Better Choice of Tertiary Alkanolamines for Postcombustion CO2

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Better Choice of Tertiary Alkanolamines for Postcombustion CO2 Capture: Structure with Linear Alkanol Chain Instead of Branched Sen Liu, Hao Ling, Hongxia Gao,* Paitoon Tontiwachwuthikul, and Zhiwu Liang* Joint International Center for CO2 Capture and Storage (iCCS), Hunan Provincial Key Laboratory for Cost-Effective Utilization of Fossil Fuel Aimed at Reducing CO2 Emissions, College of Chemistry and Chemical Engineering, Hunan University, Changsha 410082, People’s Republic of China

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S Supporting Information *

ABSTRACT: To give better guidance of tertiary alkanolamine selection with proper structures for postcombustion CO2 capture, effects of the alkanol chain structure and intramolecular hydrogen bond were investigated. First, water solubilities of 2.5 M aqueous solutions were experimentally investigated for 11 tertiary alkanolamines, and data of boiling point and vapor pressure for pure chemicals were collected also. The physical property results showed a biphasic phenomenon were observed in 2.5 M fresh 1diethylamino-2-propanol (1DEA2P) solution and CO2-loaded 4diethylamino-2-butanol (4DEA-2B) solution, and 1-dimethylamino-2-propanol (1DMA2P) and 1DEA2P have a dramatically relative lower boiling point and a higher vapor pressure, indicating the existence of an intramolecular hydrogen bond, especially in amines with branched alkanol chains. In addition, the pH values for fresh and equilibrium CO2-loaded amine solutions (2.5 M, 313 K), equilibrium CO2 solubility (2.5 M, 313 K, and 15 kPa), first-order reaction rate constant (0.1−0.4 M and 313 K), and dissociation constant (293−333 K) were measured, and the results showed amines with linear alkanol chains can have relatively strong basicity, absorption capacity, and a reaction rate with CO2 owing to weaker influences of the intramolecular hydrogen bond. Finally, the CO2 capture performance was comprehensively compared by ΔrGm and ΔrHm screening methods as well as the fast solvent screening method with cyclic capacity, and the results further confirmed that amines with linear alkanol chains can present better CO2 capture performance, like 3-dimethylamino-1-propanol (3DMA1P) and 3-diethylamino-1-propanol (3DEA1P). It also suggests that the selection or design of tertiary alkanolamines should with the linear alkanol chain instead of branched to have better CO2 capture performance for industrial applications. structures. Chowdhury et al.11 investigated 24 tertiary alkanolamines with the absorption rate, desorption rate, and cyclic capacity in a period of 60 min using 30 wt % fresh aqueous solution. The study found out that 1-dimethylamino2-propanol (1DMA2P), 2-diethylaminoethanol (DEEA), 3diethylamino-1,2-propanediol (DEA-1,2-PD), 3-diethylamino1-propanol (3DEA1P), 1-methyl-2-piperidineethanol (1 M2PPE), 1-(2-hydroxyethyl)pyrrolidine (1-(2HE)PRLD), and 1-(2-hydroxyethyl) piperidine (1-(2HE)PP) were the preferred amines with the faster absorption rate and lower heats of reactions. Rayer et al.12 investigated the influences of the steric hindrance, −OH group number, and alkyl chain length on the CO2 capture performance of various amines based on the dissociation constant (pKa) and enthalpy. The results revealed that DEEA and 3-dimethylamino-1-propanol (3DMA1P) with

1. INTRODUCTION To mitigate global warming, carbon dioxide (CO2) capture has been considered as one of the most efficient technology by reducing CO2 emissions from large fossil-powered plants.1,2 Amine solution scrubbing has been known as the most developed and doable technology for postcombustion CO2 capture (PCC) owing to its fast absorption rate and ease to scale-up advantages.3 It is, nevertheless, restricted by vast costs for investment and operation, and about 70% of the operating cost is consumed as heat energy in order to recover amines.4,5 Under the circumstance, tertiary amines with low regeneration of heat energy have been attracted much attentions due to its absorption product of (bi)carbonate.6,7 However, its disadvantage is also notable for the slow absorption rate with CO2, which is directly related to basicity of tertiary amines.8,9 Hence, absorption and desorption rates are required to be coordinated with the proper basicity and molecular structure.10 Toward efficient tertiary amine design and selection for CO2 capture, some researchers investigated the CO2 capture performance of tertiary amines with various molecular © XXXX American Chemical Society

Received: Revised: Accepted: Published: A

April 24, 2019 June 25, 2019 July 29, 2019 July 29, 2019 DOI: 10.1021/acs.iecr.9b02244 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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Industrial & Engineering Chemistry Research

Figure 1. Abbreviations and chemical structures of tested tertiary alkanolamines.

