalysts for BZA-AEP were examined.

Catalytic Desorption Performance of CeO 2 -γ-Al 2 O 3
The desorption capacity of BZA-AEP was investigated for 2 h using various catalysts, including a commercial VWT catalyst, three precursor catalysts, and two M-γ-Al 2 O 3 (M = ZnO or CeO 2 ) Lewis acid catalysts. It can be seen from Figure 1 that CeO 2 -γ-Al 2 O 3 exhibited the best catalytic effect, and the capacity of carbon dioxide desorption increased by 30% compared with the case without catalyst. The characterization analysis showed that this benefited from the large specific surface area and acidity of CeO 2 -γ-Al 2 O 3 , which increased active catalytic surface. The catalyzed desorption performance follows the order: CeO 2 -γ-Al 2 O 3 > CeO 2 > VWT ≈ ZnO > γ-Al 2 O 3 > ZnO-γ-Al 2 O 3 > without catalyst. During the initial 20 min, the desorption capacity for each catalyst in the high carbon dioxideloaded absorbent exhibit minimal differences, indicating that each catalyst possesses a strong catalytic desorption capability. In the mechanism of Lewis acid catalyst catalysis [3], the acidic sites on the catalyst's surface interact with oxygen atoms in the absorbent, facilitating the decomposition of carbamate and the subsequent release of carbon dioxide.
Molecules 2023, 28, x FOR PEER REVIEW 3 of 17 additional active sites. Additionally, γ-Al2O3 exhibits commendable chemical stability, enabling it to maintain catalytic activity even under harsh operating conditions [11][12][13][14]. Furthermore, in various other domains, metal-supported catalysts have demonstrated excellent catalytic performance [15][16][17]. Consequently, the utilization of a metal oxide-supported γ-Al2O3 catalyst holds great potential for facilitating carbon dioxide desorption and reducing regeneration energy consumption. Mao et al. [18] investigated the absorption and desorption characteristics of the BZA-AEP mixed amine absorbent, which demonstrated remarkable performance. In comparison to mono-ethanolamine (MEA), the BZA-AEP absorbent exhibited a 48% increase in average CO2 absorption rate, a 120% enhancement in CO2 desorption capacity, and a 161% rise in average CO2 desorption rate. However, the regeneration efficiency of BZA-AEP was approximately 55%. To further augment its desorption efficacy, two Lewis acid catalysts were synthesized in this study, and the catalytic desorption properties of these catalysts for BZA-AEP were examined.

Catalytic Desorption Performance of CeO2-γ-Al2O3
The desorption capacity of BZA-AEP was investigated for 2 h using various catalysts, including a commercial VWT catalyst, three precursor catalysts, and two M-γ-Al2O3 (M = ZnO or CeO2) Lewis acid catalysts. It can be seen from Figure 1 that CeO2-γ-Al2O3 exhibited the best catalytic effect, and the capacity of carbon dioxide desorption increased by 30% compared with the case without catalyst. The characterization analysis showed that this benefited from the large specific surface area and acidity of CeO2-γ-Al2O3, which increased active catalytic surface. The catalyzed desorption performance follows the order: CeO2-γ-Al2O3 > CeO2 > VWT ≈ ZnO > γ-Al2O3 > ZnO-γ-Al2O3 > without catalyst. During the initial 20 min, the desorption capacity for each catalyst in the high carbon dioxideloaded absorbent exhibit minimal differences, indicating that each catalyst possesses a strong catalytic desorption capability. In the mechanism of Lewis acid catalyst catalysis [3], the acidic sites on the catalyst's surface interact with oxygen atoms in the absorbent, facilitating the decomposition of carbamate and the subsequent release of carbon dioxide.  Figure 2 compares the average CO2 desorption rate and regeneration efficiency of BZA-AEP catalyzed by the no catalyst, commercial VWT catalyst, three precursor catalysts and two M-γ-Al2O3 Lewis acid catalysts. The results indicate that the CeO2-γ-Al2O3 Lewis acid catalyst outperforms the others, exhibiting superior average CO2 desorption rates and regeneration efficiencies. As seen in Figure 2, the average CO2 desorption rate order is as follows: CeO2-γ-Al2O3 > CeO2 > VWT > γ-Al2O3 > ZnO ≈ ZnO-γ-Al2O3 > without catalyst. The CeO2-γ-Al2O3 Lewis acid catalyst increases the average CO2 desorption rate of BZA-AEP by 87% compared to the no catalyst and by 17% compared to the commercial VWT catalyst, highlighting its exceptional performance in enhancing desorption rates. In  Figure 2 compares the average CO 2 desorption rate and regeneration efficiency of BZA-AEP catalyzed by the no catalyst, commercial VWT catalyst, three precursor catalysts and two M-γ-Al 2 O 3 Lewis acid catalysts. The results indicate that the CeO 2 -γ-Al 2 O 3 Lewis acid catalyst outperforms the others, exhibiting superior average CO 2 desorption rates and regeneration efficiencies. As seen in Figure 2, the average CO 2 desorption rate order is as follows: CeO 2 -γ-Al 2 O 3 > CeO 2 > VWT > γ-Al 2 O 3 > ZnO ≈ ZnO-γ-Al 2 O 3 > without catalyst. The CeO 2 -γ-Al 2 O 3 Lewis acid catalyst increases the average CO 2 desorption rate of BZA-AEP by 87% compared to the no catalyst and by 17% compared to the commercial VWT catalyst, highlighting its exceptional performance in enhancing desorption rates. In terms of regeneration efficiency, the catalyst order is as follows: CeO 2 -γ-Al 2 O 3 > CeO 2 > VWT ≈ ZnO > γ-Al 2 O 3 > ZnO-γ-Al 2 O 3 > without catalyst. Under the catalysis of the terms of regeneration efficiency, the catalyst order is as follows: CeO2-γ-Al2O3 > CeO2 > VWT ≈ ZnO > γ-Al2O3 > ZnO-γ-Al2O3 > without catalyst. Under the catalysis of the CeO2γ-Al2O3 Lewis acid catalyst, the regeneration efficiency of BZA-AEP reaches 73%, which is 30% and 7% higher than that of the no catalyst and the commercial VWT catalyst, respectively. This further substantiates the superior performance of the CeO2-γ-Al2O3 Lewis acid catalyst in promoting BZA-AEP regeneration.

