Mechanism and Kinetics of CO2 Absorption into an Aqueous Solution

Jan 5, 2017 - A novel ionic liquid (1-aminoethyl-3-methylimidazolium lysinate, [C2NH2MIm][Lys]) functionalized with three amino groups had been develo...
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Mechanism and Kinetics of CO2 Absorption into an Aqueous Solution of a Triamino-Functionalized Ionic Liquid Xiaobin Zhou, Guohua Jing,* Fan Liu, Bihong Lv, and Zuoming Zhou College of Chemical Engineering, Huaqiao University, Xiamen 361021, China ABSTRACT: A novel ionic liquid (1-aminoethyl-3-methylimidazolium lysinate, [C2NH2MIm][Lys]) functionalized with three amino groups had been developed for carbon dioxide (CO2) capture in this work. The CO2 absorption loading of [C2NH2MIm][Lys] solution was found to be 1.59 mol CO2/mol IL with a concentration of 0.5 mol/L at 313.15 K, which was much higher than that of the most existing dual functionalized ILs. Besides, [C2NH2MIm][Lys] also owned a good regenerability even after 6 regeneration cycles. According to the results of 13C nuclear magnetic resonance (13C NMR), CO2 absorption into [C2NH2MIm][Lys] solution could be divided into two stages, which began with the formation of carbamate, and followed by the hydration of CO2 to form carbonate/bicarbonate. Meanwhile, the desorption of the saturated [C2NH2MIm][Lys] was proven to be a reverse process of the adsorption. Based on the mechanism results, the kinetics of CO2 capture into [C2NH2MIm][Lys] solution was investigated by using a double stirred-cell absorber at temperatures ranging from 303 to 333 K. Under the pseudofirst-order regime, the overall reaction rate constants (kov) and the forward second-order rate constants (k2) under different concentrations and temperature were obtained, which were both increased considerably as temperature increased. Moreover, the values of enhancement factor (E) were linear with CRNH2 and temperature. The Arrhenius equation of CO2 absorption was also estimated, and the activation energy was calculated to be 25.5 kJ·mol−1.

1. INTRODUCTION As one of the main greenhouse gases, carbon dioxide (CO2), which generates from fossil fuel combustion and many other industrial processes, is responsible to contribute the greenhouse effect and global warming. In this regard, postcombustion carbon capture and sequestration (CCS) technology is a promising route to achieve a meaningful reduction in CO2 emissions.1,2 The traditional industrial approach commonly employed for CO2 capture is based on the use of aqueous amine solutions (e.g., monoethanol amine (MEA)) or chilled ammonia.3,4 However, these solvents not only have disadvantages, such as toxicity, equipment corrosion, and solvent loss but also require a large amount of energy for solvent regeneration.2,5 Recently, ionic liquids (ILs) have been emerging as potential contenders for CO2 capture due to their superior physicochemical characteristics, including low melting point, high thermal stability, adjustable structure, and good recyclability.6−9 However, the solubility of CO2 in conventional ILs is limited due to the physical absorption. In order to achieve better performance, some special groups (e.g., −NH2, −OH) were introduced to the anion or the cation of ILs. The aminefunctionalized IL has been chosen as the most promising candidate for CO2 capture. In 2002, Bates and co-workers had reported the first example for CO2 chemisorption by an aminefunctionalized, imidazole-derived task-specific IL.10 After that, the amine-functionalized ILs for CO2 capture had been developed widely.12−14 Also, Zhang et al.13 found that a dual amine-functionalized cation-tethered IL absorbed 1.08 mol CO2/mol IL at 1 atm and 293 K. Hu et al.14 found that a poly amino-based IL exhibited a 2.04 mol CO2/mol IL with the water content of 40% at 298 K. These reported works indicated that the CO2 absorption capacity of the functionalized ILs related to the amount of amino groups, and the greater amount © XXXX American Chemical Society

of amino groups in the ILs, the more CO2 could be absorbed by the solvents. Thus, designing novel multi amine-functionalized ILs was inspired for CO2 capture during these years. Due to having identical functional group as alkanolamines, amine-functionalized ILs are expected to have high reactivity toward CO 2 . The knowledge of investigation on CO 2 absorption mechanism is of great importance to understand the characteristics of the CO2 capture process. It was reported that cation-functionalized ILs capture CO2 in a way similar to the reaction mechanism between CO2 and the primary amines proposed by Crooks and Donnellan.15 During the CO2 absorption, amine groups first reacted with CO2 to form carbamate and ammonium-appended amine species in a stoichiometry of 2:1 (−NH2:CO2)10,11 and then the carbamate decomposed further into bicarbonate with the continuous addition of CO2.7,12 Some researchers reported anion-functionalized ILs which could react with CO2 to form carbamic acid in a ratio of one CO2 per one amine (1:1 stoichiometry).16 Nevertheless, the attention is mainly focused on the CO2 absorption of single and dual group functionalized ILs, there are few works on the multi amine-functionalized ILs. Especially, the component of the multi amine-functionalized ILs solution will be more complex when amino groups are introduced to both anion and cation. Therefore, CO2 reaction into a multi amine-functionalized ILs solution system needs further research. Furthermore, the understanding of the kinetics phenomena of CO2 reaction with absorbents is essential for the effective design of the CO2 absorption process. Some researchers have already studied the reactive absorption kinetics Received: November 10, 2016 Revised: January 5, 2017 Published: January 5, 2017 A