and the feasible heterocyclic structures were proposed. In addition, to experimentally demonstrate existence of the intramolecular hydrogen bond, the physical property parameters such as boiling point, vapor pressure, and water solubility were investigated. Additionally, to understand the exact influences of the intramolecular hydrogen bond on CO2 absorption and desorption performances of tested tertiary alkanolamines, experiments were performed for the equilibrium CO2 solubility (2.5 M, 313 K, and 15 kPa), pH values (313 K, fresh and equilibrium CO2-loaded solutions), firstorder reaction rate constant (k0, 0.1−0.4 M and 313 K), acid dissociation constant (ln Ka, 293−323 K), and fast solvent screening (absorption at 313 K and 15 kPa, desorption at 353 K, and with nitrogen stripping). Finally, the influences of the intramolecular hydrogen bond on CO2 capture performance were comprehensively analyzed and compared with experimental results for all tested tertiary alkanolamines.

linear alkanol chains were the preferred tertiary amines. Xiao et al.13 studied the structure−activity characteristics of 10 tertiary amines for the commercial application of PCC based on the second-order reaction rate constant (k2), equilibrium solubility, pKa value, and absorption heat. The work concluded that the electron-donating effect, steric hindrance effect, and side carbon chain (branched alkanol chain) may promote the activity of tertiary amines, but the heterocyclic structure (fiveor six-membered ring caused by the intramolecular hydrogen bond) and the addition of the hydroxyl group can reduce the values of equilibrium CO2 solubility, k2, and pKa. Also, it suggested that the hydroxy group can have its proper position to achieve better CO2 capture performance. Based on these researches, it can be concluded that the steric hindrance and electron-donating effect of alkyl group addition on the N atom can be in favor of CO2 capture, while the addition of the hydroxyl group can decrease the basicity and reaction rate of amines. But meanwhile, differentiated conclusions can be observed, mainly focused on the preferred molecular structure and proper position of the hydroxyl group (like 1DMA2P and 3DMA1P). In addition, Przesl̷awska et al.14 also proved that 3DMA1P has a heterocyclic structure of sixmembered ring resulted by the intramolecular hydrogen bond, and Xiao et al.13 also claimed that the internal hydrogen bond can significantly affect the CO2 capture performance of tertiary amines. However, there is no systematic research about the effect of the alkanol chain structure and intramolecular hydrogen bond on CO2 capture performance of tertiary alkanolamines. Therefore, further researches are required to address these problems, including (i) the existence and accurate effect of the intramolecular hydrogen bond on physicochemical properties of tertiary alkanolamines and (ii) the exact influences of the alkanol chain structure connected to the N atom (linear or branched) and the intramolecular hydrogen bond on CO2 capture performance of tertiary alkanolamines. To figure out these two issues, 11 tertiary alkanolamines were selected as objects of this study, including 2(dimethylamino)ethanol (DMEA), 1DMA2P, 3DMA1P, 3(dimethylamino)-1,2-propanediol (DMA-1,2-PD), 2-(diethylamino) ethanol (DEEA), 1-diethylamino-2-propanol (1DEA2P), 3DEA1P, 3-(diethylamino)-1,2-propanediol (DEA-1,2-PD), 4-diethylamino-2-butanol (4DEA-2B), 2-(dimethylamino)-2-methyl-1-propanol (DMA-2 M-1P), and Nmethyldiethanolamine (MDEA), as presented in Figure 1. Selection of tertiary alkanolamines instead of alkylamines in this work was mainly considering the effect of volatility and environmental pollutions. First, the existence possibilities of the intramolecular hydrogen bond were theoretically analyzed,

2. THEORY AND INFERENCE 2.1. Reaction Mechanism. Reacting CO2 with tertiary amines in aqueous solution follows the base-catalyzed hydration mechanism, which is originally proposed by Donaldson and Nguyen,15 and the overall reaction can be presented as eq 1. The R1R2R3N is the free tertiary amine, and the R1R2R3NH+ is the protonated amine. R1R 2R3N + CO2 + H 2O ↔ R1R 2R3NH+ + HCO−3

(1)

8

According to the latest research work, there are three parallel reactions occurred when CO2 absorption is initiated into aqueous tertiary amine solution, as follows: Reaction of H2O with CO2: CO2 + H 2O ↔ HCO−3 + H+

(2)

Reaction of molecular pair of R1R2R3N and H2O with CO2:16

Reaction of hydroxyl ion (OH−) with CO2: R1R 2R3N + H 2O ↔ R1R 2R3NH+ + OH− −

OH + CO2 ↔

HCO−3

(4) (5)

The rate contribution of each reaction to the overall reaction rate is dependent on basicity (pKa) of tertiary amine and the B