Effect of Temperature on Catalytic Desorption Performance of CeO2-γ-Al2O3
The impact of CeO2-γ-Al2O3 catalyst on the catalytic desorption of BZA-AEP was investigated for a duration of 2 h at desorption temperatures ranging from 90 °C to 110 °C. It can be clearly seen from Figure 3 that, with the increase of the desorption temperature, the difference in the desorption capacity within 2 h between catalyst catalysis and no catalyst catalysis presents a first increasing and then decreasing trend. Notably, the CeO2-γ-Al2O3 catalyst demonstrates the most significant increase in BZA-AEP desorption capacity at a 100 °C desorption temperature. Under the catalysis of CeO2-γ-Al2O3, the desorption capacity of BZA-AEP increased by 2-46% compared to no catalyst, signifying the substantial advantage of the CeO2-γ-Al2O3 catalyst in improving BZA-AEP desorption capacity. Furthermore, in comparison with the MEA desorption capacity of no catalyst, the desorption amount of BZA-AEP catalyzed by CeO2-γ-Al2O3 increased by 40-222%. As depicted in Figure 4c, the maximum desorption rate of BZA-AEP under CeO2-γ-Al2O3 catalysis increased with rising desorption temperature, and the time required to reach the maximum desorption rate gradually shortened. Figure 4a-c show that CeO2-γ-

Effect of Temperature on Catalytic Desorption Performance of CeO 2 -γ-Al 2 O 3
The impact of CeO 2 -γ-Al 2 O 3 catalyst on the catalytic desorption of BZA-AEP was investigated for a duration of 2 h at desorption temperatures ranging from 90 • C to 110 • C. It can be clearly seen from Figure 3 that, with the increase of the desorption temperature, the difference in the desorption capacity within 2 h between catalyst catalysis and no catalyst catalysis presents a first increasing and then decreasing trend. Notably, the CeO 2 -γ-Al 2 O 3 catalyst demonstrates the most significant increase in BZA-AEP desorption capacity at a 100 • C desorption temperature. Under the catalysis of CeO 2 -γ-Al 2 O 3 , the desorption capacity of BZA-AEP increased by 2-46% compared to no catalyst, signifying the substantial advantage of the CeO 2 -γ-Al 2 O 3 catalyst in improving BZA-AEP desorption capacity. Furthermore, in comparison with the MEA desorption capacity of no catalyst, the desorption amount of BZA-AEP catalyzed by CeO 2 -γ-Al 2 O 3 increased by 40-222%. terms of regeneration efficiency, the catalyst order is as follows: CeO2-γ-Al2O3 > CeO2 > VWT ≈ ZnO > γ-Al2O3 > ZnO-γ-Al2O3 > without catalyst. Under the catalysis of the CeO2γ-Al2O3 Lewis acid catalyst, the regeneration efficiency of BZA-AEP reaches 73%, which is 30% and 7% higher than that of the no catalyst and the commercial VWT catalyst, respectively. This further substantiates the superior performance of the CeO2-γ-Al2O3 Lewis acid catalyst in promoting BZA-AEP regeneration.

