Energy-Efficient and Sustainable CO₂ Capture in MEA Systems Enabled by FeOOH Catalysts

Abstract

Carbon dioxide (CO₂) capture is a cornerstone of global carbon neutrality, yet the high energy penalty associated with solvent regeneration—particularly for monoethanolamine (MEA) systems—remains a major barrier to its sustainable deployment. This study presents a sustainable and high-performance catalytic solution using micro-sized iron oxyhydroxide (β-FeOOH). Characterized by a high specific surface area ($287 m²/g) and a synergistic distribution of abundant Lewis and Brønsted acid sites, the β-FeOOH catalyst significantly enhances CO₂ desorption kinetics. Experimental results demonstrate that the incorporation of β-FeOOH into a 30 wt% MEA solution increases the CO₂ desorption rate by 10.9% while simultaneously lowering the regeneration temperature from the conventional 120 °C to 85 °C. Such a reduction in thermal requirements offers a pathway to utilize low-grade industrial waste heat, drastically improving the process’s energy efficiency. Furthermore, the catalyst exhibited remarkable cyclic stability over ten consecutive cycles, maintaining its structural integrity and catalytic activity. These findings highlight β-FeOOH as an eco-friendly, cost-effective, and robust catalyst that aligns with the principles of green chemical engineering, offering a scalable strategy to enhance the sustainability of carbon capture operations.

1. Introduction

The rising concentrations of carbon dioxide (CO₂) in the atmosphere, mainly due to the burning of fossil fuels, are a major challenge for climate stability [1,2]. The development of high-efficiency carbon capture technologies is therefore not only a technical imperative but also a key enabler for sustainable industrial development and the achievement of the Paris Agreement targets. Among various mitigation strategies, post-combustion carbon capture using amine-based solvents—particularly aqueous monoethanolamine (MEA)—remains one of the most mature and widely implemented approaches due to its high reactivity toward CO₂ and established industrial applicability [3,4]. However, conventional MEA-based systems suffer from several inherent limitations, including high energy consumption during solvent regeneration, solvent degradation, and limited long-term cyclic stability [5,6]. These challenges significantly increase operational costs and carbon footprint, thereby constraining the large-scale deployment and overall sustainability of amine-based carbon capture processes.

To address these limitations, considerable effort has been devoted to the development of solid-phase additives and heterogeneous catalysts that can accelerate CO₂ absorption kinetics and reduce the regeneration energy demand of MEA-based solvents. A wide range of materials have been explored, including metal oxides, zeolites, porous carbons, metal–organic frameworks (MOFs), and ionic liquid-modified supports [7]. These materials generally improve system performance by facilitating gas–liquid mass transfer, promoting carbamate decomposition, and providing acid–base active sites for reversible CO₂ interactions.

Among metal oxides, TiO(OH)₂ nanosheets have demonstrated moderate activity, improving desorption rates by approximately 6.2% due to their hydrophilic surfaces and proton-conductive properties [8]. Similarly, boehmite (AlOOH) nanorods function as Lewis acidic promoters and enhance solvent fluidity, yielding a 5.5% increase in CO₂ release [9]. Zeolitic materials such as Fe-BEA, with high specific surface area (~520 m²/g) and bifunctional acid sites, have shown greater efficacy, achieving 8.4% enhancement in desorption under cyclic conditions [10]. These improvements are largely attributed to their ability to lower the energy barrier for carbamate decomposition through surface-catalyzed pathways. These studies suggest that hydroxylated metal oxides with accessible acidic surface sites can influence amine-based CO₂ absorption–desorption behavior, providing useful guidance for the design of heterogeneous promoters in MEA systems.

In parallel, MOFs with tunable functional groups and large surface areas—such as MIL-101 and ZIF-8—have been investigated for reversible CO₂ adsorption and desorption [11,12]. However, their application in hot, aqueous amine systems remains limited due to poor hydrothermal stability and degradation under acidic environments. Similarly, carbonaceous materials like nitrogen-doped graphene or mesoporous carbons have been evaluated for CO₂ capture, often requiring surface oxidation or amine grafting to impart sufficient chemical affinity. While these approaches offer potential, their catalytic activity is often indirect, depending on textural modulation rather than active chemical participation in the desorption process.

Recent efforts have also explored hybrid materials, such as ionic liquid-functionalized oxides and metal nanoparticle composites, to create synergistic surface interactions [13]. For example, IL@SiO₂ hybrids can modulate local microenvironments and promote CO₂ diffusion, though their cost and long-term stability remain concerns [14,15]. Despite these advances, many existing materials exhibit limitations including poor durability, complex synthesis procedures, or reliance on high surface area rather than chemically active sites.

