Mechanistic Study of CO₂ Absorption in Alkanolamine Solutions Based on Density Functional Theory

Abstract

Among the various CO₂ capture technologies, chemical absorption is currently one of the most widely applied methods in industrial practice. In this study, density functional theory was employed to investigate the reaction mechanisms of CO₂ absorption by typical alkanolamine solvents. Reaction pathways between CO₂ and four representative alkanolamines—monoethanolamine (MEA), diethanolamine (DEA), triethanolamine (TEA), and methyldiethanolamine (MDEA)—were constructed and analyzed. By evaluating the activation energy barriers of different amines, the thermodynamic characteristics and reaction feasibility of the CO₂ absorption process were systematically elucidated. The results show that the primary amine MEA exhibits the lowest activation energy barrier (32.02 kJ/mol), indicating the most favorable reaction kinetics, while the secondary amine DEA shows a slightly higher barrier of 47.35 kJ/mol. As tertiary amines, TEA and MDEA exhibit significantly higher activation energy barriers, indicating slower reaction kinetics; however, they generally possess higher CO₂ loading capacities and less stable reaction products, which facilitate solvent regeneration. The activation energy barriers of MDEA and TEA were calculated to be 54.53 kJ/mol and 94.17 kJ/mol, respectively, indicating that MDEA reacts more readily with CO₂ than TEA.

1. Introduction

Research on CO₂ absorption by alkanolamine solutions primarily originates from the growing concerns over climate change mitigation and greenhouse gas emissions [1]. CO₂ is one of the major greenhouse gases, and its emissions are widely recognized as a key driver of global climate change. The combustion of fossil fuels releases large amounts of CO₂. Among these, coal is extensively used worldwide for power generation and is characterized by a higher carbon content and lower hydrogen content, leading to greater CO₂ emissions upon combustion. Therefore, coal-fired processes are considered one of the dominant sources of CO₂ emissions in fossil fuel utilization [2]. Capturing and treating flue gas from thermal power plants is thus regarded as a critical pathway for reducing carbon emissions. Consequently, significant efforts have been devoted to developing efficient CO₂ capture technologies to lower atmospheric CO₂ concentrations [3].

Among various carbon capture technologies, chemical absorption using liquid solvents has emerged as one of the most promising approaches for industrial applications due to its high separation efficiency, operational stability, and technological maturity [4].

The strong research interest in alkanolamine-based solvents stems from their high affinity for CO₂. When CO₂ passes through aqueous alkanolamine solutions, it reacts chemically with amines to form carbonate or bicarbonate species, thereby transferring CO₂ from the gas phase into the liquid phase [5]. Owing to its high absorption efficiency and excellent industrial applicability, this technology has been widely implemented in post-combustion CO₂ capture processes in power plants and chemical industries [6].

In recent years, increasing attention has been paid to integrating alkanolamine absorption technology with carbon capture, utilization, and storage (CCUS) systems to further reduce greenhouse gas emissions and mitigate climate change impacts. Due to their compatibility with existing industrial infrastructure and high level of technological maturity, alkanolamine solvents are considered one of the most promising near-term solutions for large-scale CO₂ capture [7]. Therefore, studies on CO₂ absorption by alkanolamine solutions are of great environmental significance and practical engineering value.

Bishnoi et al. [8] systematically investigated and modeled the absorption kinetics, solubility, and mass transfer behavior of piperazine (PZ)-activated MDEA systems, providing an experimental foundation for the industrial application of activated MDEA solvents. Danckwerts [9] presented an early comprehensive review of the reaction mechanisms and kinetics of CO₂ with MEA and other alkanolamines, establishing a relatively complete intermediate-reaction framework for engineering applications. Samanta et al. [10] studied the absorption rates and reaction mechanisms of PZ/MDEA systems under conditions close to industrial operation using wetted-wall column and stirred-cell (disk) apparatus, and compared the effects of different activator concentrations on reaction rates. Sema et al. [11] conducted systematic experimental studies on the coupling between mass transfer and chemical kinetics in MDEA–MEA blended solvent systems, and proposed rate models suitable for process simulation, demonstrating the advantages of mixed alkanolamine systems over single-component solvents.