DOI: 10.1021/acs.energyfuels.6b02963 Energy Fuels XXXX, XXX, XXX−XXX

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Figure 1. Schematic diagram of synthesis of [C2NH2MIm][Lys]. spectra of the samples were obtained in the range of 400−4000 cm−1 with an IR spectrometer (8400s, Shimadzu Corporation, Japan). 13C nuclear magnetic resonance (13C NMR) analysis was carried out on a NMR spectrometer (Bruker AVIII500 MHz) to explore the products of CO2 absorption into [C2NH2MIm][Lys]. The sample (0.5 mL) was added to the nuclear magnetic tubes for analysis, using an internal standard of 0.1 mL D2O for the deuterium lock. 2.4. Absorption and Desorption Experiments. The absorption experiments were carried out in a laboratory bubbling reactor with a diameter of 2.5 cm and a height of 25 cm, which had been used in our previous work.19 The concentration of [C2NH2MIm][Lys] solution was 0.5 mol·L−1, and its volume was 25 mL. The gas of CO2 was continuously supplied to the reactor with a flow rate of 30 mL·min−1 by using a mass flow controller. The reactor was immersed in a constant temperature water basin at a determined temperature. The inlet and outlet gas flow rates from the reactor were measured by a soap-film flowmeter. The reaction time was recorded by electric stopwatch. During the absorption process, temperature and pH of the solvent were measured by acidity meter (FE20, Mettler-Toledo, International Inc.). The solvent was considered to be saturated when the outlet gas flow rate was equal to the inlet gas flow rate. After absorption, the CO2saturated solutions were regenerated in a three-necked flask in oil bath at 120 °C for 90 min. During desorption, a reflux condenser was connected to the reactor outlet to minimize the water vapor loss. After regenerated, the solvent was repeated to absorb CO2. 2.5. Kinetics Experiments. The kinetics experiments for CO2 absorption into [C2NH2MIm][Lys] solution were carried out in a double stirred-cell absorber, which had been described in our previous work.20 This absorber has a defined gas−liquid interface as the mass transfer area, and is convenient for kinetic investigation. CO2 gas was continuously supplied to the absorber with a flow rate of 100 mL· min−1. The flow rates of the inlet and outlet gas were measured by soap-film flow meter. The CO2 absorption rate can be calculated as follows:

of CO2 in some amine-functionalized ILs.17,18 However, kinetics information for triamino-functionalized ILs absorbent is rare in literature and needs to be revealed. In the present work, a novel multi amine-functionalized IL (1-aminoethyl-3-methylimidazolium lysinate, [C2NH2MIm][Lys]) was developed for CO2 capture, which contained three amino groups for the purpose of higher CO2 absorption capacity. The characterization and performance of [C2NH2MIm][Lys] were investigated. In order to get a deep understanding on the reaction mechanism of CO2 capture, 13C nuclear magnetic resonance (NMR) was used to analyze the reaction intermediate during the absorption and desorption. Moreover, the kinetics of the reaction between CO2 and [C2NH2MIm][Lys] was also studied. The enhancement factor (E), the overall reaction kinetics constant (kov), and the secondorder rate constant (k2) were all calculated.

2. MATERIALS AND METHODS 2.1. Chemicals. 1-Aminoethyl-3-methylimidazolium bromide ([C2NH2MIm][Br]) (>98.0%) was purchased from Lanzhou Greenchem ILS, LICP, CAS, China. Lysine (>99.0%) was received from Chengdu Xiya Chemical Reagent Co. Ltd., China. D2O was provided by J&K Scientific Ltd. The gas of CO2 (>99.999%) was supplied by Fujian Xiamen Kongfente Gas Co., Ltd., China. In addition, all other chemicals used in the study were of analytical grade. All solutions were prepared with ultrapure water. 2.2. Synthesis of the Functionalized ILs. The multi aminefunctionalized IL ([C2NH2MIm][Lys]) was synthesized by introducing functionalized groups of amine and lysine to the imidazole-based ILs. The preparation process containing replacement and neutralization had been described in detail in our previous work,19 and the schematic diagram of synthesis is depicted in Figure 1. To be brief, 0.5 mol·L−1 of [C2NH2MIm][Br] solution was slowly added to the ion exchange column to form [C2NH2MIm][OH]. Then the equalmolar neutralization reaction between [C2NH2MIm][OH] and lysine was carried out in a three-necked flask for 24 h at 25 °C under stirring. The product was rinsed by methanol and ethyl acetate and concentrated with a rotary evaporator method. Finally, the product was dried under vacuum overnight before use. 2.3. Characterization. The thermal characterization of the samples was analyzed by thermo gravimetric analysis (TGA, DSC910/SDT2960, TA Instruments Co. Ltd., USA) by heating the samples from room temperature to 800 °C with a heating rate of 10 °C min−1 in N2 atmosphere. Fourier transform infrared (FTIR)

N=

P(Q out ‐Q in) (1)

ART −2 −1

where, N denotes CO2 absorption rate (mol·m ·s ). P, A, R, and T are pressure (Pa), gas−liquid interfacial area (m2), ideal gas law constant (kmol·m−3·s−1), and temperature (K), respectively. Qin and Qout are the inlet and outlet gas flow rates (m3·s−1), respectively. 2.6. Physicochemical Properties. 2.6.1. Viscosity. The viscosity of the solution was measured by using the digital rotational viscometer (NDJ-5S, Shanghai Bangxi Instrument Technology Co., Ltd., China). Herein, the sample of the solvent was set into a small beaker and B