Effect of Temperature on Catalytic Desorption Performance of CeO2-γ-Al2O3
The impact of CeO2-γ-Al2O3 catalyst on the catalytic desorption of BZA-AEP was investigated for a duration of 2 h at desorption temperatures ranging from 90 °C to 110 °C. It can be clearly seen from Figure 3 that, with the increase of the desorption temperature, the difference in the desorption capacity within 2 h between catalyst catalysis and no catalyst catalysis presents a first increasing and then decreasing trend. Notably, the CeO2-γ-Al2O3 catalyst demonstrates the most significant increase in BZA-AEP desorption capacity at a 100 °C desorption temperature. Under the catalysis of CeO2-γ-Al2O3, the desorption capacity of BZA-AEP increased by 2-46% compared to no catalyst, signifying the substantial advantage of the CeO2-γ-Al2O3 catalyst in improving BZA-AEP desorption capacity. Furthermore, in comparison with the MEA desorption capacity of no catalyst, the desorption amount of BZA-AEP catalyzed by CeO2-γ-Al2O3 increased by 40-222%. As depicted in Figure 4c, the maximum desorption rate of BZA-AEP under CeO2-γ-Al2O3 catalysis increased with rising desorption temperature, and the time required to reach the maximum desorption rate gradually shortened. Figure 4a-c show that CeO2-γ- by 292%, 159%, 89%, 61% and 64% and, compared with MEA without catalysis, increased by 211%, 380%, 286%, 138% and 102%, respectively. Figure 4d reveals that the average CO 2 desorption rate linearly increases with desorption temperature. The CeO 2 -γ-Al 2 O 3 catalyst can reduce the desorption temperature of BZA-AEP from 110 • C to 100 • C. Under CeO 2γ-Al 2 O 3 catalysis, the average desorption rate of BZA-AEP surpasses that of BZA-AEP without catalysis by 87-354%, and exceeds that of MEA without catalysis by 141-400%. These data conclusively confirm that the CeO 2 -γ-Al 2 O 3 Lewis acid catalyst significantly increases the CO 2 desorption rate of BZA-AEP at various desorption temperatures.

Cyclic Catalytic Desorption Performance of CeO2-γ-Al2O3
According to the results in Section 3.2, the cycle desorption temperature was set at 100 °C. As observed in Figure 5a, the loading of the BZA-AEP rich solution reached 0.63438 mol CO2/mol amine during the initial absorption and desorption process. In the subsequent two cycles, the rich solution load experienced a slight decrease, stabilizing at approximately 0.6 mol CO2/mol amine. This decrease can be attributed to some amine absorbent and carbamates were adsorbed on the surface of the fresh catalyst during the first cycle, so that this part of amine absorbent failed to enter the absorber to participate in the absorption process. Concurrently, the lean liquid loading increased marginally in the last two cycles compared to the first cycle and stabilized at around 0.2 mol CO2/mol amine, further corroborating the hypothesis that some amine absorbent and carbamates were adsorbed on the catalyst surface. It can be seen from Figure 5b that the desorption capacity of the absorbent tends to be stable after the first cycle. Under the catalysis of