Therefore, a catalyst that integrates mesoporosity, thermal stability, and synergistic surface acid–base properties is highly desirable for amine-based CO₂ capture. Among potential candidates, iron oxyhydroxides—particularly the β-phase of FeOOH—remain underexplored, despite their hydroxyl-rich surfaces and potential to provide Brønsted/Lewis acidic sites similar to those reported for AlOOH and TiO(OH)₂. β-FeOOH exhibits a mesoporous architecture, excellent thermal robustness, and coexisting Brønsted and Lewis acid sites, suggesting its promise as a bifunctional catalyst for facilitating CO₂ desorption [16]. Despite these advances, the potential of FeOOH remains under-explored, particularly regarding the synergistic effect of its crystalline architecture and active site distribution. Unlike the relatively dense structures of conventional metal oxides or the layered arrangements in some hydroxides, β-FeOOH possesses a distinctive 2 × 2 tunnel structure [17,18]. This framework provides an optimized platform for high-density surface hydroxyl groups and, more importantly, facilitates superior proton mobility through its internal channels—a critical requirement for the deprotonation of BH⁺ and the subsequent breakdown of carbamate. Furthermore, the strong Lewis acidity of the Fe³⁺ centers in FeOOH offers a more potent electronic effect in polarizing the C-N bond compared to Al- or Ti-based materials. Such structural and electronic advantages, combined with its earth-abundance and low cost, position FeOOH as a more commercially viable and efficient candidate for large-scale CO₂ capture applications.

Herein, we report the synthesis, characterization, and evaluation of mesoporous β-FeOOH as a dual-function catalyst for CO₂ capture in 30 wt% MEA solutions. The catalyst was prepared via hydrothermal precipitation and characterized using BET, XRD, TEM, Py-IR, and XPS techniques to elucidate its textural and surface acid–base properties. A custom-designed bubbling reactor was used to assess CO₂ absorption and desorption performance at various FeOOH loadings. Our findings reveal that β-FeOOH not only improves CO₂ uptake capacity and absorption time but also significantly accelerates desorption kinetics, reducing regeneration time and increasing CO₂ recovery. This study demonstrates the catalytic potential of mesoporous FeOOH in improving the efficiency of amine-based carbon capture and provides mechanistic insights into its surface-mediated enhancement. This work aims not only to improve catalytic performance but also to address the sustainability challenges associated with energy-intensive solvent regeneration in CO₂ capture systems.

2. Materials and Methods

2.1. Chemicals

Iron (III) chloride hexahydrate (FeCl₃⋅6H₂O, 99%) and monoethanolamine (MEA, ACS, ≥99.0%) were purchased from Macklin Inc., Shanghai, China. The deionized water (DI water) was self-prepared in the laboratory with a resistivity greater than 18 MΩ·cm. Carbon dioxide gas (CO₂, 99.99%) and nitrogen gas (N₂, 99.99%) were provided by Zhejiang Nankai Gas Co., Ltd., Zhejiang, China. All chemicals and materials were used as received without further purification.

2.2. Preparation of β-FeOOH

The β-FeOOH was synthesized using a hydrolysis method. The β-FeOOH catalyst was prepared by dissolving 1.62 g of FeCl₃⋅6H₂O in 150 mL of DI-water and stirring the solution at 0 °C for 15 h. The resulting suspension was centrifuged, washed three times with deionized water, and dried at 60 °C for 6 h to yield the final product.

2.3. Characterizations

Physical characterization methods were performed using scanning electron microscopy (SEM) with a ZEISS Sigma instrument (ZEISS Group, Jena, German) to observe surface morphology and microstructure of the sample. Brunauer–Emmett–Teller (BET) analysis using an Anton Paar Nova 800 instrument (Anton Paar GmbH, Graz, Austria) was employed to measure specific surface area and pore size distribution. Samples were degassed at 120 °C for 12 h under vacuum (<10⁻² mbar). BET surface area was calculated in the P/P₀ range 0.05–0.30; pore-size distribution was derived from the Barrett–Joyner–Halenda (BJH) desorption branch assuming cylindrical pores. Thermogravimetric analysis (TG) was conducted using a TA Q600 (TA-Instruments, New Castle, DE, USA) under a N₂ at a flow rate of 100 mL/min, with a temperature range from room temperature to 800 °C and a heating rate of 10 °C/min, to determine thermal stability and decomposition behavior.