DOI: 10.1021/acs.iecr.9b02244 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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Industrial & Engineering Chemistry Research ion strength of its (bi)carbonate, and with stronger alkalinity, it will perform more likely as a hydroxyl ion supplier and presenting a faster reaction rate for CO2 absorption.8,17 Hence, the basicity is the key parameter for tertiary amines, and its water solubility (hydration and dissociation) can also be of significance for CO2 capture performance. 2.2. Intramolecular Hydrogen Bond and Heterocyclic Structure. Normally, most of the hydrogen bond existed in aqueous amine solution was intermolecular hydrogen bonding. According to eqs 3 and 4, the intermolecular hydrogen bonding is in favor of hydrolysis of amines and water solubility of amine in solution. However, tertiary alkanolamines have at least one amino group and a hydroxyl group, and then the special structure provides possibilities for the existence of the intramolecular hydrogen bond and formation of the heterocyclic structure. According to the reference,18 five- and sixmembered rings can be easier to be formed with higher reactivity. Przesl̷awska et al.14 also proved that 3DMA1P has a heterocyclic structure of a six-membered ring resulted by the intramolecular hydrogen bond. Hence, it can be inferred that tertiary alkanolamines presented in Figure 1 could all have intramolecular hydrogen bonds due to their similar molecular structures. Then, the free tertiary alkanolamines with intramolecular hydrogen bonds and heterocyclic structures were proposed and shown in Figure 2. It can be seen that MDEA, DMEA, DEEA,

Figure 3. Protonated tertiary alkanolamines with heterocyclic structures driven by intramolecular hydrogen bonds.

According to the proposed heterocyclic structures, the tertiary alkanolamines with branched alkanol chains could have higher existence possibilities and stronger forces of the intramolecular hydrogen bond due to the branched alkanol chain may create a natural and favorable angle for the hydrogen bond of N···H−O (for free amines) or O···H+−N (for protonated amines). Second, the intramolecular hydrogen bond can change the physical and chemical properties of tertiary alkanolamines. On the one hand, the intramolecular hydrogen bond can weaken the intermolecular hydrogen bonding between amine molecules or amine molecules with water in solution, and it can lead to a lower boiling point, a higher vapor pressure, and poor water solubility in solution.19 On the other hand, the intramolecular hydrogen bond can disperse the charge density of the N atoms in molecules20 and then decrease the forces and opportunities of amines combining with water, resulting in weaker alkalinity in solution.10,21 Therefore, the existence of heterocyclic structures caused by the intramolecular hydrogen bond can result in poor CO2 absorption performance, leading to the lower equilibrium CO2 solubility and reaction rate. In addition, R1R2R3NH+ with the heterocyclic structure also needs more energy (not only for breaking down the ionic bond of N−H+ but also overcoming the intramolecular hydrogen bond) to release the proton for amine recovery, which may lead to a slower desorption rate for CO2 stripping. 2.3. ΔrGm and ΔrHm. The molar Gibbs free energy change (ΔrGm) and molar reaction enthalpy (ΔrHm) for protontransfer processes of amines or protonated amines were calculated by the van’t Hoff and Gibbs−Helmholtz equations,12,22 which are described as follows:

Figure 2. Tertiary alkanolamines with heterocyclic structures driven by intramolecular hydrogen bonds.

1DMA2P, 1DEA2P, and DMA-2 M-1P may form a fivemembered ring structure in molecules, while 3DMA1P, 3DEA1P, and 4DEA-2B may be more stable with a sixmembered ring in molecules. Both kinds of heterocyclic structures may be formed in DMA-1,2-PD and DEA-1,2-PD molecules due to two hydroxyl groups that were attached on the alkanol chain. In addition, when CO2 absorption is initiated into aqueous tertiary alkanolamine solutions, free amines can be transformed into the R1R2R3NH+. Also, heterocyclic structures of five- and six-membered rings may be also existed in protonated amines, as shown in Figure 3. Based on the proposed heterocyclic structure, the existence possibilities and influences of the intramolecular hydrogen bond on physicochemical properties of tertiary alkanolamines for CO2 capture can be further analyzed. First, it can be observed that DMEA, DEEA, 3DMA1P, and 1DMA2P have linear alkanol chains connected to the N atoms; 1DMA2P, 1DEA2P, DMA-2 M-1P, and 4DEA-2B have branched alkanol chains; and DMA-1,2-PD and DEA-1,2-PD have an extra hydroxyl group on the linear alkanol chain.

Δr Gm = −RT ln K1

Δr Hm = − R

dln K 2 d(1/T )

(6)

(7)

where R is the molar gas constant with a value of 8.314 J· mol−1·K−1, K1 (Ka) and K2 (1/Ka) are the equilibrium constants of forward and reverse reactions for tertiary amines combining the proton. Based on the reaction mechanism and Bro̷ nsted relationships,23,24 the CO2 absorption and desorption rates of tertiary amines can be directly related to alkalinity and the two parameters (K1 and K2), which can, therefore, be applied for evaluating CO2 capture performance of amines in terms of the reaction rate and energy efficiency.12 The lower value of ΔrGm means more energy release for tertiary amine C

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Industrial & Engineering Chemistry Research proton attraction, which can endow stronger basicity and is favorable for CO2 absorption. The lower ΔrHm value means less energy requirement for protonated amine releasing the proton, which is good for tertiary amine recovery and CO2 desorption.