Cyclic Catalytic Desorption Performance of CeO 2 -γ-Al 2 O 3
According to the results in Section 3.2, the cycle desorption temperature was set at 100 • C. As observed in Figure 5a, the loading of the BZA-AEP rich solution reached 0.63438 mol CO 2 /mol amine during the initial absorption and desorption process. In the subsequent two cycles, the rich solution load experienced a slight decrease, stabilizing at approximately 0.6 mol CO 2 /mol amine. This decrease can be attributed to some amine absorbent and carbamates were adsorbed on the surface of the fresh catalyst during the first cycle, so that this part of amine absorbent failed to enter the absorber to participate in the absorption process. Concurrently, the lean liquid loading increased marginally in the last two cycles compared to the first cycle and stabilized at around 0.2 mol CO 2 /mol amine, further corroborating the hypothesis that some amine absorbent and carbamates were adsorbed on the catalyst surface. It can be seen from Figure 5b that the desorption capacity of the absorbent tends to be stable after the first cycle. Under the catalysis of CeO 2 -γ-Al 2 O 3 , the cycle capacity of BZA-AEP reaches 0.40147 mol CO 2 /mol amine, which is 31% higher than that of uncatalyzed BZA-AEP and 108% higher than that of uncatalyzed MEA. CeO2-γ-Al2O3, the cycle capacity of BZA-AEP reaches 0.40147 mol CO2/mol amine, which is 31% higher than that of uncatalyzed BZA-AEP and 108% higher than that of uncatalyzed MEA.
(a) (b) It can be seen from Figure 6 that the average absorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 is higher than that of uncatalyzed BZA-AEP during the second and third cycles. The reason for this phenomenon is that the BZA-AEP catalyzed by CeO2-γ-Al2O3 releases more carbon dioxide during the desorption process. The concentration of amines not bound to carbon dioxide in BZA-AEP lean solution catalyzed by CeO2-γ-Al2O3 is higher than that in BZA-AEP lean solution without catalysis: the higher the amine concentration, the faster the reaction rate of amine and carbon dioxide. From Figure 7, we can observe that the average CO2 desorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 starts to stabilize after the second cycle, indicating the catalyst's good stability. Among them, the average CO2 desorption rate of the second cycle is lower than that of the first cycle due to a decrease in the rich solution load caused by the exclusion of a portion of the amine solution from the absorption cycle. Consequently, the desorption capacity and average CO2 desorption rate decrease accordingly. Compared with the uncatalyzed BZA-AEP, the average CO2 desorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 increased by 144%. Compared with uncatalyzed MEA, the average CO2 desorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 increased by 268%. It can be seen from Figure 6 that the average absorption rate of BZA-AEP catalyzed by CeO 2 -γ-Al 2 O 3 is higher than that of uncatalyzed BZA-AEP during the second and third cycles. The reason for this phenomenon is that the BZA-AEP catalyzed by CeO 2 -γ-Al 2 O 3 releases more carbon dioxide during the desorption process. The concentration of amines not bound to carbon dioxide in BZA-AEP lean solution catalyzed by CeO 2 -γ-Al 2 O 3 is higher than that in BZA-AEP lean solution without catalysis: the higher the amine concentration, the faster the reaction rate of amine and carbon dioxide. CeO2-γ-Al2O3, the cycle capacity of BZA-AEP reaches 0.40147 mol CO2/mol amine, which is 31% higher than that of uncatalyzed BZA-AEP and 108% higher than that of uncatalyzed MEA.
(a) (b) It can be seen from Figure 6 that the average absorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 is higher than that of uncatalyzed BZA-AEP during the second and third cycles. The reason for this phenomenon is that the BZA-AEP catalyzed by CeO2-γ-Al2O3 releases more carbon dioxide during the desorption process. The concentration of amines not bound to carbon dioxide in BZA-AEP lean solution catalyzed by CeO2-γ-Al2O3 is higher than that in BZA-AEP lean solution without catalysis: the higher the amine concentration, the faster the reaction rate of amine and carbon dioxide. From Figure 7, we can observe that the average CO2 desorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 starts to stabilize after the second cycle, indicating the catalyst's good stability. Among them, the average CO2 desorption rate of the second cycle is lower than that of the first cycle due to a decrease in the rich solution load caused by the exclusion of a portion of the amine solution from the absorption cycle. Consequently, the desorption capacity and average CO2 desorption rate decrease accordingly. Compared with the uncatalyzed BZA-AEP, the average CO2 desorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 increased by 144%. Compared with uncatalyzed MEA, the average CO2 desorption rate of BZA-AEP catalyzed by CeO2-γ-Al2O3 increased by 268%. From Figure 7, we can observe that the average CO 2 desorption rate of BZA-AEP catalyzed by CeO 2 -γ-Al 2 O 3 starts to stabilize after the second cycle, indicating the catalyst's good stability. Among them, the average CO 2 desorption rate of the second cycle is lower than that of the first cycle due to a decrease in the rich solution load caused by the exclusion of a portion of the amine solution from the absorption cycle. Consequently, the desorption capacity and average CO 2 desorption rate decrease accordingly. Compared with the uncatalyzed BZA-AEP, the average CO 2 desorption rate of BZA-AEP catalyzed by CeO 2 -γ-Al 2 O 3 increased by 144%. Compared with uncatalyzed MEA, the average CO 2 desorption rate of BZA-AEP catalyzed by CeO 2 -γ-Al 2 O 3 increased by 268%.

SEM Characterization
The morphology of the catalyst was examined utilizing the Sigma 300 scanning electron microscope (SEM) both before and after its usage. Figure 8a-d display the SEM images of γ-Al 2 O 3 , CeO 2 , fresh CeO 2 -γ-Al 2 O 3 , and CeO 2 -γ-Al 2 O 3 catalyst after three times cycle, respectively. The γ-Al 2 O 3 (Figure 8a) shows that the catalyst is aggregated from extremely fine particles. The SEM images of CeO 2 (Figure 8b) corroborate the spherical morphology of CeO 2 nanoparticles [19]. The fresh CeO 2 -γ-Al 2 O 3 catalyst (Figure 8c) displays a high dispersion of γ-Al 2 O 3 among CeO 2 nanoparticles. In the SEM image of the CeO 2 -γ-Al 2 O 3 catalyst (Figure 8d) after three times cycle, we can see that the surface of the catalyst has not changed significantly, but the aggregation between the particles has become tighter. This observation aligns with the experimental data presented in previous sections, further substantiating the exceptional stability of the CeO 2 -γ-Al 2 O 3 catalyst. Figure 9 reveals that the catalyst particle size distribution follows a Gaussian distribution. The γ-Al 2 O 3 catalyst has a mean particle size of 6.35 nm, with a standard deviation of 1.06 nm. The CeO 2 catalyst exhibited an average particle size of 67.17 nm and a standard deviation of 24.59 nm. The CeO 2 -γ-Al 2 O 3 catalyst displayed an average particle size of 26.76 nm and a standard deviation of 8.40 nm.