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Energy & Fuels immersed in a constant temperature water basin at a determined temperature. 2.6.2. Diffusion Coefficient. Since [C2NH2MIm][Lys] solution is electrolyte solution, the diffusivity of CO2 in the aqueous electrolyte solution can be calculated from the relations, given by Ratcliff and Holdcroft:21

DCO2 = DCO2,w (1 − KC)

(2)

⎛ − 2119 ⎞ ⎟ DCO2,w = 2.35 × 10−6exp⎜ ⎝ T ⎠

(3)

3. RESULTS AND DISCUSSION 3.1. Characterization. The thermal stability is a vital physicochemical property of ILs because high thermal stability helps IL to achieve good reversibility for CO2 capture. The thermogravimetric (TG) curves of [C2NH2MIm]Br and [C2NH2MIm][Lys] are compared in Figure 2a. It was found

where DCO2 and DCO2,w denote the diffusion coefficient of CO2 in salt solvent and pure water, respectively (m2·s−1), C is the concentration of salt solution (mol·L−1), K is the constant which depends on the temperature and can be expressed as follows: (4)

K = 0.03 + 0.55ω ω can be obtained by eq 5: μw = μ(1 − ωC)

(5)

where μw and μ is the viscosity of pure water, respectively (mPa·s). According to the Ratcliff’s equation, the diffusion coefficient of CO2 (D) is related to solution viscosity and electrolyte concentration. Thus, the diffusion coefficient of CO2 into the IL aqueous solution can be simplified as follows: μw ⎞ ⎡ ⎛ − 2119 ⎞⎤⎛ ⎟ ⎜0.45 − 0.03C + 0.55 DCO2 = ⎢2.35 × 10−6exp⎜ ⎟ ⎥ ⎝ T ⎠⎦⎝ ⎣ μ⎠ (6) 2.6.3. Solubility. The solubility of CO2 into amines and functionalized ILs in the most reported works were estimated by N2O analogy. Nevertheless, [C2NH2MIm][Lys] would form ions in the solution, and the ions had significant effect on the solubility of CO2 gas. Thus, the N2O analogy cannot be simply extended for [C2NH2MIm][Lys] solution. Instead, the model of Schumpe,22 which takes the salting out effect of electrolyte solution into account, can be used to describe the solubility of gases in ionic solution. ⎛ HCO ,w ⎞ 2 ⎟⎟ = log⎜⎜ ⎝ HCO2 ⎠

∑ (hi + hg)Ci (7)

here, HCO2, w and HCO2 denote the Henry constants (mol·L−1·Pa−1) in pure water and in the salt solution, respectively. hi is the ion-specific parameter for either the cation (h+) or the anion (h−), hg denotes the gas specific parameter (L·mol−1), and Ci is the concentration of ion i in the solution. HCO2, w can be calculated as a function of temperature according to the following equation:

⎛ 2044 ⎞ ⎟ HCO2,w = 3.54 × 10−7exp⎜ ⎝ T ⎠

Figure 2. Characterization of [C2NH2MIm]Br and [C2NH2MIm][Lys]. (a) TG analysis; (b) FTIR spectra.

that the maximum weight loss temperature of [C2NH2MIm][Lys] was 230 °C, which was lower than that of [C2NH2MIm] Br (360 °C). The results indicated that the thermal stability of [C 2 NH 2 MIm][Lys] was slightly lower than that of [C2NH2MIm]Br. Fortunately, the regeneration temperature of the absorbent for CO2 capture in the practical application often sets at 120 °C, which is far below the decomposition temperature of [C2NH2MIm][Lys]. Thus, the thermal stability of [C2NH2MIm][Lys] was still high enough for CO2 capture. The FTIR spectra of [C2NH2MIm]Br and [C2NH2MIm][Lys] are presented in Figure 2b. As seen in Figure 2b, there are some obvious peaks appearing in the spectra of [C2NH2MIm]Br, such as 3423 and 1630 cm−1 assigned to −NH2, 1576 cm−1 assigned to CN and/or CC in the imidazole ring, 1173 cm−1 assigned to C−N.13 Besides the similar results with [C2NH2MIm]Br, there existed some special points in the spectra of [C2NH2MIm][Lys]. The peak around 3400 cm−1 in the spectra of [C2NH2MIm][Lys] became broad, which could be attributed to the vibrations of −NH2 groups in [Lys]−.

(8)

When the value of hi for an ion is unknown, the value for a similar species can be used. Therefore, the values of hRNH3+ can be replaced by the value of hNH4+.23 The values of the lysine anion specific parameter (hLys‑) and hg were taken from literature.24 The values of hNH4+, hLys‑, and hT are presented in Table 1.

Table 1. Values of hNH4+, hLys‑, and hT T (°C)

hNH4+ (dm3·mol−1)

hLys‑ (dm3·mol−1)

hg (dm3·mol−1)

