Carbon Dioxide Absorption Characteristics of Aqueous Alkanolamine Using Nuclear Magnetic Resonance Spectroscopy

Introduction

The Intergovernmental Panel on Climate Change proposes a scenario in which the primary cause of global warming is the presence of greenhouse gases, of which carbon dioxide (CO₂) accounts for a high percentage, 76.7%. Among the world energy sources, fossil fuels contribute most to CO₂ emissions, 56.6%, and this level is expected to persist beyond 2030. Therefore, international regulations for controlling CO₂ emissions have become necessary. CO₂ reductions through carbon capture and sequestration provide a potential countermeasure (Pachauri et al., 2007).

Postcombustion capture technologies mainly treat CO₂ in flue gases generated by the combustion of fossil fuels and air in thermoelectric power plants. Industry capture technologies commonly rely on chemical absorption by aqueous absorbents. Alkanolamines are the most widely commercialized absorbents, and diverse studies have examined their structural characteristics, absorption capacities, gas flux properties, reaction temperature sensitivities, and correlation between CO₂ partial pressure and absorbent concentration (Bonenfant et al., 2003; Yoon and Lee, 2003). Monoethanolamine (MEA) is an excellent absorbent with a high CO₂ absorption rate; however, it has several disadvantages, such as a high cost associated with the need for companion corrosion inhibitors and a high vapor pressure. MEA is strongly corrosive, and the absorption capacity decreases over long-term operation as a result of degradation. Diethanolamine (DEA) and triethanolamine (TEA) are less corrosive than MEA, and their vapor pressures are lower, which minimizes evaporation loss; however, DEA and TEA provide low rates of absorption. To ameliorate these disadvantages, blends of the alkanolamines or promoters were prepared in an effort to improve the reaction rates (Glasscock et al., 1991; Closmann et al., 2009).

Understanding the reaction mechanisms that underlie CO₂ absorption can guide development of new absorbents and more efficient processes by articulating models of the chemical equilibria in mixed absorbents or absorbent–promoter blends. Thus far, studies of species formation in absorbents with absorbed CO₂ have relied on molecular modeling (Ma'mun et al., 2006; Edali et al., 2009). Although researchers have experimentally studied alkanolamine species formation in absorbents (Jakobens et al., 2008; Yang et al., 2009; Barzagli et al., 2009), a few studies have examined the absorption characteristics of primary, secondary, and tertiary alkanolamines. This study measured the CO₂ absorption rates and capacities by aqueous solutions of MEA, DEA, and TEA by using a vapor–liquid equilibrium (VLE) apparatus, and the distributions and absorption characteristics of the species were characterized by ¹H and ¹³C nuclear magnetic resonance (NMR) spectral analysis.

Reaction and Mechanisms

CO₂ absorption in aqueous MEA (primary alkanolamine) or DEA (secondary alkanolamine) solutions has been described from a mechanistic perspective by Caplow (1968) and Danckwerts (1979). Danckwerts proposed that zwitterions formed by a reaction between amines and CO₂ with base removal of the zwitterion protons to form a carbamate. The mechanism of reaction between CO₂ and an amine is also a zwitterion mechanism. The major mechanisms of CO₂ absorption by amine aqueous solutions are presented in Equations (1)–(4). Primary and secondary amines absorb CO₂ to form carbamate or bicarbonate/carbonate and are regenerated as free amines at high temperatures accompanied by CO₂ discharge.

(1)

(2)

(3)

MEA: R=CH₂CH₂OH, R₁=H

DEA: R=CH₂CH₂OH, R₁=CH₂CH₂OH

TEA, a tertiary alkanolamine, does not directly react with CO₂ to form carbamate; rather, it forms bicarbonate through hydrolysis of CO₂. Donaldson and Nguyen (1980) and Singh and Versteeg (2008) studied the kinetics of CO₂ reaction with TEA and reported that TEA acted as a base catalyst. They concluded that carbamate did not form. In this study, ¹³C NMR peak analysis showed that TEA formed bicarbonate.

(4)

TEA: R=CH₂CH₂OH, R₁=CH₂CH₂OH, R₂=CH₂CH₂OH

The reactions of primary, secondary, and tertiary amines suggested that MEA, DEA, and TEA aqueous solutions would form the species shown in Fig. 1 after absorbing CO₂.