Hwang [12] revisited and refined the key reaction pathways of the MEA/CO₂ system using density functional theory (DFT) and ab initio molecular dynamics (AIMD), explicitly accounting for the aqueous environment. This work clarified the transition states and the influence of solvent effects on the free-energy surface, marking a shift in the field from empirical descriptions toward molecular-level mechanistic understanding. Putta [13] proposed an activity-based kinetic model capable of more accurately predicting absorption rates near equilibrium conditions, thereby overcoming the limitations of conventional models at high loadings and high amine concentrations. Orestes et al. [14] systematically investigated the interactions between CO₂ and amine molecules using DFT, focusing on the stability of zwitterionic intermediates and the influence of different exchange–correlation functionals on reaction energy calculations, thus providing a theoretical evaluation of reaction pathways and the applicability of computational methods.

The use of density functional theory (DFT) simulations to evaluate the feasibility and reaction difficulty of CO₂ absorption by different alkanolamines has become an important research focus in recent years. In wet absorption processes, aqueous solutions of monoethanolamine (MEA), diethanolamine (DEA), triethanolamine (TEA), and methyldiethanolamine (MDEA) are the most widely used solvent systems. These compounds possess advantages such as simple molecular structures, low cost, wide availability, and high reactivity toward CO₂, and are therefore extensively applied in both single-component and blended alkanolamine systems [15].

At present, research on CO₂ capture using alkanolamine solutions mainly focuses on multicomponent blended systems, with experimental studies dominating the field. These studies primarily investigate the effects of different components on CO₂ absorption capacity and absorption rate [16]. However, systematic investigations of the reaction mechanisms between single-component alkanolamines and CO₂ at the molecular level remain relatively limited, particularly with respect to reaction pathways and energy barrier distributions. A detailed understanding of the reaction mechanisms in single-amine systems can provide essential theoretical insights into the synergistic effects observed in blended amine systems, thereby elucidating the roles of different amines and activators in CO₂ absorption processes [17].

To address these issues, this study employs the DMol³ module within the Materials Studio package and applies density functional theory (DFT) to systematically investigate the microscopic reaction mechanisms of MEA, DEA, TEA, and MDEA with CO₂. By constructing the geometries of reactants, transition states, and intermediates, and combining transition state searches with intrinsic reaction coordinate (IRC) analysis, the typical reaction pathways were identified. The energy changes and activation barriers of key elementary steps were also calculated [18]. In addition, solvent effects were incorporated using an implicit solvation model to simulate aqueous environments, thereby providing a more realistic representation of the actual absorption systems [19].

On this basis, the effects of different alkanolamine molecular structures on CO₂ absorption reactivity and reaction pathway characteristics were systematically compared. The fundamental differences among primary, secondary, and tertiary amines in terms of reaction mechanisms and energy barrier distributions were elucidated. Furthermore, key steps such as proton transfer and the formation of bicarbonate and carbamate species were analyzed in detail, identifying the dominant factors governing reaction rates and absorption performance. The results of this study provide important theoretical guidance for the molecular design of novel alkanolamine absorbents, optimization of reaction mechanisms, and the development of low-energy CO₂ capture technologies.

2. Materials and Methods

2.1. Computational Methods

All density functional theory (DFT) calculations were performed using the DMol³ module implemented in the Materials Studio 2019 software package (Dassault Systèmes BIOVIA, San Diego, CA, USA) [20]. The exchange–correlation interactions were described using the generalized gradient approximation (GGA) with the Perdew–Burke–Ernzerhof (PBE) functional [21]. Valence electrons were described using a double numerical plus polarization (DNP) basis set. Geometry optimizations were performed without symmetry constraints.