Chemical characterization methods included Fourier-Transform Infrared Spectroscopy (FTIR, Shimadzu Corporation, Kyoto, Japan) using a Shimadzu IRTracer-100 with a MIR TGS detector, with a resolution of 4 cm⁻¹, scanning from 400 to 4000 cm⁻¹, to analyze functional groups and chemical bonds. X-ray Photoelectron Spectroscopy (XPS, Thermo Scientific, Waltham, MA, USA) was performed using Thermo Scientific K-Alpha with a monochromatic Al Kα source (E = 1486.68 eV) to study surface chemical composition, oxidation states, and elemental distribution. X-ray Diffraction (XRD, Rigaku Corporation, Tokyo, Japan) using a Rigaku SmartLab SE, with a scanning range from 20° to 90° at a rate of 5°/min was applied to identify crystalline phases, crystallinity, and structural properties of the sample.

2.4. Catalyst Performance Test

The CO₂ capture experiments were conducted in a 30% MEA aqueous solution, as shown in Figure 1, using a 250 mL three-necked glass flask to simulate a batch reactor. In these experiments, 100 mL of a 30% MEA solution and the pre-calculated β-FeOOH catalyst were placed in the reactor. Before initiating CO₂ absorption, the reactor was placed in an ultrasonic bath for 10 min to ensure the complete dispersion of the catalyst in the solution. A simulated flue gas mixture containing 20% CO₂ and 80% N₂ was bubbled into the reactor at a flow rate of 800 mL/min. The simulated flue gas entered the reactor at 25 °C to mimic post-combustion flue gas after water knock-out. The outlet gas was monitored and recorded using a gas analyzer (GT-903, Shenzhen Korno Electronic Technology Co., Ltd., Shenzhen, China), and the absorption process was kept for 2 h. For the CO₂ desorption process, the MEA solution was gradually heated to 85 °C. N₂ at a flow rate of 500 mL/min served as the carrier gas, and the CO₂ concentration in the outlet gas was continuously monitored and recorded by the gas analyzer. The desorption process lasted for 1 h, and then the gas valves, heating, and stirring were turned off.

During the CO₂ capture process, key parameters such as the CO₂ absorption and desorption rates, the total amounts of CO₂ absorbed and desorbed, and the change in the desorption rate over time were used to evaluate the catalyst’s performance. The CO₂ absorption rate was calculated using the following equation:

$${CO}_2 absorption rate=\left({\displaystyle \frac{C_{in}-C_{out}}{C_{in}}}\right)\times 100\%$$

where $C_{in}$ is the CO₂ concentration at the inlet, and $C_{out}$ is the CO₂ concentration at the outlet.

The CO₂ desorption rate was calculated using the equation:

$${CO}_2 desorption rate=Q_{Total}\times C_{out}\times {\displaystyle \frac{1}{wt (MEA solution)}}$$

where $Q_{Total}$ is the total desorption flow rate and calculated using the following equation:

$$Q_{Total}={\displaystyle \frac{Q_{N_2}}{100-C_{out}}}$$

where $Q_{N_2}$ is the N₂ served as the carrier gas. The increase in the desorption rate was calculated using:

$$Increase in desorption rate=\left({\displaystyle \frac{R_{cat}-R_{no-cat}}{R_{no-cat}}}\right)$$

where $R_{cat}$ is the desorption rate in the presence of the catalyst, and $R_{no-cat}$ is the desorption rate without the catalyst.

3. Results and Discussion

3.1. Properties of Prepared FeOOH

The morphology and elemental distribution of the synthesized FeOOH, intended for application as a catalyst in MEA-based CO₂ capture systems, were characterized using scanning electron microscopy (SEM), energy-dispersive X-ray spectroscopy (EDS), and elemental mapping, as presented in Figure 2. The low-magnification SEM image (Figure 2a) reveals that the FeOOH particles are agglomerated, irregular in shape, and vary in size, suggesting a significant surface area. At higher magnification (Figure 2b), the particles exhibit a porous microstructure, which are advantageous for catalytic and adsorption processes due to enhanced surface interactions. The elemental mapping (Figure 2c) and EDS spectrum (Figure 2d) confirm the presence of Fe and O as the primary elements, with weight percentages of 65.6% Fe and 34.4% O. Furthermore, the Fe (Figure 2e) and O (Figure 2f) elemental maps distribution maps further reveal a homogeneous spatial dispersion of both elements, confirming the uniform composition of the synthesized FeOOH. These morphological and compositional characteristics collectively support the material’s suitability for use in catalytically active systems.

As shown in Figure 3a, the synthesized FeOOH particles exhibit an average diameter of approximately 3.1 μm. As shown by TG curve in Figure 3b, the prepared FeOOH exhibits an obvious weight loss of 16.8 wt% due to vaporization of surface-adsorbed water, observed under 150 °C. Previous research has demonstrated that water adsorption on catalyst surfaces enhances catalytic activity by facilitating reaction mechanisms. Thus, a higher surface water content may contribute to increased catalytic efficiency. Furthermore, TG analysis reveals no significant mass loss above 400 °C, confirming the material’s thermal stability. The continuous weight loss observed up to 800 °C further confirms the dehydration of β-FeOOH, leading to the formation of Fe₂O₃ [19]. The DSC results are consistent with this weight loss trend, showing endothermic heat absorption corresponding to the dehydration process.

