Kinetics of carbon dioxide absorption into aqueous [Hmim][Gly] solution

International Journal of Greenhouse Gas Control 16 (2013) 197–205

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Kinetics of carbon dioxide absorption into aqueous [Hmim][Gly] solution Hongxiu Guo, Zuoming Zhou, Guohua Jing ∗ Department of Environmental Science & Engineering, College of Chemical Engineering, Huaqiao University, Xiamen, Fujian 361021, China

a r t i c l e

i n f o

Article history: Received 21 December 2012 Received in revised form 2 March 2013 Accepted 28 March 2013 Available online 4 May 2013 Keywords: Carbon dioxide 1-Hexyl-3-methylimidazolium glycine Absorption Regeneration Kinetics

a b s t r a c t CO2 absorption performance in the [Hmim][Gly] solution was investigated in detail at 298, 303, 313, 323 K within the concentration range of 0.5–1.2 M under an atmospheric pressure by using a double stirred cell absorber with a planar gas–liquid interface. The density, viscosity and pH of the [Hmim][Gly] solution were obtained, and the solubility and diffusivity of CO2 in [Hmim][Gly] solution were calculated. To explore the regeneration property of the CO2 -loaded solution, method under vacuum pressure was used by varying the regeneration temperature from 328 to 343 K and the regeneration time in the range of 30–120 min. The effect of regeneration cycles on regeneration efficiency was also determined. The kinetic results demonstrated that the reaction took place in a fast pseudo-first order regime. Hatta number Ha, the enhancement factor E, the overall reaction kinetic constant kov and the second-order rate constant k2 were obtained using the two-film model. At the absorption temperature of 303 K, the enhancement 1/2 factor E was linear with CIL . The second-order reaction rate was determined as the following equation: k2 = 3.04 × 107 exp(−3050/T ). © 2013 Elsevier Ltd. All rights reserved.

1. Introduction The environmental pollution caused by energy consumption is a globe issue. The use of fossil fuel resources increases greenhouse gas emission and provokes global warming (Chen et al., 2012a). Carbon dioxide (CO2 ), a major greenhouse gas, exerts profound impacts on global climate change and has recently attracted increasing interest in carbon mitigation schemes. Meanwhile, CO2 is also an important carbon source (Chen et al., 2012b). Therefore, many scientists interest in capturing and storing CO2 for various commercial applications. The removal of CO2 from gas streams has traditionally been accomplished by using liquid absorbents such as aqueous amines. This approach has many benefits including maturity, low cost, and high CO2 capacity, but it suffers from several disadvantages, namely the corrosive nature of the aqueous amine, amine degradation, and fugitive losses through solvent evaporation. Consequently, more efficient separation and capture technologies that overcome the limitations of current methods are greatly needed (Mahurin et al., 2012).

∗ Corresponding author at: Jimei Road 668, Xiamen, Fujian 361021, China. Tel.: +86 592 6166216; fax: +86 592 6162345. E-mail addresses: [email protected] (H. Guo), [email protected] (Z. Zhou), [email protected] (G. Jing). 1750-5836/$ – see front matter © 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.ijggc.2013.03.024

The possible use of ionic liquids (ILs) for CO2 capture from power plant flue gases has received remarkable attention in recent years, in both industry and academia. Due to their special characteristics such as negligible vapor pressures, high thermal stability, relatively large electrochemical windows, high absorption capacity for CO2 and easy to be recycled, ILs are considered as a suitable alternative for traditional absorbents (Aparicio et al., 2011; Lee et al., 2007; Zhang et al., 2006). Moreover, ILs have an inherent designer nature in which many of their physicochemical properties can be tuned through appropriate variation in the cation and anion pair comprising the solvent. Properties such as melting point, viscosity, hydrophobicity, hydrophilicity, hydrogen bonding, and coordinating ability can potentially be tailored to preferentially absorb and transport specific gases such as CO2 (Mahurin et al., 2010; Xue et al., 2011). 1-Butyl-3-methylimidazolium hexafluorophosphate [Bmim] [PF6 ] was firstly reported to have a high solubility for carbon dioxide (Blanchard et al., 1999). After that, many scientists turned to study the solubility of CO2 into the imidazolium-based ILs, such as 1,2-dimethyl-3-propylimidazolium bis(trifluoromethylsulfonyl)imide ([Pmmim][Tf2 N]), 1-butyl-3methylimidazolium bis(trifluoromethylsulfonyl)imide ([Bmim] [Tf2 N]), 1-butyl-3-methylimidazolium tetrafluoro-borate ([Bmim] [BF4 ]) (Hou and Baltus, 2007) and 1-ethyl-3-methylimidazolium trifluoromethanesulfonate ([emim][CF3 SO3 ]) (Camper et al., 2005). Then the mechanisms of the high solubility of CO2 in imidazoliumbased ILs were analyzed by experiments and molecular models