The convergence criteria for the self-consistent field (SCF), total energy, maximum displacement, and maximum force were set to 2.0 × 10⁻⁶ eV/atom, 2.0 × 10⁻⁵ eV/atom, 2.0 × 10⁻³ Å, and 0.05 eV/Å, respectively. All structures were fully optimized until the convergence criteria were satisfied.

The interaction energy between CO₂ and the alkanolamine system was calculated according to the following equation:

$$E_{int}=E_{complex}-E_{amine}-E_{{CO}_2}$$

where $E_{complex}$ refers to the total energy of the optimized CO₂–amine system in the presence of solvation effects, while $E_{amine}$ and $E_{{CO}_2}$ correspond to the energies of the isolated species calculated under identical computational conditions.

Solvent effects were taken into account using the conductor-like screening model (COSMO) with water as the solvent.

2.2. Models

MEA has the chemical formula HO–C₂H₄–NH₂. Owing to its fast absorption kinetics and low cost, MEA is widely used for the removal of acidic gases such as CO₂. However, it also exhibits several drawbacks, including relatively low absorption capacity, certain corrosiveness, high energy consumption during regeneration, and susceptibility to oxidation by SO₂ and O₂ present in flue gas [22]. Despite these limitations, MEA remains an important solvent in industrial applications, particularly in petroleum and natural gas processing, coal chemical industries, and ammonia synthesis. The molecular model of MEA is shown in Figure 1a.

DEA has the chemical formula [(HO–C₂H₄)₂NH]. As a secondary alkanolamine, DEA exhibits basic chemical properties and is capable of absorbing acidic gases such as CO₂ and H₂S. In aqueous solution, amines behave as weak bases and can react with CO₂ through nucleophilic addition and proton transfer to form carbamate species. In the presence of water, bicarbonate species may also be formed [23]. The molecular model of DEA is shown in Figure 1b.

TEA has the molecular formula N(CH₂CH₂OH)₃. In aqueous solution, TEA behaves as a weak base and primarily absorbs CO₂ by promoting its hydration reaction to form bicarbonate species. Unlike primary and secondary amines, tertiary amines lack N–H bonds and therefore do not form carbamate intermediates; instead, they act as base catalysts to facilitate the hydration of CO₂ in water. The molecular model of TEA is shown in Figure 1c.

MDEA has the molecular formula CH₃N(CH₂CH₂OH)₂. Similar to TEA, MDEA behaves as a weak base in aqueous solution and primarily absorbs CO₂ by promoting its hydration to form bicarbonate species. The molecular model of MDEA is shown in Figure 1d.

2.3. Reaction Mechanism

The reaction mechanism between MEA and CO₂ can be expressed as:

CO₂ + 2HO–C₂H₄–NH₂ → HO–C₂H₄–NH₃⁺ + HO–C₂H₄–NHCOO⁻

This overall reaction can be divided into two elementary steps:

CO₂ + HO–C₂H₄–NH₂ → HO–C₂H₄–NHCOO⁻ + H⁺

H⁺ + HO–C₂H₄–NH₂ → HO–C₂H₄–NH₃⁺

The reaction proceeds via a two-step mechanism. In the first step, MEA approaches the CO₂ molecule and undergoes a nucleophilic addition reaction. The nitrogen atom attacks the electrophilic carbon atom of CO₂, accompanied by proton transfer from the amine group, leading to the formation of a carbamate species (HO–C₂H₄–NHCOO⁻) and a proton (H⁺), as shown in Figure 2a.

In the second step, the released proton is rapidly captured by another MEA molecule, forming a protonated amine species (HO–C₂H₄–NH₃⁺), as illustrated in Figure 2b.