The N₂ adsorption–desorption isotherms (Figure 4a) exhibit a Type IV isotherm with a pronounced hysteresis loop, indicating the mesoporous nature of the prepared FeOOH material [20]. In liquid MEA, the catalytically relevant interactions are expected to occur primarily on the external surface and within mesopores, where solvated MEA and ionic species (e.g., HCO₃⁻/carbamate) can readily access active sites. Micropores, if present, are likely to contribute less due to steric and diffusion constraints in the viscous, strongly solvating MEA solution. The slight N₂ uptake at low relative pressures (P/P₀ < 0.2) suggests the presence of micropores, while the sharp increase in adsorption at higher pressures (P/P₀ > 0.4) confirms the dominance of mesopores. The specific surface area, calculated using the Brunauer–Emmett–Teller theory, is 287 m²/g, providing abundant active sites for CO₂ adsorption. The large surface area, combined with the presence of surface OH groups, is a key factor in improving the catalytic performance of FeOOH as a CO₂ capture catalyst in MEA solution [8]. The pore size distribution (Figure 4b) reveals a narrow and uniform peak within the mesopore range (2–50 nm), further confirming the well-controlled pore structure. This well-developed mesoporous architecture and high surface area facilitate efficient CO₂ diffusion, adsorption, and activation, making FeOOH a promising catalyst for CO₂ capture applications in MEA solution.

FT-IR reveals crucial information about the functional groups present in FeOOH, as shown in Figure 5a. The broad absorption band observed around 3400–3600 cm⁻¹ indicates the presence of hydroxyl groups (-OH), which are essential for CO₂ capture [21,22]. A strong peak near 1618 cm⁻¹ corresponds to bending vibrations of hydroxyl groups [23]. Peaks in the 1130 cm⁻¹ is attributed to Fe-O stretching vibrations, confirming the structural integrity of FeOOH [24]. The presence of hydroxyl groups is particularly significant as they facilitate CO₂ absorption by providing active sites for CO₂ molecules, enhancing the efficiency of the capture process. The identification of hydroxyl groups implies that FeOOH can interact with CO₂ through proton transfer mechanisms, contributing to improved CO₂ capture efficiency in systems like MEA.

As shown in Figure 5b, Pyrolysis-Infrared Spectroscopy (Py-IR) is employed to detect Lewis acid and Brønsted acid sites on the surface of FeOOH. The peak observed at 1450 cm⁻¹ is associated with Lewis acid sites (denoted as L), while the peak at 1540 cm⁻¹ corresponds to Brønsted acid sites (denoted as B) [25]. A peak observed at 1490 cm⁻¹ indicates the simultaneous presence of both Brønsted and Lewis acid sites. These results suggest that FeOOH exhibits a combination of acid sites, contributing to enhanced CO₂ adsorption through both proton donation (Brønsted acid sites) and acceptor mechanisms (Lewis acid sites) [26]. The presence of these acidic sites will improve the interaction between FeOOH and CO₂, enhancing its catalytic performance for CO₂ capture and desorption processes.

The XRD pattern (Figure 6a) confirms the crystalline phase of FeOOH, indicating the material adopts an orthorhombic structure. The β-FeOOH phase shows distinct diffraction peaks at 26.7°, 34.0°, 35.2°, 39.2°, 46.5°, 55.9°, 61.3°, 64.2°, and 67.9°, corresponding to crystal lattice planes (310, 400, 211, 420, 411, 600, 541, and 604). These peaks align well with the reference pattern of akaganeite (JCPDS 75-1549), confirming the β-FeOOH phase [27]. However, the peaks appear relatively broad, which is consistent with small crystallites and/or microstrain. A quantitative crystallite size evaluation requires peak-width (FWHM) analysis with instrumental broadening correction. Despite the broadness, the overall crystalline nature contributes to the stability of FeOOH, providing a solid structural foundation for its catalytic performance. These structural properties are essential for its application in CO₂ capture processes.