3. CHEMICALS AND EXPERIMENTS 3.1. Chemicals. The information of CAS numbers, purities, and manufacturers for all the used chemicals are presented in Table 1. The 4DEA-2B was experimentally synthetized as Table 1. CAS Number, Purity, and Manufacturer of 11 Tertiary Alkanolamines amines

CAS number

purity (wt %)

manufacturer

MDEA DMEA 1DMA2P 3DMA1P DMA-1,2-PD DEEA 1DEA2P 3DEA1P DEA-1,2-PD DMA-2M-1P DEAB

105-59-9 108-01-0 108-16-7 3179-63-3 623-57-4 100-37-8 4402-32-8 622-93-5 621-56-7 7005-47-2 5467-48-1

99 99 98 99 98 99 >98 >95 97 >97 94

Aladdin reagenta Aladdin reagenta Aladdin reagenta Aladdin reagenta TCI reagenta Aladdin reagenta TCI reagenta Aladdin reagenta Alfa Aesar reagenta TCI reagenta synthesized

Figure 4. Observation of biphasic phenomenon. The left one is 2.5 M fresh aqueous 1DEA2P solution; the right one is 2.5 M aqueous CO2loaded 4DEA-2B solution.

molar concentration of amines ([R1R2NH]), as shown in eq 8,23 and the calculations were elaborated in detail in the Supporting Information. The Ka values of MEA and MDEA were compared with reference data in good agreement as shown in Figure S1, and the repeated experiments were performed for measuring Ka values of MDEA, DMEA, and DEEA with good reliability also, as shown in Table S1.

a

Aladdin reagent, Aladdin Industrial Corporation, Shanghai, China. TCI reagent, Tokyo Chemical Industry Co., Ltd. cAlfa Aesar reagent, Thermo Fisher Scientific-CN, Shanghai, China. b

described in our previous work13 by following the method of Tontiwachwuthikul et al.25 The CO2 and nitrogen (N2) gases with purity of 99.9 vol % were provided by Changsha Rizhen Gas Co. Ltd., China. The aqueous amine and CO2 solutions were prepared by deionized water produced by a reverse osmosis ultra-pure water equipment (model: TS-RO-10 L/H, ≤0.1 μs/cm, Taoshi Water Equipment Engineering Co. Ltd.). 3.2. Water Solubility and Biphasic Phenomenon Observation. The aqueous 1DEA2P and 4DEA-2B solutions with a concentration of 2.5 M were prepared with deionized water. The two phases can be observed for fresh aqueous 1DEA2P solution but not for fresh aqueous 4DEA-2B solution. Then, the pure CO2 gas was bubbled into two solutions, and a biphasic phenomenon gradually disappeared with CO 2 absorption into 1DEA2P solution (at 313 K and CO2 loading of 0.7 mol/mol), while two phases appeared in 4DEA-2B solution with increasing CO2 loading (at 313 K and 0.12 mol/ mol), as displayed in Figure 4. No phase change phenomenon was observed in fresh or CO2-loaded solutions for the other 8 tested tertiary alkanolamines. 3.3. Determination of pH and Dissociation Constant. The pH values and acid dissociation constants of tested alkanolamines in aqueous solutions are determined as our previous work.23 A thermostat bath (model HX020, ±0.05 °C, Hannuo Instruments, Shanghai, China) was used to keep the temperature of the amine solution constant. Then, a pH meter (accuracy of ±0.01, model E-201-D, INESA Scientific Instruments, China) was first calibrated by buffer solutions with standard pH values (6.86 and 9.18 at 298 K). Afterwards, it was used to measure the pH value of fresh amine solutions with a molar concentration of 0.05 M. The Ka value can be calculated based on the parameters of the molar concentration of the proton ([H+]), water dissociation constant (Kw), and

+ ji [R R NH][H ] zy K a = [H+]jjj 1 2 − 1zzz j z Kw (8) k { 3.4. Solvent Screening and Cyclic Capacity. The solvent screening process is just the same as in our former works.26 An absorption process was first performed at 313 K, and a volume of 0.3 L amine solution was bubbled by mixture gases of CO2 and N2, with a partial pressure of 15 kPa, and a total flow rate of 1 L/min. The absorption process was terminated when the reaction was close to the equilibrium (in which the outlet gas CO2 concentration of outlet gas was more than 90% of that of the inlet gas). Then, the desorption process was carried out with rich CO2-loaded amine solutions produced by the absorption process, the temperature was constant at 353 K, and a stream of N2 gas was bubbled into the solution with a flow rate of 0.85 L/min for CO2 stripping. Desorption processes were stopped when the outlet gas CO2 concentration of the outlet gas was close to that of the inlet gas (with a CO2 concentration of less than 1 vol %). The detailed method is shown in the Supporting Information, while the cyclic capacity can be calculated by the fixed time period of CO2-loading curves with times for absorption and desorption processes. The definition of cyclic capacity and its calculations can also be seen in the Supporting Information. 3.5. Measurement of Equilibrium CO2 Solubility. The apparatus and procedure for measuring the equilibrium CO2 solubility of amine solution are the same as our previous work.27 The amine solution was bubbled with mixed gases of CO2 and N2, and the CO2 partial pressure was fixed by gas rationing with mass flow controllers (model D07, ±1.5%