OPEN IN VIEWER

FIG. 1.

Molecular structures of the species present in CO₂-loaded MEA, DEA, or TEA aqueous solutions. MEA, monoethanolamine; DEA, diethanolamine; TEA, triethanolamine.

Experimental

MEA (99%), DEA (98.5%), and TEA (99%) used in preparing samples for VLE absorption experiments and NMR analysis were purchased from Samchun Chemical Co. All reagents were used as received. MEA, DEA, and TEA absorbents were used to prepare 30 mass% aqueous solutions in deionized water. The gas used in all experiments was 99.99 vol% CO₂.

A diagram of the VLE apparatus is shown in Fig. 2. The apparatus included a stainless steel CO₂ gas reservoir, a stainless steel reactor for reactions between the absorbents and CO₂, an indicator for indicating the temperature and pressure, and a recorder for storing the CO₂ pressure measurements in real time. The reactor was charged with a batch of absorbent solution and a magnetic stirrer to maximize the contact area between CO₂ and absorbents. The internal volumes were 300.29 cm³ for the reservoir and 322.56 cm³ for the reactor. The volume difference was calibrated by using 99.999 vol% N₂. CO₂ was injected into the reservoir after preheating to 313K by using a constant-temperature water bath. The reactor was maintained at 333K to measure the absorption capacities as a function of temperature. Before injecting CO₂, residual gases in the reactor were removed by using a vacuum pump. When the temperature of the gas reservoir and reactor reached the experimental temperature, a CO₂ injection valve was opened to inject CO₂ into the reactor. The pressure of CO₂ in the reactor was shown to decrease over time. The pressure in the reactor was measured every five seconds. When the temperature and pressure stabilized, a state of equilibrium was reached, and the pressure in that state was measured to calculate the CO₂ equilibrium partial pressure. The pressure of the injected gas and the pressure of the gas after equilibrium were measured to calculate the moles of CO₂ absorbed by the absorbent. Since CO₂ reacted with the absorbents only in the closed reactor, the absorption rate could be calculated from the pressure difference. After the initial CO₂ injection into the reactor, CO₂ was injected again after equilibrium had been reached to establish a new equilibrium state.

OPEN IN VIEWER

FIG. 2.

Schematic diagram of experimental vapor–liquid equilibrium apparatus: 1, CO₂ gas; 2, gas reservoir; 3, reactor; 4, magnetic stirrer; 5, vacuum pump; 6, pressure measuring instrument; 7, recorder; 8, pressure transmitter; 9, temperature indicator.

The CO₂ loading capacity of an absorbent was calculated by using the ideal gas equations. The experiments in this study were conducted at low pressures (20–800 kPa), and the ideal gas equation can be used to calculate the CO₂ loading capacity.

The CO₂ loading capacity in the reactor can be calculated by using Equations (5–8): \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} n_{{S_{{\rm CO}_2}}} = \frac {(P_{Si} - P_{St}) \times V_S} {RT_S} \tag{5} \end{align*} \end{document}

The amount of CO₂ \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$(n_{{S_{{\rm CO}_2}}})$$ \end{document} that reacted at the reservoir can be obtained by multiplying the terminal pressure (P_St) difference after an injection to the reservoir at the initial pressure (P_Si) of the reservoir with the volume of the reservoir (V_S) followed by the division of universal gas constant (R) and the temperature of the reservoir (T_S). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} n_{{R_{{\rm CO}_2}}} = \frac {(P_{Ri} - P_{Rt}) \times V_R} {RT_R} \tag{6} \end{align*} \end{document}

The amount of CO₂ \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$n_{{R_{{\rm CO}_2}}}$$ \end{document} that reacted to a reactor can be also obtained in the same way. The pressure before the injection of a reactor is (P_Ri), and the remaining pressure after the injection of a reactor is (P_Rt). The volume of a reactor is (V_R), and the temperature of a reactor is (T_R).

From Equations (5) and (6), the amount of CO₂ absorbed by reaction can be calculated by \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} n_{{{\rm absorbed}_{{\rm CO}_2}}} = n_{{S_{{\rm CO}_2}}} - n_{{R_{{\rm CO}_2}}} \tag{7} \end{align*} \end{document}

The CO₂ loading capacity can then be calculated by \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*} {\rm CO}_2 \ {\rm loading} = \frac {n_{{{\rm absorbed}_{{\rm CO}_2}}}} {n_{a \min e}} \tag{8} \end{align*} \end{document}

The reactor was designed so that the internal pressure, measured during the absorption reactions, could be stored in real time. Each sample was sealed and analyzed by NMR.