SEM Characterization
The morphology of the catalyst was examined utilizing the Sigma 300 scanning electron microscope (SEM) both before and after its usage. Figure 8a-d display the SEM images of γ-Al2O3, CeO2, fresh CeO2-γ-Al2O3, and CeO2-γ-Al2O3 catalyst after three times cycle, respectively. The γ-Al2O3 (Figure 8a) shows that the catalyst is aggregated from extremely fine particles. The SEM images of CeO2 (Figure 8b) corroborate the spherical morphology of CeO2 nanoparticles [19]. The fresh CeO2-γ-Al2O3 catalyst ( Figure 8c) displays a high dispersion of γ-Al2O3 among CeO2 nanoparticles. In the SEM image of the CeO2-γ-Al2O3 catalyst (Figure 8d) after three times cycle, we can see that the surface of the catalyst has not changed significantly, but the aggregation between the particles has become tighter. This observation aligns with the experimental data presented in previous sections, further substantiating the exceptional stability of the CeO2-γ-Al2O3 catalyst. Figure 9 reveals that the catalyst particle size distribution follows a Gaussian distribution. The γ-Al2O3 catalyst has a mean particle size of 6.35 nm, with a standard deviation of 1.06 nm. The CeO2 catalyst exhibited an a verage particle size of 67.17 nm and a standard deviation of 24.59 nm. The CeO2-γ-Al2O3 catalyst displayed an average particle size of 26.76 nm and a standard deviation of 8.40 nm.