D

DOI: 10.1021/acs.iecr.9b02244 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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Industrial & Engineering Chemistry Research accuracy, Seven Star, China), while the reaction temperature was controlled by a water bath (±0.05 °C). When the CO2 absorption of amine solutions reaches the equilibrium, liquid samples were taken to determine the equilibrium CO2 solubility by a Chittick apparatus as our previous work.27 The experimental equilibrium CO2 solubility data of MDEA were compared and in good agreement with the reference data, as shown in Figure S4. 3.6. Determination of k0. The observed first-order reaction rate constant of CO2 absorption into amine solution (k0, s−1) was determined using a stopped-flow apparatus (SF61DX, Hi-Tech Scientific, Ltd. (U.K.)), and the experimental procedure was similar to our previous work.6,8,28 Two seal syringes were filled with aqueous amine solution and aqueous CO2 solution separately, the two solutions with a volume of 60 μL was synchronously injected into a cell, and then the solution will be stopped and instantaneously mixed. The solutions will be fully mixed with a time of less than 0.002 s where the mass transfer can be ignored in solutions. Then, the conductivity of the mixed solution in the unit was tested and plotted as time. Finally, the k0 was correlated by the timeconductivity curve with the equation of Y = − A × exp ( − k0 × t) + C. The reacted solutions were heated, and the temperature was maintained by a circulating water bath in a sandwich glass, with an accuracy of ±0.1 K. It should be specially noted that the amine-to-CO2 molar ratio in mixed solution was fixed at about 20 to make sure the same CO2 loading of reacted solutions.8 The aqueous CO2 solution is prepared by water saturation with pure CO2 gas bubbling, and the molar concentration of CO2 can be known with reference data29 at equilibrium or dilution with water. Then, the amineto-CO2 molar ratio is fixed by certain molar concentrations of the two solutions.

Figure 5. Comparing physical properties such as boiling point (box for experiment and circle open for prediction at 760 mmHg) and vapor pressure (tilted square open for experiment at 20 °C and triangle tilted open for prediction at 25 °C) for 11 tertiary alkanolamines.

3DMA1P are isomers, DMEA has a lower molecular weight than that of the 1DMA2P, and DMA-2 M-1P has a higher molecular weight than that of 3DMA1P. However, with similar structures and molecular weights for the four amines, the 1DMA2P has dramatically lower boiling point, and then it can be inferred that the presence of the unusual 1DMA2P boiling point can be mainly caused by different hydrogen bonding. Since the 3DMA1P has been demonstrated to have an intramolecular hydrogen bond14 and with the fact that the intramolecular hydrogen bond can lead to the decrease in the boiling point and higher vapor pressure for pure amines,19 it can then be inferred that 1DMA2P can have a more intramolecular hydrogen bond instead of an intermolecular hydrogen bond in solution. On the other hand, the 1DEA2P also follows the same reason to present the lowest boiling point among the DEEA, 1DEA2P, and 3DEA1P. Furthermore, DEEA and DMA-2 M-1P are isomers, and their boiling points are similar; however, the vapor pressure value of DEEA is less than that of DMA-2 M-1P. Actually, DMA-2 M-1P (pKa = 10.34) has stronger basicity than that of DEEA (pKa = 10.01),11 which means DMA-2 M-1P can have greater polarity and better water solubility. However, DMA-2 M-1P has a higher vapor pressure than that of DEEA, and then it can be inferred that DMA-2 M-1P has more intramolecular hydrogen bonding than that of DEEA. This also follows the inference of that branched alkanol chain connected to the N atom can be favorable for formation of heterocyclic structures in molecules. MDEA, DMA-1,2-PD, and DEA-1,2-PD have a much higher boiling point than those of the other tertiary alkanolamines that is mainly owing to one more OH− group addition on their molecule (with more intramolecular hydrogen bond) and a higher molecular weight. 4DEA-2B has a higher boiling point due to its higher molecular weight. Second, the heterocyclic structures caused by the intramolecular hydrogen bond can decrease the water solubilities of amines and protonated amines, and the biphasic phenomenon was observed in 2.5 M fresh aqueous 1DEA2P and CO2-loaded 4DEA-2B solutions, as shown in Figure 4. This result can assist to prove the existence of intramolecular hydrogen bonds in 1DEA2P and 4DEA-2B. No phase change was observed in