The reaction mechanism between DEA and CO₂ can be expressed as:

CO₂ + 2(HO–CH₂CH₂)₂–NH → (HO–CH₂CH₂)₂–NH₂⁺ + (HO–CH₂CH₂)₂–NCOO⁻

This overall reaction can be divided into two elementary steps:

CO₂ + (HO–CH₂CH₂)₂–NH → (HO–CH₂CH₂)₂–NCOO⁻ + H⁺

H⁺ + (HO–CH₂CH₂)₂–NH → (HO–CH₂CH₂)₂–NH₂⁺

The reaction proceeds via a two-step mechanism. In the first step, DEA approaches the CO₂ molecule and undergoes a nucleophilic addition reaction. The nitrogen atom attacks the electrophilic carbon atom of CO₂, accompanied by proton transfer from the amine group, leading to the formation of a carbamate species (HO–CH₂CH₂)₂–NCOO⁻ (as shown in Figure 3a) and a proton (H⁺).

In the second step, the released proton is rapidly captured by another DEA molecule, forming the protonated amine species (HO–CH₂CH₂)₂–NH₂⁺, as illustrated in Figure 3b.

The reaction mechanism between TEA and CO₂ can be expressed as:

CO₂ + H₂O + (HO–CH₂CH₂)₃–N →HCO₃⁻ + (HO–CH₂CH₂)₃–NH⁺

This overall reaction can be divided into three elementary steps:

CO₂ + H₂O → H₂CO₃

H₂CO₃ → H⁺ + HCO₃⁻

H⁺ + (HO–CH₂CH₂)₃–N → (HO–CH₂CH₂)₃–NH⁺

The reaction proceeds via a three-step mechanism. In the first step, CO₂ reacts with water in the weakly basic environment of the TEA solution to form carbonic acid (H₂CO₃). In the second step, carbonic acid undergoes its first dissociation in solution, producing H⁺ and HCO₃⁻. In the third step, the released proton is rapidly captured by a TEA molecule, forming an N–H bond with the nitrogen atom and yielding the protonated amine species (HO–CH₂CH₂)₃–NH⁺, as shown in Figure 4 [24].

The reaction mechanism between MDEA and CO₂ can be expressed as:

CO₂ + H₂O + (HO–CH₂CH₂)₂–NCH₃ → HCO₃⁻ + (HO–CH₂CH₂)₂–N(CH₃)H⁺

This overall reaction can be divided into three elementary steps:

CO₂ + H₂O → H₂CO₃

H₂CO₃→H⁺ + HCO₃^-

H⁺ + (HO–CH₂CH₂)₂–NCH₃ → (HO–CH₂CH₂)₂–N(CH₃)H⁺

The reaction proceeds via a three-step mechanism. In the first step, CO₂ reacts with water in the weakly basic environment of the MDEA solution to form carbonic acid (H₂CO₃). In the second step, carbonic acid undergoes its first dissociation in solution, generating H⁺ and HCO₃⁻. In the third step, the released proton is rapidly captured by an MDEA molecule, forming an N–H bond with the nitrogen atom and yielding the protonated amine species (HO–CH₂CH₂)₂–N(CH₃)H⁺, as shown in Figure 5.

2.4. Model Validation

In this study, the COSMO solvation model was employed in the simulations. To verify the reliability of the established computational model, the simulation results obtained using only the implicit solvation model were compared with those obtained using the water-assisted model, while all other computational conditions remained consistent with those described in Section 2.1 Computational Methods.

Under the water-assisted model, the energy barriers for the reaction between MEA and CO₂ are presented in Figures S1 and S2. As summarized in Table S1, the comparison between the two models shows that the maximum absolute error between the COSMO solvation model and the water-assisted model is 3.19 kJ/mol. In addition, the calculated energy barrier diagrams obtained from both models exhibit similar trends.

Based on these results, it can be concluded that the COSMO solvation model possesses reliability comparable to that of the water-assisted model and can reasonably reproduce the actual reaction process. Therefore, in the subsequent investigations, the COSMO solvation model was continuously adopted to simplify the computational procedure and data visualization.

3. Results and Discussion

3.1. Reaction Analysis of MEA with CO₂

By calculating the energy barriers of the two reaction steps, the energy profile of the reaction process can be clearly determined, and the rate-determining step can be identified. As shown in Figure 6 and Figure 7, the energy barrier of the first step is 32.02 kJ/mol, with an endothermic energy of 7.4 kJ/mol, while the second step exhibits an energy barrier of 1.41 kJ/mol and an exothermic energy of 61.5 kJ/mol. The text continues here.