XRD Characterization
To determine any changes in the structure of the catalyst, both before and after use, the Ultima IV X-ray powder diffractometer was employed. Prior to the test, the catalyst was dried and pressed. The diffraction patterns were obtained using continuous scanning with a range of 10-80 • (2θ) and a scanning rate of 5 • /min. Figure 10 displays the crystal structure diffraction patterns of the two precursor catalysts (γ-Al 2 O 3 and CeO 2 ), the fresh CeO 2 -γ-Al 2 O 3 catalyst, and the CeO 2 -γ-Al 2 O 3 catalyst after three times cycle. The diffraction pattern of γ-Al 2 O 3 exhibits a weak characteristic peak intensity, indicating a small grain size. Moreover, only the diffraction peak of CeO 2 can be observed on the CeO 2 -γ-Al 2 O 3 catalyst, with no diffraction peak of crystalline γ-Al 2 O 3 evident. This suggests that γ-Al 2 O 3 may be highly dispersed on CeO 2 or present as clusters, surpassing the XRD detection limit [20][21][22][23][24]. The crystal structures of γ-Al 2 O 3 (PDF-ICDD 01-079-1558) [25], CeO 2 (PDF-ICDD 00-043-1002) [26], and CeO 2 -γ-Al 2 O 3 belong to the cubic system. The characteristic peaks of the CeO 2 -γ-Al 2 O 3 catalyst before and after cycling remain essentially unchanged, indicating that the cycling process does not impact the catalyst's structure. The diffraction peak intensity varies among these catalysts. A narrower peak corresponds to a larger grain size and better crystallinity, while a wider peak signifies a smaller grain size and poorer crystallinity [27][28][29]. The CeO 2 -γ-Al 2 O 3 catalyst possesses a highly ordered crystal structure, distinct diffraction peaks, and narrow peak width, which implies good crystallinity and high stability-crucial for long-term catalytic applications. Table 1 presents the grain sizes of γ-Al 2 O 3 , CeO 2 , CeO 2 -γ-Al 2 O 3 , and CeO 2 -γ-Al 2 O 3 after three times cycle, as determined using the Debye-Scherrer equation [30,31] (Dβ = Kλ/βcosθ), based on the strongest diffraction peak. Additionally, the interplanar spacing of the strongest peaks of these catalysts is computed using the Bragg equation [32] (2dsinθ = nλ). The CeO 2 -γ-Al 2 O 3 catalyst exhibits a smaller grain size compared to the CeO 2 catalyst, resulting in a larger active surface during the catalytic reaction. The addition of CeO 2 to γ-Al 2 O 3 also enhances its thermal stability and mitigates the sintering of CeO 2 nanoparticles [33], thereby maintaining the catalyst's long-term activity. Given that the ionic radius of Al 3+ (0.54 Å) is smaller than that of Ce 4+ (0.92 Å), the lattice constant of the CeO 2 -γ-Al 2 O 3 catalyst is slightly reduced compared to the CeO 2 catalyst. This observation indicates that some Al 3+ ions may be doped into the surface lattice of the CeO 2 catalyst. Such doping aids in enhancing the stability and catalytic activity of the catalyst, allowing the CeO 2 -γ-Al 2 O 3 catalyst to exhibit superior performance during the desorption process.
peak intensity, indicating a small grain size. Moreover, only the diffraction peak of CeO2 can be observed on the CeO2-γ-Al2O3 catalyst, with no diffraction peak of crystalline γ-Al2O3 evident. This suggests that γ-Al2O3 may be highly dispersed on CeO2 or present as clusters, surpassing the XRD detection limit [20][21][22][23][24]. The crystal structures of γ-Al2O3 (PDF-ICDD 01-079-1558) [25], CeO2 (PDF-ICDD 00-043-1002) [26], and CeO2-γ-Al2O3 belong to the cubic system. The characteristic peaks of the CeO2-γ-Al2O3 catalyst before and after cycling remain essentially unchanged, indicating that the cycling process does not impact the catalyst's structure. The diffraction peak intensity varies among these catalysts. A narrower peak corresponds to a larger grain size and better crystallinity, while a wider peak signifies a smaller grain size and poorer crystallinity [27][28][29]. The CeO2-γ-Al2O3 catalyst possesses a highly ordered crystal structure, distinct diffraction peaks, and narrow peak width, which implies good crystallinity and high stability-crucial for long-term catalytic applications.  Table 1 presents the grain sizes of γ-Al2O3, CeO2, CeO2-γ-Al2O3, and CeO2-γ-Al2O3 after three times cycle, as determined using the Debye-Scherrer equation [30,31] (Dβ = Kλ/βcosθ), based on the strongest diffraction peak. Additionally, the interplanar spacing of the strongest peaks of these catalysts is computed using the Bragg equation [32] (2dsinθ = nλ). The CeO2-γ-Al2O3 catalyst exhibits a smaller grain size compared to the CeO2 catalyst, resulting in a larger active surface during the catalytic reaction. The addition of CeO2 to γ-Al2O3 also enhances its thermal stability and mitigates the sintering of CeO2 nanoparticles [33], thereby maintaining the catalyst's long-term activity. Given that the ionic radius of Al 3+ (0.54 Å) is smaller than that of Ce 4+ (0.92 Å), the lattice constant of the CeO2-γ-Al2O3 catalyst is slightly reduced compared to the CeO2 catalyst. This observation indicates that some Al 3+ ions may be doped into the surface lattice of the CeO2 catalyst. Such doping aids in enhancing the stability and catalytic activity of the catalyst, allowing the CeO2-γ-Al2O3 catalyst to exhibit superior performance during the desorption process.

BET and NH 3 -TPD Characterization
The specific surface area of the catalyst was measured using the fully automatic surface area and porosity analysis of the ASAP 2460. Prior to testing the catalyst, a vacuum degassing pre-treatment was carried out at a temperature of 200 • C for 4 h. The BET (Brunauer Emmett Teller) method was used to calculate the specific surface area of the catalyst. To examine the concentration distribution of catalyst acid, including TCD detector, the AutoChem II 2920 chemisorption instrument was utilized. Before the sample test, the temperature was raised to 400 • C at a rate of 10 • C/min in an argon atmosphere and maintained for 1 h to eliminate any physically adsorbed water and impurities from the sample surface. The temperature was then reduced to 50 • C. Next, a 10% NH 3 -He gas flow was introduced onto the catalyst surface for adsorption saturation, followed by a highpurity He gas blow for 1 h to remove any weak physical adsorption of NH 3 on the surface. Finally, the NH 3 -TPD curve was obtained by heating up to 450 • C at a rate of 10 • C/min. The Gaussian deconvolution method was applied to perform a semi-quantitative analysis of the TPD curve to determine the acidity of the catalyst. Table 2 reveals the variations in specific surface area and acid strength among the three catalysts. γ-Al 2 O 3 exhibits a significantly larger specific surface area, suggesting the availability of a greater number of active surfaces for catalytic reactions. In contrast, the specific surface area of CeO 2 is relatively small, measuring only 2.3596 m 2 /g. Loading CeO 2 onto γ-Al 2 O 3 leads to an increase in the specific surface area of CeO 2 , but simultaneously reduces the specific surface area of γ-Al 2 O 3 . Figure 11 displays the catalyst's acid strength distribution characterized by NH 3 -TPD. The NH 3 desorption peak at 100-200 • C typically corresponds to the weak acid site, the NH 3 desorption peak in the range of 200-400 • C corresponds to the medium acid site, and the NH 3 desorption peak above 400 • C corresponds to the strong acid site [34]. As evident in Figure 11, the low-temperature desorption peak near 120 • C arises from the catalyst's weak acid site, while the medium-temperature desorption peak near 350 • C is attributable to the catalyst's medium-strong acid site. Table 2 shows that the weak acid sites of the CeO 2 -γ-Al 2 O 3 catalyst are significantly stronger than those of the two parent catalysts. The weak acid sites of the CeO 2 -γ-Al 2 O 3 catalyst increase by 287% and 20% compared to γ-Al 2 O 3 and CeO 2 parent catalysts, respectively. This suggests that CeO 2 loading enhances the weak acid sites of the γ-Al 2 O 3 catalyst.  Figure 11. Acid strength distribution of CeO2-γ-Al2O3 catalyst.