4. RESULTS AND DISCUSSION 4.1. Verification of Heterocyclic Structure by Physical Properties. To demonstrate the existence of heterocyclic structures displayed in Figures 2 and 3, some key physical properties of 11 tertiary alkanolamines were collected, including the boiling point and vapor pressure, as shown in Table S2. Then, the boiling point (measured by experiment) and vapor pressure (partially tested by experiments) properties for these pure chemicals were plotted with amines, as shown in Figure 5. It should be specially noted that these data were obtained from documentation of safety data sheets (SDSs) for these chemicals30,31 or predicted by ACD/Labs Module.32 First, the predicted values are systematically off to the experimental vapor pressure values; this could be due to the purity of chemicals was less than 100%, as shown in Table S2. However, based on the qualitative comparison for the experimental and predicted values, it can be observed that the predicted vapor pressure values of DMEA, 1DMA2P, 3DMA1P, DEEA, and MDEA have the same relative order as those of the experiment data (from chemical SDS), indicating the qualitative comparison of the predicted vapor pressure can be reliable. In addition, it can be seen that 1DMA2P has the lowest boiling point and the highest vapor pressure among DMEA, 1DMA2P, 3DMA1P, and DMA-2 M-1P, while 1DEA2P also shows the same results among DEEA, 1DEA2P, and 3DEA1P. The factors that affecting the boiling point of chemicals mainly include the molecular weight, distance between molecules, intermolecular contacting area, and hydrogen bond. According to Figure 1, 1DMA2P and E

DOI: 10.1021/acs.iecr.9b02244 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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these amines displayed a similar trend because the basicity is directly related to the stability of reaction products, that is, the protonated amines and (bi)carbonate.8 In addition, the influences of the intramolecular hydrogen bond and heterocyclic structure on basicity and equilibrium CO2 solubility could also be observed. It can be typically presented by the data of 1DMA2P, 3DMA1P, 1DEA2P, and 3DEA1P. It can be seen that tertiary alkanolamines with linear alkanol chains between amino and hydroxyl groups (3DMA1P and 3DEA1P) have stronger alkalinity and higher equilibrium CO2 solubility than those with branched alkanol chains for isomers. This is because of the branched alkanol chain may result in a stronger intramolecular hydrogen bond and promote the formation of the heterocyclic structure, which will lead to poor basicity and lower CO2 solubility of amines. Additionally, the pH value absence of fresh 2.5 M 1DEA2P solution was because of two phases presented in solution as shown in Figure 4, which is mainly caused by the intramolecular hydrogen bond. Moreover, 3DEA1P solution had slightly higher solubility than that of DMA-2 M-1P, and it may be both caused by different molecular structures and stronger forces of the intramolecular hydrogen bond. On the one hand, owing to the force driven by steric hindrance of two methyl groups connected to α-C, it can weaken the electronwithdrawing inductive effect of the hydroxyl group. On the other hand, electron-donating effects of two methyl groups can also help DMA-2 M-1P with stronger alkalinity. However, the branched alkanol chain may result in stronger forces of intramolecular hydrogen bonds and less possibility of amines toward CO2, then leading to lower equilibrium CO2 solubility. 4.3. Effect of Intramolecular Hydrogen Bond on Reaction Kinetics. To understand the effect of the intramolecular hydrogen bond and heterocyclic structure on the CO2 absorption rate for tested tertiary alkanolamine solutions, the first-order reaction rate constants (k0) of CO2 absorption into fresh tertiary alkanolamine solutions were experimentally determined at 313 K and with molar concentrations of 0.1−0.4 M using stopped-flow techniques, as shown in Figure 7 and

fresh or loaded 1DMA2P solutions that may be owing to a weaker charge density on its N atom of the amino group, leading to the relatively weak intramolecular hydrogen bond and less possibility for heterocyclic structure formation. From the above, it can be concluded that the heterocyclic structure may exist in all tested tertiary alkanolamines, but, obviously, the tertiary alkanolamines with branched alkanol chains (like 1DMA2P and 1DEA2P) can endow a stronger intramolecular hydrogen bond and lead to more amine with a heterocyclic structure of five- or six- membered ring. Correspondently, the amine with the linear alkanol chain (like 3DMA1P and 3DEA1P) can have a relatively weaker intramolecular hydrogen bond due to the possible larger angle strain for heterocyclic structure formation. Despite of these inferences, the existence and strength of the intramolecular hydrogen bond in tested tertiary alkanolamines still need further demonstration with direct evidence or molecule simulations. 4.2. Effect of Intramolecular Hydrogen Bond on Basicity and Equilibrium CO2 Solubility. To further verify the influences of the intramolecular hydrogen bond and heterocyclic structures on CO2 absorption performance, pH values of 2.5 M fresh amine solutions were measured for 10 tested tertiary alkanolamine at 313 K, and pH values as well as equilibrium CO2 solubilities of 2.5 M CO2-loaded amine solutions were measured at 313 K and with a CO2 partial pressure of 15 kPa, as presented in Figure 6. It can be observed

Figure 6. pH values of 2.5 M fresh amines solution at 313 K as well as the pH values and equilibrium CO2 solubility of 2.5 M loaded solutions at 313 and 15 kPa for 10 tested tertiary alkanolamines.