303 313 323 333

−0.0737 −0.0737 −0.0737 −0.0737

0.1020 0.0886 0.0786 0.0686

−0.01908 −0.02296 −0.02573 −0.02916 C

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Energy & Fuels Meanwhile, the peak around 1630 cm−1 of [C2NH2MIm][Lys] turned into smoothness because the vibration signal of −COO group from [Lys]− overlapped with that of CN and/or C C in the imidazole ring. It was worth noting that there was a new peak emerged at 1400 cm −1 in the spectra of [C2NH2MIm][Lys], which was caused by the vibration of C−N in the [Lys]−. Based on the results, it was proven that [C2NH2MIm][Lys] was synthesized in this work. 3.2. Absorption and Desorption Performance of [C2NH2MIm][Lys]. 3.2.1. CO2 Absorption into [C2NH2MIm][Lys] Solution. The absorption loadings of CO 2 into [C2NH2MIm][Lys], [C2NH2MIm]Br, lysine solutions are compared in Figure 3. It was found that the absorption rate

comparable with that of [APmim][Lys] (6.71 mol of CO2/kg of IL). Meanwhile, [C2NH2MIm][Lys] had a higher CO2 absorption loading than the reported dual amino-functionalized ILs due to a higher number of the functionalized amino groups. Thus, increasing the number of the functionalized amino group of ILs could significantly increase the absorption loading of the absorbent. 3.2.2. Effect of Temperature on CO2 Absorption. Since the temperature of the typical gas after flue gas desulfurization (FGD) process often fluctuated in the range of 45−55 °C, CO2 capture into [C2NH2MIm][Lys] solution at different temperature varied from 30 to 60 °C is investigated, and the results are presented in Figure 4a. It was easy to observe that the CO2

Figure 3. CO2 absorption into [C2NH2MIm][Lys], [C2NH2MIm]Br, lysine solutions. {QCO2: 30 mL/min; C: 0.5 mol/L; VL = 25 mL; T: 313.15 K}.

and loading of [C2NH2MIm][Lys] were significantly higher than that of [C2NH2MIm]Br and lysine. The CO2 absorption loading of the saturated [C2NH2MIm][Lys] was found to be 1.59 mol CO2/mol IL, which was higher than the sum loading of [C2NH2MIm]Br (0.39 mol CO2/mol IL) and lysine (0.76 mol CO 2 /mol IL). Moreover, the comparison of CO 2 absorption by [C2NH2MIm][Lys] and the reported aminofunctionalized ILs are shown in Table 2. The results showed that with the same number of the amino groups, [C2NH2MIm][Lys] exhibited slightly lower CO2 absorption loading than that of [APmim][Lys]. However, due to the relatively low molecular weight, the CO2 loading per kg of [C 2NH2MIm][Lys] (6.01 mol of CO2/kg of IL) was

Figure 4. Effect of temperature on (a) CO2 loading and (b) viscosity of [C2NH2MIm][Lys] solution. {QCO2: 30 mL/min; C: 0.5 mol/L; VL = 25 mL }.

Table 2. Comparison of CO2 Absorption by [C2NH2MIm][Lys] and Reported Amino-Functionalized ILs

ionic liquids DAIL [APmim][Gly] [aP4443][AA] [P66614][Pyr] [aemmim] [Tau] [APmim][Lys] [C2NH2MIm] [Lys]

number of amino group

CO2 loading, temperature, K mol CO2/mol IL

loading increased as temperature increased from 30 to 60 °C in early stage of the reaction. Nevertheless, the loading of the saturated solvents at different temperature was close to each other and all kept around at 1.55 mol CO2/mol IL. The results indicated that [C2NH2MIm][Lys] was stable for CO2 capture at different temperature. Many researchers proved that the viscosity of absorbent increased after CO2 absorption, which might impact the diffusion of CO2 in solution and then decrease the absorption rate.29−31 Herein, the viscosity of [C2NH2MIm][Lys] aqueous solution during the absorption at different temperature was also measured. As shown in Figure 4b, the viscosity of

reference

two two two two two

303.15 303.15 298.15 303.15 303.15

1.05 1.23 1.0 1.02 0.9

13 19 25 26 27

three three

323.15 313.15

1.8 1.59

28 this work D

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Energy & Fuels [C2NH2MIm][Lys] aqueous solution decreased as the temperature increased at the same reaction time. Meanwhile, it increased slightly as the reaction time increased at a certain temperature. The results confirmed that the viscosity of [C2NH2MIm][Lys] also increased after CO2 absorption, but the increase was very small, which was beneficial to the actual operation. 3.2.3. Regeneration of CO2-Saturated [C2NH2MIm][Lys] Solution. The regenerability of [C2NH2MIm][Lys] was investigated to evaluate its long-term stability. As shown in Figure 5, the regeneration efficiency of [C2NH2MIm][Lys]

was an acid gas, pH value of the solution continuously decreased as the gas flowed into the absorber. It was worth to note that the temperature increased first and decreased subsequently during the absorption process. The results indicated that the absorption could be divided into two stages according to the change trend of the temperature. In the first stage, the temperature increased from 28.5 °C to its maximum of 33.5 °C and pH decreased from 10.5 to 8.5. In the second stage, the temperature of the solution decreased gradually and finally closed to the ambient temperature, meanwhile, pH tended to be stable. Compared to the second stage, the CO2 absorption rate of the first stage was extremely fast. The result indicated the exothermic reaction in the first stage and the endothermic reaction in the second stage, which was similar to the CO2 capture into MEA and [APmim][Gly] solution.19,32 3.3.2. 13C NMR. To gain a deep insight into the absorption mechanism, the 13C NMR spectra of [C2NH2MIm][Lys] solution at different CO2 absorption loading were analyzed. Since pH value of the solution was continuously decreased as the reaction time increased, pH value was used to evaluate the CO2 capture process with different CO2 loading. The result is depicted in Figure 7.