The chemical composition and surface states of the prepared FeOOH were further investigated using X-ray Photoelectron Spectroscopy (XPS), as shown in Figure 6b–d. The XPS survey spectrum (Figure 6b) reveals the dominant presence of Fe and O, indicating high purity with no significant impurity peaks. The Fe 2p high-resolution spectrum (Figure 6c) displays two prominent peaks at approximately 711 eV (Fe 2p_3/2) and 724 eV (Fe 2p_1/2), along with satellite peaks, confirming the presence of Fe³⁺ in the FeOOH phase [28]. These binding energies are consistent with the formation of FeOOH, which is essential for catalytic activity in CO₂ capture. The O 1s high-resolution spectrum (Figure 6d) shows two main peaks at approximately 531.5 eV, attributed to hydroxyl groups (–OH), and 529.8 eV, assigned to lattice oxygen (O²⁻) [29,30]. The presence of these surface hydroxyl groups is particularly significant as they enhance CO₂ adsorption and activation, contributing to the catalytic performance of FeOOH. Overall, the XPS analysis highlights the presence of Fe³⁺ and surface hydroxyl groups, both of which are critical for CO₂ capture applications. The FeOOH structure exhibits chemical characteristics that support its efficiency in catalytic CO₂ absorption processes.

3.2. Catalytic CO₂ Capture Performance

In this study, a 30 wt% MEA solution was used as the control (blank) sample, while FeOOH catalysts were added at various concentrations (0.1, 0.2, 0.4, and 0.8 wt%) to evaluate their effects on CO₂ absorption and desorption behavior. Figure 7a illustrates the CO₂ concentration in the reactor’s exhaust gas over time, which reflects the portion of CO₂ not captured by the MEA solution. Initially, the CO₂ concentration rapidly dropped from 0.11 mmol/s to 0 mmol/s, indicating efficient absorption. It remained at zero for approximately 2000 s, signifying complete CO₂ removal during this stage. Thereafter, the concentration rose sharply between 2000 and 3500 s, reflecting the depletion of the MEA absorption capacity. In the presence of FeOOH, the outlet CO₂ concentration remained lower throughout the process, indicating enhanced absorption performance. This improvement can be attributed to increased mass transfer efficiency facilitated by the dispersed catalyst particles. Following the peak period, the absorption rate slowed down and entered a steady-state phase, governed primarily by reaction kinetics. Under lower MEA availability, the CO₂ capture rate decreased accordingly. Therefore, the enhancement observed in this study likely arises from a combined effect of surface catalytic reactions and improved gas–liquid mass transfer, with catalytic proton-transfer pathways playing a dominant role under the investigated conditions.

As defined by the International Energy Agency, the effective absorption time refers to the duration required to reduce the outlet CO₂ concentration to 10% of its initial value [31]. According to Figure 7a, the blank sample exhibited an effective absorption time of 3457 s with a total absorbed amount of 284.62 mmol. In comparison, FeOOH-modified samples achieved extended absorption times of 3397 s (0.1 wt%), 3572 s (0.2 wt%), 3754 s (0.4 wt%), and 3620 s (0.8 wt%), with corresponding CO₂ uptake values of 281.82, 290.00, 296.70, and 291.37 mmol, respectively. These results indicate that catalyst incorporation improves both absorption time and capacity. Figure 7b presents the total amount of CO₂ absorbed over the test period. The blank sample absorbed 287.95 mmol, while catalyst-modified samples absorbed 285.12 (0.1 wt%), 296.88 (0.2 wt%), 303.17 (0.4 wt%), and 298.28 mmol (0.8 wt%). These results indicate that FeOOH primarily enhances the kinetic performance of CO₂ absorption, accelerating the approach to equilibrium, while the thermodynamic equilibrium capacity of the MEA system remains essentially unchanged. At a higher catalyst loading of 0.8 wt%, a slight decline in performance was observed. This behavior may be related to increased mass-transfer resistance and reduced effective utilization of catalytic sites, potentially arising from particle agglomeration, partial blockage of the gas–liquid interface, or local heterogeneity in catalyst dispersion. Further quantitative studies on slurry rheology and dispersion stability are required to clarify the dominant factors. Among all loadings, 0.4 wt% consistently yielded the best performance in both absorption duration and total uptake. This enhancement is attributed to surface-mediated CO₂ adsorption by FeOOH, which increases the local concentration of reactants near the gas–liquid interface.

The CO₂ desorption performance under the influence of FeOOH catalysts at different weight loadings (blank, 0.1 wt%, 0.2 wt%, 0.4 wt%, and 0.8 wt%), presented in Figure 8. From Figure 8a, CO₂ desorption exhibits an initial increase followed by a gradual decline. With the addition of catalysts, CO₂ desorption occurs earlier compared to the blank sample. The desorption time is reduced sequentially from 0.1 wt% to 0.4 wt% catalyst loading and slightly delayed at 0.8% but still greater than the blank sample. This improvement is mainly due to the catalytic sites provided by FeOOH—specifically Lewis acid and Brønsted acid sites. The Brønsted acid sites facilitate proton transfer, accelerating the decomposition of carbamates, while Lewis acid sites enhance proton transfer during the protonated amine process, thus boosting CO₂ desorption rates [32,33]. For the blank sample without catalysts, the peak desorption time occurs at 1496 s, while with catalysts, this peak time decreases to 1484 s, 1440 s, and 1335 s. The peak area increases and then decreases, indicating that the catalysts enhance the desorption process up to a certain concentration.