that the pH values of fresh solutions gave out an order of MDEA < DMA-1,2-PD < DMEA < DMEA < 1DMA2P < 3DMA1P < DEA-1,2-PD < DEEA < 3DEA2P < DMA-2 M-1P. The pH values of CO2-loaded solutions at equilibrium gave out a similar result of about 8.83 (range from 8.63 to 8.98). The equilibrium CO2 solubility showed a similar trend of pH value order of fresh solutions, but the CO2 solubility of DMA-2 M-1P was a little less than that of 3DEA1P, despite of its higher basicity. First, the displayed basicity order of fresh tertiary amine solutions was mainly resulted by the electron-donating effect of the alkyl group connected to the amino group and an electronwithdrawing inductive effect of the hydroxyl group.21,23 It is observed that the ethyl groups connected to the amino group can obviously increase the alkalinity of tertiary alkanolamines in solution than those with methyl groups, while the addition of the hydroxyl group or by closing it to the amino group can significantly decrease the basicity of tertiary alkanolamines in solution. Spontaneously, equilibrium CO2 solubility values of

Figure 7. k0 of CO2 absorption into the 11 tested aqueous tertiary alkanolamine solutions at 313 K and with molar concentrations of 0.1−0.4 M.

Table S3. It can be seen that the k0 values also gave out a similar order as that of the pH values for 2.5 M fresh tertiary alkanolamine solutions, that is, MDEA < DMA-1,2-PD < DMEA < DMEA < 1DMA2P < 3DMA1P < DEA-1,2-PD < DEEA < 3DEA2P < DMA-2 M-1P. The order of k0 values F

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CO2 capture performance for tested tertiary alkanolamines, the acid dissociation constant was determined at temperatures of 293−323 K, the ln Ka values were presented in Table 2 and plotted with temperatures in Figure 9. While it can be seen that the ln Ka is directly proportional to the 1/T, in that case, its slope can be constant, which can be used for further calculations of ΔrGm and ΔrHm.

further confirmed the effect of the intramolecular hydrogen bond on the CO2 absorption performance. In addition, the pH values (to determine the [H+]) with CO2 loadings were measured for 0.4 M aqueous MDEA, DMEA, 3DMA1P, and DEEA solutions at 313 K, the calculated molar concentrations of hydroxyl ion ([OH−]) were plotted with CO2 loading in Figure 8, and then the pH

Figure 8. Concentration of hydroxyl ion curves with CO2 loading for aqueous MDEA, DMEA, 3DMA1P, and DEEA solutions.

Figure 9. Plots of ln Ka with 1000/T for 11 tested tertiary alkanolamines. The dashed lines are the fitted results.

values, [OH−], and corresponding CO2 loadings can be seen in Table S4. It should be noted that [OH−] of fresh amines solutions was treated as the CO2 loading equal to 0.001 instead of 0 since the value of X-ray was presented in logarithm to distinguish the change of [OH−] with CO2 loading increasing. It can be observed that [OH−] can be first decreased very fast with little CO2 loading increasing and then a trend to be gentle. This can be mainly explained by the presence of (bi)carbonate and basicity of amines. The reaction products can restrain the dissociation of free amine to the release hydroxyl ion, while the ion strength can be also directly related to the basicity of amines. The most important of all is that [OH−] for 4 tested tertiary alkanolamines kept invariant with different CO2 loadings. Since the basicity and ion strength can be the key-influencing parameter for CO2 absorption,8 the inferred rule may be also fitted for CO2-loaded tertiary alkanolamine solutions also. 4.4. Effect of Heterocyclic Structure on CO2 Capture Performance. It is more appropriate to compare amine CO2 capture performance by comprehensively taking consideration of absorption and desorption rates,33 and the reaction rates of tertiary amines with CO2 can be directly related to basicity (pKa) also.8,9,24 Therefore, to comprehensively evaluate the

Then, the calculated values of ΔrGm and ΔrHm for 11 tested alkanolamines were plotted, as shown in Figure 10. The results

Figure 10. ΔrGm and ΔrHm plots for tested 11 tertiary alkanolamines.

showed that tertiary alkanolamines with linear alkanol chains (DMEA, DEEA, 3DMA1P, and 3DEA1P) can have better CO2 capture performance comparing with the amines with branched alkanol chains or the extra addition of the hydroxyl

Table 2. ln Ka Values of 11 Tested Tertiary Alkanolamines at 293.15−313.15 K ln Ka amines

293.15 K

298.15 K

303.15 K

313.15 K

323.15 K

333.15 K

MDAE DMA-1,2-PD DMEA 1DMA2P 3DMA1P DEA-1,2-PD 4DEA-2B DEEA 1DEA2P 3DEA1P DMA-2 M-1P

−19.94 −20.31 −21.53 −21.90 −22.10 −22.22 −23.86 −22.89 −23.51 −23.65 −23.85

−19.71 −20.01 −21.28 −21.64 −21.85 −21.92 −23.51 −22.61 −23.26 −23.39 −23.58

−19.46 −19.73 −21.07 −21.39 −21.58 −21.62 −23.25 −22.29 −22.98 −23.12 −23.33

−18.86 −19.23 −20.52 −20.84 −21.11 −21.11 −22.79 −21.81 −22.43 −22.61 −22.70

−18.49 −18.74 −20.06 −20.31 −20.59 −20.56 −22.22 −21.30 −21.83 −22.08 −22.06

−17.91 −18.19 −19.48 −19.80 −20.15 −20.06 −21.59 −20.72 −21.23 −21.55 −21.58

G

DOI: 10.1021/acs.iecr.9b02244 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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avoid the influences of the intramolecular hydrogen bond to present relatively better CO2 capture performance.