Figure 5. Regenerability of [C2NH2MIm][Lys] solution after multiple desorption. {QCO2: 30 mL/min; C: 0.5 mol/L; VL = 25 mL; Tabsorption: 313.15 K; Tdesorption: 393.15 K}.

slightly decreased as the regeneration cycles increased. Nevertheless, the regeneration efficiency still kept above 90% after six absorption−desorption cycles, which indicated [C2NH2MIm][Lys] could be repeatedly used after multiple thermal desorption. 3.3. Mechanism of CO2 Capture into [C2NH2MIm][Lys] Solution. 3.3.1. pH and Temperature Change of the Solution During CO2 Absorption. To understand the CO2 absorption process, the changes of temperature and pH of the [C2NH2MIm][Lys]−CO2 system were investigated at ambient temperature, and the results are shown in Figure 6. Since CO2 Figure 7. 13C NMR spectra of [C2NH2MIm][Lys] during absorption process.

As mentioned in Section 3.3.1, CO2 absorption into [C2NH2MIm][Lys] could be divided into exothermic and endothermic process, indicating the existence of different reactions in these two stages. The 13C NMR spectra also proved this conclusion. At the beginning of the reaction, there was a signal (2) at 163.8 ppm revealed the formation of carbamate carboxyl carbon.13 As the reaction carried on, the peak intensity of carbamate was increased, and pH value of the solution changed from initial 10.49 to 9.30. In this period, there was no signal about HCO3−/CO32−, which indicated that the formation of carbamate was the main reaction. When pH value of the solution was lower than 8.83, another new signal (3) appeared at 160.72 ppm revealing the formation of HCO3−/ CO32−. Herein, HCO3− and CO32− appeared as a single peak due to the rapid proton exchange.12 As the reaction carried on, the peak intensity of HCO3−/CO32− increased, while the peak

Figure 6. Changes in solution temperature and pH during CO2 absorption. {QCO2: 80 mL/min; C: 0.5 mol/L; VL = 25 mL}. E

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IL (i.e., C2NH2MIm-COO−, Lys-COO−, C2NH2MIm-NH3+, Lys-NH3+), which followed the zwitterion mechanism and was an exothermic reaction. As the absorbent became saturated, the hydration reaction of CO2 into water and decomposition of carbamate were enhanced, which was an endothermic reaction. Moreover, the desorption of the saturated [C2NH2MIm][Lys] was proven to be a reverse process of the adsorption. 3.4. Kinetics Analysis of CO 2 Absorption into [C2NH2MIm][Lys] Solution. 3.4.1. Chemical Reaction Rate. As mentioned above, CO2 absorption into [C2NH2MIm][Lys] first followed the zwitterion mechanism, which was generally used to describe the reaction between CO2 and primary and secondary alkanolamine.33,34 During this stage, CO2 reacted with [C2NH2MIm][Lys] to form unstable zwitterion and subsequently deprotonated by a base in the solution (here, [C2NH2MIm][Lys] is denoted as RNH2).

intensity of carbamate decreased gradually. During the whole CO2 absorption process, the signal of the carbonyl carbon in lysine anion shifted from 182.47 ppm to its neutral position 174.71 ppm, confirming that the CO2 absorption began with the absorption of the jettisoned protons during carbamate formation and continued with the consumption of carbamate, giving way to carbonate/bicarbonate.12 Eventually, the main CO2 capture approach is the formation of carbonate/ bicarbonate. Based on the 13C NMR spectra, it could be found that the carbamate was hydrolyzed by the H+, while HCO3−/CO32− generated as the excess CO2 supplied in this stage. Besides, the mechanism of CO2 desorption from CO2saturated solution was also investigated. The 13C NMR spectra of the samples at different CO2 loadings during the regeneration process are compared in Figure 8. Similar to the

k2

CO2 + RNH 2 XooY RNH+2 COO−

(9)

k −1

k Bi

RNH+2 COO− + Bi → RNHCOO− + BiH+

(10) −

where Bi was supposed to act as a base (such as H2O, OH , or [C2NH2MIm][Lys]) in the solution, which deprotonated the zwitterion. When assuming quasi-steady-state conditions for intermediate zwitterion concentration and a first-order behavior of CO2, the reaction rate of CO2 in [C2NH2MIm][Lys] solutions can be given by eq 11: −rCO2 − RNH2 =

k 2CCO2C RNH2 1+

k −1 ∑i k BiC Bi

(11)

where k2 and k−1 are the second-order forward and reverse reaction rate constant, respectively, L·mol−1·s−1, and ∑kBiCBi indicates the contribution of the bases present in solution for proton removal. Since the deprotonation of the zwitterion is relatively fast when compared to the reversion rate of CO2, kBiCBi is very large and (k−1/(∑ikBiCBi)) can be neglected.35 Eq 11 can be simplified as follows:

Figure 8. 13C NMR results of [C2NH2MIm][Lys] during desorption process.

absorption, the desorption process also could be divided into two stages. In early stage of the desorption, pH increased gradually as the CO2 desorbed from the solution. When pH increased from 7.28 to 8.78, the signal of HCO3−/CO32− faded away while that of the carbamate increased. During this period of regeneration, some of the HCO3−/CO32− desorbed CO2 under thermolysis and some of them reacted with RNH3+ to form carbamate. Whereafter, the peak intensity of carbamate decreased slightly as the pH increased from 8.78 to 10.53, and the signal of carbonyl carbon in lysine anion returned back from the neutral position to its basic position. On this stage, carbamate began to react with H+ to form [C2NH2MIm][Lys] and desorbed CO2 after the complete decomposition of HCO3−/CO32−. It was observed that a weak characteristic peak of carbamate still existed, which indicated this absorbent could not be completely regenerated. Nevertheless, the regeneration efficiency after sixth thermal desorption even kept at 90% in our experiments, which indicated a well regeneration of [C2NH2MIm][Lys]. On the basis of the above results, the schematic mechanism of CO2 absorption and desorption in [C2NH2MIm][Lys] solution is depicted in Figure 9. CO2 absorption into [C2NH2MIm][Lys] started to form carbamate and protonated