In the first step, MEA reacts with CO₂ to form HO–C₂H₄–NHCOO⁻ and a free proton (H⁺). This step has the highest energy barrier (32.02 kJ/mol) in the overall reaction, indicating that it is the rate-determining step. The energy barrier of this step mainly originates from the energy required to break the N–H bond. In the second step, the free proton rapidly combines with an MEA molecule to form HO–C₂H₄–NH₃⁺. This step releases 61.5 kJ/mol of energy, representing the largest energy release in the overall reaction process.

Based on the calculated energy barriers, the reaction of MEA with CO₂ is a typical exothermic process [25], indicating that the reaction is thermodynamically favorable and can proceed readily [26].

3.2. Reaction Analysis of DEA with CO₂

By calculating the energy barriers of the two reaction steps, the energy profile of the reaction process can be clearly elucidated and the rate-determining step can be identified. As shown in Figure 8 and Figure 9, the energy barrier of the first step is 47.35 kJ/mol, with an endothermic energy of 14.37 kJ/mol, while the second step exhibits an energy barrier of 14.57 kJ/mol and an exothermic energy of 53.77 kJ/mol.

In the first step, DEA reacts with CO₂ to form (HO–CH₂CH₂)₂–NCOO⁻ and a free proton (H⁺). This step has the highest energy barrier (47.35 kJ/mol) in the overall reaction, indicating that it is the rate-determining step. Similar to the MEA system, the energy barrier mainly arises from the energy required to break the N–H bond. In the second step, the free proton rapidly combines with a DEA molecule to form (HO–CH₂CH₂)₂–NH₂⁺. This step releases 53.77 kJ/mol of energy, representing the largest energy release in the overall reaction.

Based on the calculated energy barriers, the reaction of DEA with CO₂ is thermodynamically favorable and exothermic, indicating that the absorption process can proceed readily.

The reaction pathway of DEA is similar to that of MEA. However, the hydrogen bonding interactions in the products formed from DEA are fewer and weaker than those in the MEA system. In addition, the larger steric hindrance of DEA leads to lower stability of the reaction products. As a result, the products of DEA reacting with CO₂ are more easily desorbed compared to those of MEA.

3.3. Reaction Analysis of TEA with CO₂

By calculating the energy barriers of the reaction steps, the energy evolution of the reaction process can be clearly elucidated, and the rate-determining step can be identified. As shown in Figure 10 and Figure 11, the maximum energy barrier of the first step is 94.17 kJ/mol, accompanied by an exothermic energy of 24.77 kJ/mol. The second step exhibits an energy barrier of 12.49 kJ/mol with an endothermic energy of 0.86 kJ/mol, while the third step has a maximum energy barrier of 7.84 kJ/mol and an endothermic energy of 41.57 kJ/mol.

In the weakly basic TEA solution, CO₂ first reacts with water to form carbonic acid. This step has the highest energy barrier (94.17 kJ/mol) in the overall reaction, indicating that it is the rate-determining step. The energy barrier mainly arises from two contributions: (i) the structural reorganization of the CO₂ molecule, involving redistribution of the π-electron system, which requires energy input; and (ii) the partial cleavage of the O–H bond in the water molecule, leading to the formation of a free proton (H⁺), which also consumes energy.

Subsequently, the released proton combines with a TEA molecule to form (HO–CH₂CH₂)₃–NH⁺. This step releases 41.57 kJ/mol of energy, representing the largest energy release in the overall reaction process.

Based on the calculated energy barriers, although the reaction of TEA with CO₂ is overall exothermic, the hydration of CO₂ involves a significantly high activation barrier. As a result, the reaction proceeds at a relatively slow rate and is less favorable compared to primary and secondary amine systems.