group, and this can be mainly explained by the influences of the intramolecular hydrogen bond and electron-withdrawing effect of the hydroxyl group. In addition, 4DEA-2B and DMA2 M-1P can have fast reaction rates with CO2, but their desorption rate could be relatively slow, and this can be caused by both strong basicity and the effect of the intramolecular hydrogen bond. 1DMA2P, 1DEA2P, and DEA-1,2-PD had relatively slow reaction and desorption rates for CO2 capture also may due to the strong forces of the intramolecular hydrogen bond. Finally, the tertiary alkanolamines with dimethyl amino groups (for DMEA, 1DMA2P, 3DMA1P, and DMA-1,2-PD) and diethyl amino groups (for DEEA, 1DEA2P, 3DEA1P, and DEA-1,2-PD) showed a similar trend just as their physical properties (boiling point and vapor pressure) and chemical properties (basicity, equilibrium CO2 solubility and kinetics). However, tertiary alkanolamines with dimethyl amino groups can have faster desorption and slower absorption rates than those of tertiary alkanolamines with diethyl amino groups. This is mainly due to the electrondonating effect and different alkalinity caused by different alkyl groups connected to the amino group. To further validate the screening result above, solventscreening experiments were performed for 2.5 M aqueous DMEA, 1DMA2P, 3DMA1P, and DEEA, and the cyclic capacity plots with the cyclic time as presented in Figure 11. It

5. CONCLUSIONS This work proved that intramolecular hydrogen bonds can exist in free tertiary alkanolamines and protonated amines, mainly based on comparing physical properties such as boiling point, vapor pressure, and water solubility as well as the fact that the intramolecular hydrogen bond exist in 3DMA1P. In addition, the tertiary alkanolamines with branched alkanol chains (like 1DMA2P and 1DEA2P) can result in stronger forces of intramolecular hydrogen bonds. However, the accurate evaluations of existence and strength of intramolecular hydrogen bonds in tertiary alkanolamines still need further demonstration with direct evidence. The intramolecular hydrogen bonds will affect the chemical properties of tertiary alkanolamines, especially the absorption and desorption performance for CO2 capture. The experimental results showed that the intramolecular hydrogen bond and formation of heterocyclic structures in amines with branched alkanol chains (like 1DMA2P and 1DEA2P) can relatively decrease the basicity (pKa), equilibrium CO2 solubility, reaction rate constant (k0), slower desorption rate, and cyclic capacity for CO2 capture in aqueous solution. Hence, it can be concluded that tertiary alkanolamines with linear alkanol chains (like 3DMA1P and 3DEA1P) can present better CO2 capture performance. Also, the work suggests that the selection or design of tertiary alkanolamines should with linear alkanol chains instead of branched in terms of CO2 capture applications or future researches.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.iecr.9b02244.

Figure 11. Cyclic capacity at cyclic times of 30−90 min for 2.5 M aqueous DMEA, 1DMA2P, 3DMA1P, and DEEA solutions.

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Calculation of Ka; validations of determination of Ka values and repeated experiments; definition, calculation, and experiment procedure of cyclic capacity; validation of equilibrium CO2 solubility (PDF)

AUTHOR INFORMATION

Corresponding Authors

can be observed that the four tertiary amines gave out the value order of cyclic capacity as 1DMA2P < DMEA < DEEA < 3DMA1P. The 3DMA1P obtained the highest cyclic capacity among the four amines, while 1DMA2P showed the lowest cyclic capacity. This result also confirmed that the tertiary alkanolamines with the branched alkanol chain (like 1DMA2P) can lead to poor CO2 capture performance. The results further demonstrated that tertiary alkanolamines with branched alkanol chains could lead to a stronger intramolecular hydrogen bond than those with linear alkanol chains. The formed heterocyclic structure can decrease the charge density on the N atom and lead to lower basicity, CO2 absorption capacity, and reaction rate constant. On the other hand, the intramolecular hydrogen bond may be also occurred for protonated amines in solution, resulting in a more energy requirement for proton dissociation of the protonated amines and a slower desorption rate for CO2 stripping (like 1DMA2P, 1DEA2P, DMA-2 M-1P, and 4DEA-2B). Correspondingly, the tertiary alkanol amines with linear alkanol chains can relatively

*E-mail address: [email protected]. Phone: +8615116365674 (H.G.). *E-mail address: [email protected]. Phone: +8613618481627. Fax: +86-731-88573033 (Z.L.). ORCID

Zhiwu Liang: 0000-0003-1935-0759 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The financial support from the National Natural Science Foundation of China (NSFC nos. 21536003, 21706057, 21878073, and 51521006), the Natural Science Foundation of Hunan Province in China (no. 2018JJ3033), and the China Outstanding Engineer Training Plan for Students of Chemical Engineering & Technology in Hunan University (MOE no. 2011-40). H

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