¹H NMR and ¹³C NMR were measured by using a model AVANCE 500 MHz from Bruker Co., and the solvent used in the analytical experiments was deuterium oxide (D₂O) 99.99% from Sigma-Aldrich Co. The NMR analytical conditions of each sample are provided in Table 1.

Results and Discussion

Speciation

The hydroxyl groups (-OH) in MEA and DEA polarize the molecular electronic distribution by pulling electron density away from adjacent atoms. Since -OH groups reduce the electron density of adjacent protons, thereby reducing the local response to diamagnetic shielding effects, the peaks of these protons shift down-field. The peak positions in the ¹H NMR and ¹³C NMR spectra are shown in Figs. 3–5 and 6–8, respectively. The formation of carbamate and bicarbonate moieties was indicated by the evolution of low-field peaks at 165–160 ppm in the ¹³C MNR spectrum. In this study, carbamate and bicarbonate were formed in both MEA and DEA, as shown in Figs. 6 and 7. In the case of TEA, bicarbonate formed selectively, and carbamate did not form (Fig. 8). These results were consistent with the results of a study performed by Singh and Versteeg (2008). The CO₂ loading capacities of MEA, DEA, and TEA aqueous solutions and the resultant chemical shifts are given in Tables 2–4. In the NMR samples, mixed DOH molecules formed due to proton exchange with the D₂O solvent. The protons of amine groups and OH^– could not be separately distinguished in ¹H NMR, and the missing protons appeared as new DOH peaks. In this experiment, such protons appeared near δ=4.8. The composition of the MEA and DEA solutions differed significantly at the start of the reaction. DEA formed carbamate and bicarbonate simultaneously at the start of the CO₂ absorption process, whereas MEA formed bicarbonate only after certain levels of carbamate had accumulated. The VLE absorption results showed that CO₂ loading in MEA and DEA could not theoretically exceed 0.5 moles for stoichiometric reasons, and additional absorption was only possible on formation of free amines through the hydrolysis of carbamate. Sartori and Savage (1983) explained this observation in terms of the stability of the carbamate moiety. The reactions that formed bicarbonate in TEA were initiated by reactions between water and CO₂, so the maximum equilibrium loading exceeded that in the primary or secondary amines. However, the reaction rate was quite low.

OPEN IN VIEWER

FIG. 3.

¹H-NMR spectra during CO₂ absorption by the MEA-H₂O-CO₂ system at 333K. Proton resonances at different CO₂ absorption levels in eight 30 mass% MEA solutions as a function of CO₂ loading. NMR, nuclear magnetic resonance.

OPEN IN VIEWER

FIG. 4.

¹H-NMR spectra during CO₂ absorption by the DEA-H₂O-CO₂ system at 333K. Proton resonances at different CO₂ absorption levels in eight 30 mass% DEA solutions as a function of CO₂ loading.

OPEN IN VIEWER

FIG. 5.

¹H-NMR spectra during CO₂ absorption by the TEA-H₂O-CO₂ system at 333K. Proton resonances at different CO₂ absorption levels in eight 30 mass% TEA solutions as a function of CO₂ loading.

OPEN IN VIEWER

FIG. 6.

¹³C-NMR spectra during CO₂ absorption by the MEA-H₂O-CO₂ system at 333K. Carbon resonances at different CO₂ absorption levels in eight 30 mass% MEA solutions as a function of CO₂ loading.

OPEN IN VIEWER

FIG. 7.

¹³C-NMR spectra during CO₂ absorption by the DEA-H₂O-CO₂ system at 333K. Carbon resonances at different CO₂ absorption levels in eight 30 mass% DEA solutions as a function of CO₂ loading.

OPEN IN VIEWER

FIG. 8.

¹³C-NMR spectra during CO₂ absorption by the TEA-H₂O-CO₂ system at 333K. Carbon resonances at different CO₂ absorption levels in eight 30 mass% TEA solutions as a function of CO₂ loading.

OPEN IN VIEWER

Table 2.