The reaction mechanism of TEA with CO₂ differs significantly from that of MEA and DEA. As a tertiary amine, the nitrogen atom in TEA is bonded to three hydroxyethyl groups and does not contain any N–H bonds. Therefore, TEA cannot directly react with CO₂ to form carbamate species. Instead, it acts as a base, promoting the hydration and subsequent ionization of CO₂ in aqueous solution, leading to the formation of bicarbonate ions. It can thus be concluded that although the reaction of TEA with CO₂ is not kinetically favorable, the resulting products are relatively unstable and can be readily desorbed.

3.4. Reaction Analysis of MDEA with CO₂

By calculating the energy barriers of the reaction steps, the energy evolution of the reaction process can be clearly elucidated, and the rate-determining step can be identified. As shown in Figure 12 and Figure 13, the maximum energy barrier of the first step is 54.53 kJ/mol, accompanied by an exothermic energy of 27.98 kJ/mol. The second step exhibits an energy barrier of 9.5 kJ/mol with an endothermic energy of 0.89 kJ/mol, while the third step has a maximum energy barrier of 1.41 kJ/mol and an exothermic energy of 50.61 kJ/mol.

In the weakly basic MDEA solution, CO₂ first reacts with water to form carbonic acid. This step has the highest energy barrier (54.53 kJ/mol) in the overall reaction, indicating that it is the rate-determining step. The energy barrier mainly arises from two contributions: (i) the distortion of the linear structure of CO₂, leading to redistribution of the π-electron system and associated energy consumption; and (ii) the partial cleavage of the O–H bond in the water molecule, which generates a free proton (H⁺) and requires additional energy input.

Subsequently, the released proton combines with an MDEA molecule to form (HO–CH₂CH₂)₂–N(CH₃)H⁺. This step releases 50.61 kJ/mol of energy, representing the largest energy release in the overall reaction process.

Based on the calculated energy barriers, although the reaction of MDEA with CO₂ is overall exothermic, the relatively high activation barrier associated with CO₂ hydration results in a slow reaction rate, making the process less favorable compared to primary and secondary amines.

The reaction mechanism of MDEA with CO₂ is similar to that of TEA. However, due to its stronger basicity, MDEA exhibits a more pronounced catalytic effect on the hydration of CO₂, resulting in a lower activation energy barrier and making the reaction more favorable compared to TEA. Moreover, since the reaction products are bicarbonate species with relatively low stability, the desorption process can also proceed readily.

4. Discussion

Based on the density functional theory (DFT) investigation of the reaction mechanisms of alkanolamine solutions with CO₂, the main conclusions are summarized as follows:

1. Reaction pathways differ among amine types.

Primary and secondary amines react with CO₂ to release a proton and form a negatively charged carbamate species. The released proton subsequently combines with another amine molecule to form a protonated cation, and together they constitute carbamate. In contrast, tertiary amines react with protons generated from water ionization, promoting the hydration of CO₂ to form bicarbonate species.

2. Reaction energies vary for different amines.

The heat released during CO₂ absorption follows the order: primary amine > secondary amine > tertiary amine. Specifically, the heat release is 54.1 kJ/mol for primary amines and 49.72 kJ/mol for secondary amines. For tertiary amines, MDEA and TEA release 40.77 kJ/mol and 39.4 kJ/mol, respectively. This indicates that primary amines exhibit the highest exothermicity.

3. Reaction products are different.

Primary and secondary amines form carbamate species upon reaction with CO₂, whereas tertiary amines produce bicarbonate species.

4. Activation energy barriers differ significantly.

The highest energy barriers are 32.02 kJ/mol for primary amines and 47.35 kJ/mol for secondary amines. Tertiary amines exhibit much higher barriers, with values of 94.17 kJ/mol for TEA and 54.53 kJ/mol for MDEA. These results indicate that primary amines react most readily with CO₂, while tertiary amines show the lowest reactivity. Among tertiary amines, MDEA reacts more easily with CO₂ than TEA.

5. Desorption energy requirements vary.

The energy required for desorption is highest for primary amines and lowest for tertiary amines.