−rCO2 − RNH2 = k 2CCO2C RNH2

(12)

Meanwhile, the hydration reactions of CO2 may also take place in the aqueous solution: ks

CO2 + H 2O ⇄ HCO−3 + H+

(13)

k OH−

CO2 + OH− XooooY HCO−3

(14)

The reaction rate of CO2 in the aqueous solution can be written as −rCO2 − H2O = (ks + k OH−COH−)CCO2

(15)

Moreover, due to the low stability, carbamate may readily undergo hydrolyis in water to form HCO3− (eq 16). Therefore, bicarbonate ions will be present in a larger amount than carbamate ions. RNHCOO− + H+ + H 2O → RNH3+ + HCO3−

(16)

The overall rate of all CO2 reactions in [C2NH2MIm][Lys] aqueous solution is given by the sum of eqs 12 and 15: F

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Figure 9. Mechanism of CO2 capture into [C2NH2MIm][Lys] solution.

Table 3. Diffusivity and Solubility of CO2 in [C2NH2MIm][Lys] Aqueous Solution T, K

C, mol·L−1

μw/μ

DCO2, 10−9 m2·s−1

HCO2, 10−4 mol·m−3·Pa−1

303

0.25 0.50 0.75 1.00 0.50 0.50 0.50

0.8066 0.6432 0.5932 0.4049 0.7170 0.7294 0.7560

1.91 1.70 1.63 1.39 2.24 2.78 3.45

2.81 2.62 2.45 2.28 2.18 1.80 1.52

313 323 333

the enhancement factor, which reflects the impact of chemical reaction on mass transfer. According to the phase balance in the gas−liquid interface, the concentration of CO2 in the interface can be calculated as follows:

−roverall = [k 2C RNH2 + (ks + k OH−COH−)]CCO2 = kovCCO2 (17) −1

where kov denotes the observed reaction rate constant (s ) and is given by eq 18: kov = k 2C RNH2 + (ks + k OH−COH−)

⎛ N⎞ CCO2,i = HCO2PCO2,i = HCO2⎜PCO2 − ⎟ kG ⎠ ⎝

(18) −

The contribution of ks can be neglected. At low OH concentration, the contribution of kOH− in water can also be neglected. Therefore, kov is expressed by kov = k 2C RNH2

When CO2 concentration in the bulk liquid is negligible, based on eqs 21 and 22, the mass transfer rate can be expressed by eq 23:

(19)

3.4.2. Mass Transfer. The mass transfer rate of CO2 into [C2NH2MIm][Lys] solution was enhanced by the presence of a chemical reaction. According to the mass transfer theory, the overall mass transfer rate of CO2 can be expressed by the following equations: N = k G(PCO2 − PCO2,i)

(20)

N = EkL(CCO2,i − CCO2,L)

(21)

(22)

N=

PCO2 1 1 + k EkLHCO2 G

(23)

In eq 23, the value of 1/kG is much smaller than that of 1/ (EkLHCO2), thus, a simple equation for mass transfer rate is obtained:

N = HCO2PCO2EkL

(24)

For the fast pseudo-first-order reaction regime, requires that the Hatta number Ha can fulfill the following criterion:

where kG (mol·m−2·s−1·Pa−1) and kL (m·s−1) are the gas and liquid phase mass transfer coefficient, respectively, the values of which have been calculated in our previous work.36 E denotes

3 < Ha ≪ E∞ − 1 G

(25) DOI: 10.1021/acs.energyfuels.6b02963 Energy Fuels XXXX, XXX, XXX−XXX

Article

Energy & Fuels Table 4. Kinetic Data for CO2 Absorption in [C2NH2MIm][Lys] Aqueous Solution T, K

C, mol·L−1

N, 10−3mol·m−2·s−1

E

Ha

kov, s−1

k2, m3·kmol−1·s−1

303

0.25 0.50 0.75 1.00 0.50 0.50 0.50

7.19 7.99 8.97 9.66 8.25 8.97 9.80

99.40 142.76 193.47 224.81 155.69 180.24 196.74

99.40 142.76 193.47 224.81 155.69 180.24 196.74

2501.94 5797.85 11142.01 17643.97 7156.75 10234.43 14235.97

10007.76 11595.71 14856.01 17643.91 14313.51 20468.86 28471.93

313 323 333

The enhancement factor E is a function of the Hatta number, for the fast pseudo-first-order reaction regime, E can be considered equal to Ha: E = Ha =

DCO2kov kL

(26)

3.4.3. Reaction Kinetics. According to the eqs 19, 24, and 26, the absorption flux of CO2 into [C2NH2MIm][Lys] solution is further transformed to N = HCO2PCO2 DCO2kov = HCO2PCO2 DCO2k 2C RNH2 (27)

k2 can be transformed to k2 =

N2 (HCO2PCO2)2 DCO2C RNH2

(28)

The estimation of CO2 diffusivity (DCO2) and Henry constant (HCO2) in [C2NH2MIm][Lys] aqueous solution is presented in Table 3. At a constant temperature, both DCO2 and HCO2 of CO 2 in [C 2NH 2MIm][Lys] solution decreased as the concentration of [C2NH2MIm][Lys] increased, which was in accord with the “salting-out” effect. The DCO2 increased while HCO2 deceased as the temperature increased at a certain concentration of [C2NH2MIm][Lys] solution. The kinetics data for CO2 absorption into [C2NH2MIm][Lys] aqueous solution is summarized in Table 4. The value of E and kov for CO2−[C2NH2MIm][Lys] system increased as the concentration of [C2NH2MIm][Lys] and temperature increased. It is obvious to find that the values of E were linear with CRNH2 and T (Figure 10). According to the eq 19, the differences between kov and k2 were due to the influence induced by the concentration of [C2NH2MIm][Lys]. An increasing in the temperature caused the expected increase in rate constant (k2). The relationship between the rate constant (k2) and the temperature can be described by the Arrhenius equation (eq 29). ⎛ E ⎞ k 2 = A exp⎜ − act ⎟ ⎝ RT ⎠

Figure 10. Plots of E vs (a) CRNH2 and (b) temperature.