Catalytic Mechanism of CeO2-γ-Al2O3
According to the zwitterionic mechanism and the alkali-catalyzed bicarbonate reaction mechanism, the regeneration process of primary and secondary amines comprises two distinct steps: cleavage of the N-C bond of carbamates and deprotonation of protonated amines [35]. On the other hand, the regeneration process of tertiary amines involves bicarbonate hydrogenation decomposition and protonated amine deprotonation [36]. However, previous studies have indicated that CO2 desorption can be accelerated by providing a substantial number of Brønsted acid sites, Lewis acid sites, and HCO3 − -like alkaline groups [37,38].
Based on the above perspective, Figure 12 depicts the mechanism diagram of CO2 desorption catalyzed by CeO2-γ-Al2O3. The figure illustrates three different catalytic desorption processes for various absorbents (① for primary amine, ② for secondary amine, and ③ for tertiary amine). During the desorption process, the absorbent transitions from a rich solution to a lean solution, resulting in an increase in the pH value of the solution. Additionally, the catalytic desorption pathways vary in alkaline environments.

Catalytic Mechanism of CeO 2 -γ-Al 2 O 3
According to the zwitterionic mechanism and the alkali-catalyzed bicarbonate reaction mechanism, the regeneration process of primary and secondary amines comprises two distinct steps: cleavage of the N-C bond of carbamates and deprotonation of protonated amines [35]. On the other hand, the regeneration process of tertiary amines involves bicarbonate hydrogenation decomposition and protonated amine deprotonation [36]. However, previous studies have indicated that CO 2 desorption can be accelerated by providing a substantial number of Brønsted acid sites, Lewis acid sites, and HCO 3 − -like alkaline groups [37,38].
Based on the above perspective, Figure 12 depicts the mechanism diagram of CO 2 desorption catalyzed by CeO 2 -γ-Al 2 O 3 . The figure illustrates three different catalytic desorption processes for various absorbents ( 1 for primary amine, 2 for secondary amine, and 3 for tertiary amine). During the desorption process, the absorbent transitions from a rich solution to a lean solution, resulting in an increase in the pH value of the solution. Additionally, the catalytic desorption pathways vary in alkaline environments.
The rich solution primarily undergoes the following catalytic desorption process. Due to weak alkalinity, the AlO 2 − anion cannot form in the rich solution region [31], while the oxygen atom of CeO 2 readily receives H + , as it is more electronegative than the nitrogen atom on the amino group [10]. Consequently, the rich solution region is mainly CeO 2 to promote the deprotonation of protonated amine. The positively charged protonated amine (N atom) first adsorbs on the more negatively charged CeO 2 basic site (O atom) according to the principle of opposites attract, and the proton is transferred from the nitrogen atom of the amino group to the oxygen atom of CeO 2 , completing the deprotonation. Subsequently, CeO 2 transfers surface protons to carbamates and bicarbonates. Carbamates acquire protons, and the Lewis acid sites on the γ-Al 2 O 3 surface attack the O and N atoms of carbamates [3], promoting the stretching of N-C bonds and weakening bond strength, thus reducing the activation energy of the carbamate fracture reaction. Finally, through isomerization [4], the carbamate's proton transfers from the O atom to the nearby N atom, breaking the N-C bond and decomposing into amines and CO 2 . Bicarbonate directly decomposes into H 2 O and CO 2 after obtaining the protons transferred from CeO 2 . The rich solution primarily undergoes the following catalytic desorption process. Due to weak alkalinity, the AlO2 − anion cannot form in the rich solution region [31], while the oxygen atom of CeO2 readily receives H + , as it is more electronegative than the nitrogen atom on the amino group [10]. Consequently, the rich solution region is mainly CeO2 The lean solution primarily undergoes the following catalytic desorption process. In the strong alkaline environment of the lean solution, the surface of Al 2 O 3 exhibits electronegativity and forms an AlO 2 − basic group, as demonstrated by previous studies [38]. These anions capture protons from protonated amines and subsequently transfer them to carbamates via water. This process predominantly occurs in the lean solution region, meaning that, as CO 2 progressively desorbs, the absorbent's alkalinity gradually increases. The electronegativity of the oxygen atom for CeO 2 is less than that of the AlO 2 − basic group. hence, the AlO 2 − basic group is primarily responsible for proton transfer from the protonated amine in the lean solution region. The AlO 2 − basic group first transfers protons from the protonated amine to carbamate and bicarbonate. Then, due to the hole donor nature of CeO 2 [20], the Lewis acid site of CeO 2 can bind to the electron pair donor, allowing carbamate to accept the proton and be attacked by the Lewis acid site on the CeO 2 surface at its O and N atoms. This process results in the stretching of the N-C bond and the weakening of the bond energy, thereby reducing the activation energy of the carbamate cleavage reaction. Ultimately, through isomerization [4], the N-C bond of carbamate is broken, and the compound decomposes into amines and CO 2 . Simultaneously, bicarbonate directly decomposes into H 2 O and CO 2 after obtaining the proton transferred from the AlO 2 − basic group.