Figure 8b shows that the amount of CO₂ desorbed from the blank sample is 27.03 mmol. With catalyst additions of 0.1, 0.2, and 0.4 wt%, the desorbed amounts increase to 27.28 mmol (1.03%), 27.79 mmol (2.8%), and 29.97 mmol (10.9%), respectively, indicating enhanced diffusion and mass transfer at lower catalyst concentrations. However, at 0.8 wt% loading, the desorbed amount drops to 24.1 mmol due to particle agglomeration, which increases solution viscosity and hinders desorption efficiency—resulting in lower performance than the blank. However, higher catalyst loadings (beyond 0.4 wt%) lead to increased apparent viscosity and bubble coalescence, which reduces the effective gas–liquid interfacial area and counteracts the kinetic benefits of the additional active sites. Thus, 0.4 wt% is identified as the optimal catalyst loading.

Notably, mesoporous β-FeOOH, with a BET surface area of 287 m²/g, achieves a 10.9% increase in CO₂ desorption from a 30 wt% MEA solution at 85 °C. This enhancement surpasses that of previously reported catalysts, including TiO(OH)₂ nanosheets (180 m²/g, +6.2%), Fe-BEA zeolite (520 m²/g, +8.4%), and AlOOH nanorods (210 m²/g, +5.5%) [8,9,10]. Despite its moderate surface area, β-FeOOH exhibits superior activity, highlighting that the density and type of Brønsted/Lewis acid sites, rather than surface area alone, play a dominant role in determining catalytic performance under the tested conditions. It should be noted that the addition of FeOOH particles may also influence hydrodynamic factors such as foam formation and gas holdup during bubbling, which could affect the apparent absorption rate.

Figure 9 illustrates the comparison of CO₂ desorption rates with and without the assistance of the prepared FeOOH catalyst at the optimal loading of 0.4 wt%. The X-axis represents time, while the Y-axis indicates the CO₂ desorption rate. A prominent peak at approximately 380% demonstrates a significant enhancement in desorption rate with the FeOOH catalyst, compared to the baseline without the catalyst. This sharp peak, followed by a decline, suggests the catalyst enables a fast initial reaction, likely due to the availability of active sites, before stabilizing or potentially depleting over time. The remarkable increase in desorption rate highlights the catalytic efficiency of FeOOH, which not only accelerates the desorption process but also dramatically reduces the energy consumption required for CO₂ desorption. These findings emphasize the potential of the FeOOH catalyst in improving the efficiency and sustainability of CO₂ capture and conversion processes. It should be noted that the enhanced CO₂ desorption reported here is evaluated based on the outlet CO₂ concentration in the gas phase. The present study does not explicitly distinguish between selective CO₂ regeneration and possible solvent degradation pathways such as heat-stable salt formation or amine loss.

3.3. Stability Study

The cyclic CO₂ absorption–desorption performance of the 0.4 wt% FeOOH catalyst was tested to evaluate its efficiency and stability over multiple cycles as shown in Figure 10. Initially, the absorption capacity showed robust performance, but with repeated cycles, a gradual decline was observed, due to material degradation or inherent limitations of the cyclic process. However, after several cycles, the absorption capacity stabilized, indicating the catalyst’s ability to maintain consistent performance under the given conditions. The desorption process was also analyzed, showing insights into the completeness of CO₂ release, which is critical for practical applications in carbon capture. The effect of 0.4 wt% FeOOH loading on performance was also highlighted, suggesting the need to optimize the loading level for better long-term efficiency. Regarding the potential for solvent degradation, although some metal ions (especially dissolved iron) are known to catalyze MEA oxidation, no significant degradation was observed in our FeOOH/MEA system during the stability tests. This can be attributed to the heterogeneous nature of the FeOOH catalyst, where iron species remain integrated within the stable oxyhydroxide crystalline framework rather than being released as free ions into the bulk solution.

It is worth noting that the relatively mild regeneration temperature at 85 °C, compared to conventional conditions, suppresses both MEA degradation and potential catalyst instability. Under these conditions, β-FeOOH is expected to remain largely in its solid form due to its limited solubility in MEA solutions. XPS results confirm that iron is retained in the Fe³⁺ state, indicating preservation of the catalyst structure after repeated cycles. Although minor iron leaching cannot be fully excluded without quantitative analysis, the stable cyclic performance suggests that FeOOH predominantly operates as a heterogeneous catalyst without significant loss of active species. Overall, these results highlight its potential for cyclic CO₂ capture, while further work is needed to fully assess long-term durability.