⎛ 3070 ⎞ ⎟ k 2 = 2.785 × 108exp⎜ − ⎝ T ⎠

(30)

4. CONCLUSIONS The multi amine-functionalized ionic liquid [C2NH2MIm][Lys] containing three amino groups exhibited an excellent absorption loading of 1.59 mol CO 2 /mol IL with a concentration of 0.5 mol/L at 313.15 K, which was higher than that of the most existing dual functionalized ILs. It was proven that [C2NH2MIm][Lys] owned a good regenerability, which could maintain 90% of its initial absorption loading after six absorption−desorption cycles. Based on the experiment and 13 C NMR analysis results, the absorption could be divided into two stages, which started with the reaction between CO2 and amino groups to form carbamate and followed by the hydration of CO2 in the solution. Meanwhile, carbamate could also be

(29)

The Arrhenius plot of lnk2 vs 1/T is presented in Figure 11, and the relation can be expressed by eq 30. From this equation, the activation energy (Eact) was found to be 25.5 kJ·mol−1, which was comparable with that of other reactions between CO2 and amines (Table 5). The low Eact indicated that the reaction between CO2 and [C2NH2MIm][Lys] was easy to occur. Moreover, as shown in Figure 12, the predicated and measured k2 were compared, which indicated the good agreement between the experiment results and the predicated values. H

DOI: 10.1021/acs.energyfuels.6b02963 Energy Fuels XXXX, XXX, XXX−XXX

Energy & Fuels

■

Article

AUTHOR INFORMATION

Corresponding Author

*Tel: +86-592-6166216; Fax: +86-592-6162300; E-mail: [email protected]. ORCID

Guohua Jing: 0000-0003-1809-0705 Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was sponsored by the National Natural Science Foundation of China (21576109), and the Natural Science Foundation of Fujian Province (2016J01066 and 2016J05038). We also thank the Instrumental Analysis Center of Huaqiao University for analysis support.

■

Figure 11. Arrhenius plot of CO2 absorption into [C2NH2MIm][Lys] solution.

Table 5. Comparison of the Activation Energy with Literature Values absorbent MEA EEA+DEEA PZ AMP [C2NH2MIm] [Lys]

experimental technique stirred-cell reactor stirred-cell reactor wetted wall column wetted wall column stirred-cell reactor

temperature (K)

Eact (kJ·mol−1)

303−308

94.0

36

298−308

95.5

37

303−313

47.5

38

303−313

25.2

39,40

303−333

25.5

this work

REFERENCES

(1) IPCC. Climate Change 2013 The Physical Science Basis; Cambridge University Press: Cambridge, U.K./New York, 2013. (2) Macdowell, N.; Florin, N.; Buchard, A.; Hallett, J.; Galindo, A.; Jackson, G.; Adjiman, C. S.; Williams, C. K.; Shah, N.; Fennell, P. Energy Environ. Sci. 2010, 3 (11), 1645−1669. (3) Bai, H.; Yeh, A. C. Ind. Eng. Chem. Res. 1997, 36 (6), 2490−2493. (4) Mangalapally, H. P.; Notz, R.; Hoch, S.; Asprion, N.; Sieder, G.; Garcia, H.; Hasse, H. Energy Procedia 2009, 1 (1), 963−970. (5) Mumford, K. A.; Wu, Y.; Smith, K. S.; Stevens, G. W. Front. Chem. Sci. Eng. 2015, 9 (2), 125−141. (6) Zeng, S. J.; Wang, J.; Bai, L.; Wang, B. Q.; Gao, H. S.; Shang, D. W.; Zhang, X. P.; Zhang, S. J. Energy Fuels 2015, 29 (9), 6039−6048. (7) Hu, P. C.; Zhang, R.; Liu, Z. C.; Liu, H. Y.; Xu, C. M.; Meng, X. H.; Liang, M.; Liang, S. S. Energy Fuels 2015, 29 (9), 6019−6204. (8) Lu, J. G.; Lu, Z. Y.; Chen, Y.; Wang, J. T.; Gao, L.; Gao, X.; Tang, Y. Q.; Liu, D. G. Sep. Purif. Technol. 2015, 150, 278−285. (9) Pinto, A. M.; Rodríguez, H.; Arce, A.; Soto, A. Int. J. Greenhouse Gas Control 2013, 18 (5), 296−304. (10) Bates, E. D.; Mayton, R. D.; Ntai, I.; Davis, J. J. R. J. Am. Chem. Soc. 2002, 124 (6), 926−927. (11) Sistla, Y. S.; Khanna, A. J. Ind. Eng. Chem. 2014, 20 (4), 2497− 2509. (12) Mcdonald, J. L.; Sykora, R. E.; Hixon, P.; Mirjafari, A.; Davis, J. H. Environ. Chem. Lett. 2014, 12 (1), 201−208. (13) Zhang, J. Z.; Jia, C.; Dong, H. F.; Wang, J. P.; Zhang, X. P.; Zhang, S. J. Ind. Eng. Chem. Res. 2013, 52 (17), 5835−5841. (14) Hu, H.; Li, F.; Xia, Q.; Li, X. D.; Liao, L.; Fan, M. H. Int. J. Greenhouse Gas Control 2014, 31, 33−40. (15) Crooks, J. E.; Donnellan, J. P. J. Chem. Soc., Perkin Trans. 2 1989, 4 (4), 331−333. (16) Gurkan, B. E.; de la Fuente, J. C.; Mindrup, E. M.; Ficke, L. E.; Goodrich, B. F.; Price, E. A.; Schneider, W. F.; Brennecke, I. F. J. Am. Chem. Soc. 2010, 132 (7), 2116−2117. (17) Guo, H. X.; Zhou, Z. M.; Jing, G. H. Int. J. Greenhouse Gas Control 2013, 16 (4), 197−205. (18) Sánchez, L. M. G.; Meindersma, G. W.; de Haan, A. B. Chem. Eng. J. 2011, 166 (3), 1104−1115. (19) Lv, B. H.; Jing, G. H.; Qian, Y. H.; Zhou, Z. M. Chem. Eng. J. 2016, 289, 212−218. (20) Jing, G. H.; Zhou, L. J.; Zhou, Z. M. Chem. Eng. J. 2012, 181 (1), 85−92. (21) Ratcliff, G. A.; Holdcroft, J. G. Chem. Eng. Sci. 1961, 15 (1−2), 100−110. (22) Schumpe, A. Chem. Eng. Sci. 1993, 48 (1), 153−158. (23) Charpentier, J. C. Adv. Chem. Eng. 1981, 11 (08), 1−133. (24) Weisenberger, S.; Schumpe, A. AIChE J. 1996, 42 (1), 298−300. (25) Zhang, Y. Q.; Zhang, S. J.; Lu, X. M.; Zhou, Q.; Fan, W.; Zhang, X. P. Chem. - Eur. J. 2009, 15 (12), 3003−3011.