¹ H Nuclear Magnetic Resonance Spectra of Chemical Shift and Peak Areas of Monoethanolamine at 333K

	Loading	1 H-NMR species in solution, δ (ppm)
Amine	Mole CO2/mole absorbent	Free MEA (CH2-NH2)	Carbamate (CH2-NHCOO−)	Carbamate (CH2-OH−)	Free MEA (CH2-OH)
MEA	0.084	3.588 (0.99)	3.562 (0.08)	3.119 (0.07)	2.733 (1.00)
	0.170	3.616 (1.01)	3.565 (0.15)	3.122 (0.15)	2.785 (1.00)
	0.250	3.639 (1.01)	3.561 (0.23)	3.118 (0.23)	2.832 (1.00)
	0.333	3.666 (100)	3.559 (0.32)	3.116 (0.32)	2.887 (1.00)
	0.416	3.704 (1.00)	3.565 (0.38)	3.122 (0.38)	2.953 (1.00)
	0.496	3.728 (1.00)	3.563 (0.40)	3.119 (0.41)	3.002 (1.00)
	0.557	3.741 (1.00)	3.564 (0.36)	3.121 (0.37)	3.026 (1.00)
	0.601	3.748 (0.99)	3.562 (0.36)	3.119 (0.36)	3.041 (1.00)

The values inside parenthesis are integral calculus values.

MEA, monoethanolamine.

OPEN IN VIEWER

Table 3.

¹ H Nuclear Magnetic Resonance Spectra of Chemical Shift and Peak Areas of Diethanolamine at 333K

	Loading	1 H-NMR species in solution, δ (ppm)
Amine	Mole CO2/mole absorbent	Free DEA (CH2-NCOO−)	Carbamate (CH2-NH)	Carbamate (CH2-OH)	Free DEA (CH2-OH−)
DEA	0.147	3.696 (0.99)	3.654 (0.05)	3.363 (0.04)	2.808 (1.00)
	0.293	3.725 (0.99)	3.653 (0.08)	3.362 (0.07)	2.890 (1.00)
	0.427	3.751 (1.00)	3.654 (0.08)	3.363 (0.08)	2.963 (1.00)
	0.547	3.774 (1.00)	3.655 (0.08)	3.364 (0.08)	3.023 (1.00)
	0.642	3.786 (1.00)	3.654 (0.08)	3.363 (0.07)	3.059 (1.00)
	0.707	3.796 (1.00)	3.652 (0.08)	3.361 (0.07)	3.090 (1.00)
	0.750	3.802 (1.00)	3.652 (0.07)	3.361 (0.06)	3.106 (1.00)
	0.785	3.806 (1.00)	3.654 (0.07)	3.363 (0.06)	3.114 (1.00)

The values inside parenthesis are integral calculus values.

DEA, diethanolamine.

OPEN IN VIEWER

Table 4.

¹ H Nuclear Magnetic Resonance Spectra of Chemical Shift and Peak Areas of Triethanolamine at 333 K

	Loading	1H-NMR species in solution, δ (ppm)
Amine	Mole CO2/mole absorbent	Free TEA (CH2-N)	Carbamate (CH2-NHCOO−)	Carbamate (CH2-OH−)	Free TEA (CH2-OH)
TEA	0.179	3.708	x	x	2.838
	0.330	3.759	x	x	2.958
	0.450	3.788	x	x	3.029
	0.543	3.808	x	x	3.083
	0.618	3.809	x	x	3.086
	0.668	3.823	x	x	3.133
	0.693	3.831	x	x	3.157
	0.726	3.828	x	x	3.136

TEA, triethanolamine.

During the conversion of bicarbonate ( \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^-$$ \end{document} ) to carbonate ( \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm CO}_3^{2 -}$$ \end{document} ), the ¹³C NMR chemical shifts corresponding to aqueous amines merged into a single peak at δ=160–164 ppm due to kinetic averaging. The relative intensities of the bicarbonate and carbonate peaks indicated the position of the equilibration (Mani et al., 2006). The relative quantities of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^-$$ \end{document} and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm CO}_3^{2 -}$$ \end{document} were indicated by the position of merged ¹³C NMR peak, based on analysis of the chemical shifts of KHCO₃ (δ=160.82) and K₂CO₃ (δ=168.24). According to the method described by Mani, the relative quantities of the components could be determined from the shift in the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^- / {\rm CO}_3^{2 -}$$ \end{document} peak, shown in Figs. 6 and 7, which increased with CO₂ loading. That is, as the CO₂ loading increased, a higher rate of the reverse reaction, corresponding to the conversion of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^{-}$$ \end{document} to \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm CO}_3^{2 -}$$ \end{document} , resulted in a gradual shift toward higher field strengths. Increasing the CO₂ loading prevented the conversion of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^-$$ \end{document} to \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm CO}_3^{2 -}$$ \end{document} , as indicated by the absence of a chemical shift in the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^- / {\rm CO}_3^{2 -}$$ \end{document} peak (Fig. 8). The initial absorption of CO₂ proceeded through the carbonate.