3.4. DFT Analysis and Catalytic Mechanism

To provide energetic insight into the proposed proton-transfer-assisted desorption mechanism, density functional theory (DFT) calculations were performed on representative FeOOH surface models as shown in Figure 11. The calculations focus on the interaction between FeOOH and MEA-related species, as well as the role of Brønsted and Lewis acid sites in stabilizing intermediates and facilitating proton transfer. The optimized structures indicate that surface hydroxyl groups and coordinatively unsaturated Fe sites on FeOOH can form strong hydrogen bonding and coordination interactions with MEA-derived species, with calculated H-bond lengths typically observed in the range of 1.8 to 2.1 Å.

These interactions promote proton transfer from surface –OH groups to bicarbonate (HCO₃⁻) or carbamate (MEACOO⁻) species, facilitating the formation of unstable intermediates that decompose more readily to release CO₂. The calculated adsorption energies show that MEA-related intermediates are more strongly stabilized on acidic FeOOH surfaces compared to non-catalyzed environments, indicating a favorable thermodynamic driving force for surface-assisted activation. Specifically, the energy barrier associated with the rate-limiting proton transfer step is significantly reduced by approximately $15–20 kJ/mol in the presence of FeOOH. This substantial reduction in the activation energy barrier suggests that the catalyst provides a lower-energy pathway for CO₂ desorption. These results support the experimental observation in Section 3.2, where the 0.4 wt% FeOOH catalyst successfully reduced the peak desorption time from 1496 s to 1335 s and increased the total CO₂ desorption by 10.9%. Therefore, the enhancement in desorption kinetics can be rationalized by a feasible catalytic pathway from an energetic standpoint rather than simple physical adsorption effects. It should be noted that the present DFT calculations are based on simplified surface models and do not explicitly account for solvent effects, and therefore provide qualitative rather than quantitative insights.

The possible catalytic CO₂ reaction mechanism of FeOOH catalyst is shown in Figure 12. The absorption and desorption of CO₂ in aqueous MEA solutions proceed primarily via the zwitterion mechanism, as originally proposed by Caplow and Danckwerts [34,35]. During the initial stage of absorption, CO₂ reacts with MEA to form a zwitterionic intermediate (MEA⁺COO⁻), which subsequently interacts with another MEA molecule to generate carbamate (MEACOO⁻) and a protonated amine (MEAH⁺), as shown in reactions (R1) and (R2). Under conditions of high CO₂ loading, the carbamate species undergo hydrolysis, regenerating free MEA and forming bicarbonate ions (HCO₃⁻), as indicated in reaction (R3). These fundamental steps represent the core chemistry of CO₂ capture in MEA systems. However, the thermal desorption of CO₂ is hindered by slow reaction kinetics and high energy requirements, necessitating the use of a suitable catalyst.

β-FeOOH, possessing a large specific surface area (287 m²/g) and bifunctional acid–base character, mainly promotes the kinetics of CO₂ absorption and desorption, thereby shortening the time required to reach equilibrium, rather than significantly changing the equilibrium CO₂ capacity of the MEA system. Its role is to provide a heterogeneous acid–base interface that facilitates proton transfer and intermediate stabilization during the MEA–CO₂ reaction process. As illustrated in Reactions (R1) and (R2), FeOOH does not participate as a reactant in carbamate formation but assists this process through surface-mediated interactions. The Brønsted acid sites can facilitate proton donation and transfer, while the Lewis acidic/basic sites can stabilize electron-rich or electron-deficient species, promoting charge redistribution during the conversion between MEA, CO₂, and carbamate species. Through these surface-assisted acid–base interactions, FeOOH lowers the kinetic barriers associated with proton transfer steps and intermediate transformations, leading to enhanced absorption efficiency without altering the thermodynamic equilibrium [9].

$$\mathrm{MEA}+{\mathrm{CO}}_2\leftrightarrow {\mathrm{MEA}}^{+}{\mathrm{COO}}^{-}$$

(R1)

$${\mathrm{MEA}}^{+}{\mathrm{COO}}^{-}+\mathrm{MEA}\leftrightarrow {\mathrm{MEAH}}^{+}+{\mathrm{MEACOO}}^{-}$$

(R2)

$${\mathrm{MEACOO}}^{-}+{\mathrm{H}}_2\mathrm{O}\leftrightarrow \mathrm{MEA}+{{\mathrm{HCO}}_3}^{-}$$

(R3)