reference

Figure 12. Comparison of the predicted and measured k2 at different temperature.

decomposed to form carbonate/bicarbonate as the excess CO2 supplied. CO2 desorption from saturated [C2NH2MIm][Lys] solution was the reverse process of the absorption, in which some carbonate/bicarbonate were heated to release CO2 and some reacted with RNH3+ to form carbamate. After that, the carbamate decomposed to CO2 and [C2NH2MIm][Lys]. The reaction kinetics of CO2 into [C2NH2MIm][Lys] solution was under the pseudo-first-order regime, and the values of kov and k2 were estimated at different concentrations and temperatures. The Arrhenius equation of CO2 absorption was also obtained, and the activation energy was calculated to be 25.5 kJ·mol−1. I

DOI: 10.1021/acs.energyfuels.6b02963 Energy Fuels XXXX, XXX, XXX−XXX

Article

Energy & Fuels (26) Wang, C.; Luo, X.; Luo, H.; Jiang, D. E.; Li, H.; Dai, S. Angew. Chem., Int. Ed. 2011, 50 (21), 4918−4922. (27) Xue, Z. M.; Zhang, Z. F.; Han, J.; Chen, Y.; Mu, T. C. Int. J. Greenhouse Gas Control 2011, 5 (4), 628−633. (28) Zhou, Z. M.; Zhou, X. B.; Jing, G. H.; Lv, B. H. Energy Fuels 2016, 30 (9), 7489−7495. (29) Fu, D.; Zhang, P.; Mi, C. L. Energy 2016, 101, 288−295. (30) Fu, D.; Zhang, P.; Wang, L. M. Energy 2016, 113, 1−8. (31) Zhao, Y. S.; Zhang, X. P.; Zeng, S. J.; Zhou, Q.; Dong, H. F.; Tian, X.; Zhang, S. J. J. Chem. Eng. Data 2010, 55 (9), 3513−3519. (32) Lv, B. H.; Guo, B. S.; Zhou, Z. M.; Jing, G. H. Environ. Sci. Technol. 2015, 49 (17), 10728−10735. (33) Caplow, M. J. J. Am. Chem. Soc. 1968, 90 (24), 6795−6803. (34) Danckwerts, P. V. Chem. Eng. Sci. 1979, 34 (4), 443−446. (35) Portugal, A. F.; Derks, P. W. J.; Versteeg, G. F.; Magalhaes, F. D.; Mendes, A. Chem. Eng. Sci. 2007, 62 (23), 6534−6547. (36) Jing, G. H.; Zhou, L. J.; Zhou, Z. M. Chem. Eng. J. 2012, 181 (1), 85−92. (37) Vaidya, P. D.; Mahajani, V. V. Ind. Eng. Chem. Res. 2005, 44 (6), 1868−1873. (38) Vaidya, P. D.; Kenig, E. Y. Chem. Eng. Technol. 2009, 32 (4), 556−563. (39) Shen, S.; Yang, Y. N.; Bian, Y.; Zhao, Y. Environ. Sci. Technol. 2016, 50, 2054−2063. (40) Sun, W. C.; Yong, C. B.; Li, M. H. Chem. Eng. Sci. 2005, 60 (2), 503−516.

J

DOI: 10.1021/acs.energyfuels.6b02963 Energy Fuels XXXX, XXX, XXX−XXX