Catalyst Preparation Materials and Methods
The chemical reagents used in the experiment are shown in Table 3. The catalyst was synthesized using an ultrasonic-assisted precipitation method. Following the reported synthesis method [8], the specific synthesis pathway is depicted in Figure 13. Initially, CeCl 3 ·6H 2 O was added to 500 mL of deionized water, resulting in an aqueous solution of CeCl 3 . Subsequently, a suitable amount of γ-Al 2 O 3 powder was added to the solution, forming a suspension with a CeO 2 to γ-Al 2 O 3 molar ratio of 1. Ultrasonic treatment was applied to the suspension at room temperature, using an ultrasonic disperser operating at 10% power and 20 kHz for 0.5 h to ensure comprehensive mixing of the constituents. Next, a NaOH solution was gradually added to the suspension at room temperature with continuous stirring until the pH reached approximately 8-9. The system was then allowed to stand at room temperature for 2 h, resulting in the formation of a precipitate. The precipitate was subsequently washed with deionized water and filtered multiple times. The filtered precipitate was dried at 110 • C for 11 h in a blast drying oven. To obtain the desired catalyst CeO 2 -γ-Al 2 O 3 , the dried solid was calcined in a muffle furnace at 800 • C for 4 h. Employing the same preparation method, the ZnO-γ-Al 2 O 3 solid acid catalyst was obtained.  Figure 13. Ultrasound-assisted precipitation catalyst synthesis pathway.

Desorption Experimental Materials and Steup
The chemical compounds utilized in this study, namely ethanolamine (MEA, 99%), benzylamine (BZA, 99%), aminoethylpiperazine (AEP, 99%), carbon dioxide (CO2, 99.9%), and nitrogen (N2, 99.9%), were acquired without further purification. The specific conditions of the desorption experiment are presented in Table 4, and the experiment was executed utilizing the absorption and desorption experimental apparatus, as described in reference [11]. The specific experimental setup is shown in Figure 14. Prior to commencing the experiment, the system underwent a leak test and N2 purge. For the absorption experiments, valves 3a, 3b, and 3c were opened, and the inlet CO2 concentration was set to 5% with an inlet gas flow of 1.25 L/min, an absorption temperature of 50 °C, and an absorption solution of 50 g. For the desorption experiments, valves 3a, 3b, and 3c were closed, and the desorption temperature was set to 100 °C. Each group of experiments was repeated three times, and the results were averaged.

Name
Parameter absorbent BZA-AEP total amine concentration 3 mol/kg concentration ratio of BZA to AEP 1.5 amount of absorbent 50 g Figure 13. Ultrasound-assisted precipitation catalyst synthesis pathway.

Desorption Experimental Materials and Steup
The chemical compounds utilized in this study, namely ethanolamine (MEA, 99%), benzylamine (BZA, 99%), aminoethylpiperazine (AEP, 99%), carbon dioxide (CO 2 , 99.9%), and nitrogen (N 2 , 99.9%), were acquired without further purification. The specific conditions of the desorption experiment are presented in Table 4, and the experiment was executed utilizing the absorption and desorption experimental apparatus, as described in reference [11]. The specific experimental setup is shown in Figure 14. Prior to commencing the experiment, the system underwent a leak test and N 2 purge. For the absorption experiments, valves 3a, 3b, and 3c were opened, and the inlet CO 2 concentration was set to 5% with an inlet gas flow of 1.25 L/min, an absorption temperature of 50 • C, and an absorption solution of 50 g. For the desorption experiments, valves 3a, 3b, and 3c were closed, and the desorption temperature was set to 100 • C. Each group of experiments was repeated three times, and the results were averaged.