In the desorption phase, uncatalyzed pathways rely on the thermal decomposition of HCO₃⁻ and MEACOO⁻, which are generally slow and energy-intensive. FeOOH facilitates the breakdown of both bicarbonate (HCO₃⁻) and zwitterionic intermediates (MEA⁺COO⁻), accelerating the release of CO₂. At the onset of desorption, when CO₂ loading is high, HCO₃⁻ is the dominant source of CO₂, as outlined in reaction (R3). At 85 °C, FeOOH participates in proton generation either by abstracting OH⁻ from water to form H⁺ (R4) or by deprotonating water to yield H₃O⁺ (R5). These protons subsequently interact with HCO₃⁻ to form carbonic acid (H₂CO₃), which rapidly decomposes upon heating, as described in reactions (R6) and (R7). Compared to the uncatalyzed route, this pathway is more kinetically favorable. As desorption proceeds and CO₂ loading decreases, MEACOO⁻ becomes the primary source of CO₂. Brønsted acid sites in FeOOH facilitate proton transfer to MEACOO⁻, forming MEA⁺COO⁻, which then thermally decomposes to yield CO₂ and regenerate MEA, as depicted in reactions (R8) and (R9). These steps are further supported by DFT calculations, which show that the presence of FeOOH significantly lowers the energy barrier for proton transfer and intermediate decomposition, providing a kinetically more favorable pathway for CO₂ desorption.

$$\mathrm{FeOOH}+{\mathrm{H}}_2\mathrm{O}\leftrightarrow {{\mathrm{Fe}\left(\mathrm{OH}\right)}_2}^{-}+{\mathrm{H}}^{+}$$

(R4)

$$\mathrm{FeOOH}+{\mathrm{H}}_2\mathrm{O}\leftrightarrow {\mathrm{FeOO}}^{-}+{\mathrm{H}}_3{\mathrm{O}}^{+}$$

(R5)

$${\mathrm{H}}^{+} (\mathrm{or} {\mathrm{H}}_3{\mathrm{O}}^{+})+{{\mathrm{HCO}}_3}^{-}\leftrightarrow {\mathrm{H}}_2{\mathrm{CO}}_3$$

(R6)

$${\mathrm{H}}_2{\mathrm{CO}}_3\leftrightarrow {\mathrm{H}}_2\mathrm{O}+{\mathrm{CO}}_2$$

(R7)

$$\mathrm{FeOOH}+{\mathrm{MEACOO}}^{-}\leftrightarrow {\mathrm{FeOO}}^{-}+{\mathrm{MEA}}^{+}{\mathrm{COO}}^{-}$$

(R8)

$${\mathrm{MEA}}^{+}{\mathrm{COO}}^{-}\leftrightarrow \mathrm{MEA}+{\mathrm{CO}}_2$$

(R9)

The catalytic cycle is maintained through regeneration of active intermediates via reactions with protonated amines or bicarbonate, as shown in reaction (R10). This self-sustaining sequence allows FeOOH to preserve its catalytic functionality over multiple cycles of operation.

$${\mathrm{FeOO}}^{-}+{\mathrm{MEAH}}^{+}\leftrightarrow \mathrm{FeOOH}+\mathrm{MEA}$$

(R10)

Overall, the proposed mechanism highlights how FeOOH lowers the activation energy of both absorption and desorption processes by providing proton donors and acceptors via Brønsted and Lewis acid sites. This dual-functionality plays a key role in enhancing CO₂ capture efficiency while reducing thermal energy input, as evidenced by consistent performance over 10 consecutive cycles.

4. Conclusions

This study demonstrates that incorporating β-FeOOH into MEA-based systems effectively enhances CO₂ capture by improving both absorption and desorption kinetics. The catalyst exhibits a high surface area (287 m²/g) and abundant Brønsted/Lewis acid sites, which synergistically facilitate proton transfer and carbamate decomposition. At an optimal loading of 0.4 wt%, the absorption time increased to 3754 s with a total uptake of 303.17 mmol, while the desorption amount increased by 10.9% to 29.97 mmol, and the peak desorption time was reduced from 1496 s to 1335 s. Beyond performance enhancement, this work highlights the sustainability advantages of the FeOOH-promoted system. The ability to achieve efficient CO₂ desorption at a relatively low temperature (85 °C) significantly reduces the energy intensity of solvent regeneration, offering opportunities for the utilization of low-grade industrial waste heat and lowering the overall carbon footprint of the process. In addition, the use of earth-abundant, low-cost iron-based materials enhances resource sustainability and economic feasibility, while the demonstrated cyclic stability supports long-term operational reliability. Overall, β-FeOOH provides a practical and scalable pathway toward energy-efficient CO₂ capture, integrating catalytic performance with reduced energy demand and material sustainability. This work contributes to the development of more sustainable carbon capture technologies and supports ongoing efforts toward carbon-neutral and low-carbon industrial systems.