Vapor–liquid equilibrium

The extent of CO₂ absorption at equilibrium by 30 mass % MEA, DEA, or TEA aqueous solutions is shown in Fig. 9. The graph shows the specific partial pressure of CO₂ ( \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $$P_{{\rm CO}2}{}^*$$ \end{document} ) as a function of CO₂ loading in the absorbents. CO₂ loading is the moles CO₂ absorbed by 1 mole amine. At a given CO₂ loading, low specific partial pressures of CO₂ yield high absorption capacities. The slope of the absorption capacity of MEA rapidly increased from a CO₂ loading of 0.4 mole, and the absorption capacity reached a maximum at 0.6 mole, after which point, further absorption was not observed. Two moles of amine in MEA or DEA reacted with 1 mole of CO₂, and the equilibrium loading reached a limit at 0.5 moles (Sartori and Savage, 1983). Absorption of 0.5 mole CO₂ or higher resulted in the generation of free amine groups due to the hydrolysis of carbamate. The free amine groups also absorbed CO₂. The equilibrium CO₂ loading of DEA was higher, because the carbamate stability was lower than in MEA. Since TEA acts as a base catalyst in the reaction between water and CO₂, TEA showed an equilibrium loading similar to that of DEA under saturated conditions. Changes in pressure in the reactor over time after CO₂ was initially injected into the absorbent are shown in Fig. 10. These changes indicated the relative absorption reaction rates of individual absorbents. The order of the CO₂ absorption rates was MEA>DEA>TEA. In this study, the ¹H and ¹³C NMR data for MEA, DEA, and TEA were used to calculate the CO₂ absorption mechanism. ¹H and ¹³C NMR results provided information about the species distribution and reaction mechanism. The species distribution in the absorbents is key to the design of reliable CO₂ capture processes. The species distributions and analysis methods used in this study provide guidelines that can be used to develop new absorbents.

OPEN IN VIEWER

FIG. 9.

Equilibrium solubility of CO₂ in aqueous 30 mass% MEA, DEA, and TEA solutions at 333K.

OPEN IN VIEWER

FIG. 10.

Total pressure of the gases in aqueous MEA, DEA, and TEA 30 mass% at 333K.

Conclusions

The distribution of species formed on CO₂ absorption by primary, secondary, or tertiary amine aqueous solutions was identified by analysis of ¹H NMR and ¹³C NMR spectra. The \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^- / {\rm CO}_3^{2 -}$$ \end{document} peaks did not appear in the CO₂-absorbed MEA ¹³C NMR spectrum at low CO₂ loading but were observed at high CO₂ loading (loading=0.416 moles). These peaks appeared at δ=160.78–162.55 ppm, depending on the CO₂ loading. \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^- / {\rm CO}_3^{2 -}$$ \end{document} peaks appeared in the CO₂-absorbed DEA ¹³C NMR spectrum at δ=160.50–162.70 ppm. Although the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^- / {\rm CO}_3^{2 -}$$ \end{document} peaks appeared at low levels of TEA CO₂ loading, as with DEA, the peak intensities were very small at δ=160.27–160.37 ppm. Therefore, as the CO₂ loading increased, the fraction of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^-$$ \end{document} exceeded that of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm CO}_3^{2 -}$$ \end{document} , and the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} $${\rm HCO}_3^- / {\rm CO}_3^{2 -}$$ \end{document} peaks were shifted up-field. ¹H NMR analysis showed that at CO₂ absorption of >0.5 moles by MEA or DEA, the predicted stoichiometric point of saturation, the reactions shifted from carbamate formation to bicarbonate and carbonate formation, thus reflecting the maximum level of absorption by the individual amines.