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Nomenclature N A P R T V L H* H U VL VL t NA C r z M ¯ M D E kG kL Ha k2 kov n x

absorption rate of CO2 per unit interfacial area (mol m−2 s−1 ) gas–liquid interfacial area (m2 ) pressure (Pa) ideal gas law constant (kmol m−3 s−1 ) temperature (K) CO2 volumetric flow rate (m3 s−1 ) CO2 loading (mol CO2 mol−1 IL) Henry’s constant (atm) The solubility of CO2 in aqueous [Hmim][Gly] solution (mol m−3 Pa−1 ) lattice energy (J mol−1 ) molar volume (cm3 mol−1 ) volume of liquid phase (mL) time (min) Avogadro’s constant concentration (M or mol m−3 ) ionic radius ionic charges molecular weight (g mol−1 ) average molecular weight (g mol−1 ) diffusion coefficient (m2 s−1 ) enhancement factor gas-side physical mass transfer coefficient (mol m−2 s−1 Pa−1 ) liquid-side physical mass transfer coefficient (m s−1 ) Hatta number the second-order rate constant (m3 mol−1 s−1 ) overall kinetic constant (s−1 ) stirrer speed (rpm) mole fraction

Greek letters  regeneration efficiency (%) solubility parameter [(J cm−3 )1/2 ] ı  viscosity of solution (mPa s)  density of the pure [Hmim][Gly] (g mL−1 ) Subscripts c cation anion a 0 initial 1 IL or the first cycle CO2 , the second cycle or the second reaction order 2 3 water n the nth absorption cycle total components t i gas–liquid interface gas phase G L liquid phase the mixture of aqueous [Hmim][Gly] solution m

(Cadena et al., 2004). The results reported that the anion dominated the interactions with CO2 , and the cation played a secondary role. The presence of acidic hydrogen on the imidazolium ring was particularly intriguing as a potential additional mechanism for CO2 solvation. However, gas dissolution in imidazolium-based IL is a physical phenomenon and the absorption rate is quite slow. In order to introduce chemical reaction, a variety of ILs were integrated with ions incorporating functional groups (Lu et al., 2012).

Natural occurring amino acids are considered as alternative for this amino moiety due to their properties of easily obtained, bio-compatible and bio-degradable (Zhou et al., 2012). Amino acid based ILs prepared from 20 natural amino acids were all stable at temperature above 200 ◦ C by thermal gravimetric analysis (Fukumoto et al., 2005). Then three imidazolium amino acid ILs absorbents, 1-butyl-3-methylimidazolium glycine [Bmim][Gly], 1-butyl-3-methylimidazolium alanine [Bmim][Ala], 1-butyl-3-methylimidazolium lysine [Bmim][Lys], were synthesized and used for CO2 absorption. The results revealed that the absorption loading and the absorption rate of CO2 into the amino acid ILs were much larger than that of the amino acid salt absorbents. The regeneration efficiency of the amino acid ILs can reach 100% (Lu et al., 2011). Based on the good properties of the imidazolium amino acid ILs, they are believed to provide an industrially attractive alternative for CO2 capture, and a comprehensive study of CO2 absorption into the imidazolium amino acid ILs has to be undertaken. In this paper, 1-hexyl-3-methylimidazolium glycine ([Hmim][Gly]), chosen [Hmim]+ as cation and glycine as anion (Fig. S1), was studied in detail for CO2 capture. The effects of the [Hmim][Gly] concentration (0.5–1.2 M) and temperature (298–323 K) on CO2 absorption rate were studied. The absorbed CO2 was extruded from the aqueous [Hmim][Gly] solution upon heating under vacuum. The effect of the regeneration temperature, the regeneration time and the absorption/regeneration cycles on the regeneration efficiency was investigated. Moreover, using a gas liquid double stirred-cell absorber, the kinetic experiments were carried out to determine the reaction regime and the reaction rate constant of CO2 absorption in aqueous [Hmim][Gly] solutions. The calculations of the enhancement factor E, the overall reaction kinetic constant kov and the second-order rate constant k2 were realized using the two-film model. 2. Experimental 2.1. Chemicals [Hmim][gly] (97% purity) was obtained from Shanghai Cheng Jie Chemical Co. Ltd., China. Monoethanolamine (MEA) and diethanolamine (DEA), with a purity of 99%, were got from Xilong Chemical Co., Ltd., China. Piperazine (PZ) was supplied by Shaoxing Xingxin chemical Co., Ltd. 2-Amino-2-methyl-1-propanol (AMP) was purchased from J&K Scientific Ltd. Aqueous solutions were prepared with distilled water. The gases were pure CO2 (>99.999% purity) and N2 (>99.999% purity), obtained commercially and used as received. 2.2. Absorption and regeneration methods To explore the absorption kinetics, the effect of the [Hmim][Gly] concentration and temperature on CO2 absorption in the [Hmim][Gly] solution was carried out. A double stirred-cell absorber with a smooth gas–liquid interface was used as the absorption instrument, which was reported in detail in our previous works (Jing et al., 2012). The heated feeding gas of 15% CO2 –85% N2 , which was mixed in a gas mixing tank and calibrated by GC (Shimadzu GC-2010), flowed through a soap bubble meter to measure the flow rate. The absorption rate of CO2 can be calculated as follows: N=

P × (Vout − Vin ) ART

(1)

Here, N denotes CO2 absorption rate (mol m−2 s−1 ). A, P, R, T are gas–liquid interfacial area (m2 ), pressure (Pa), ideal gas law constant (kmol m−3 s−1 ), temperature (K), respectively. Vin and Vout are

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205

volumetric flow rate of inlet, outlet CO2 (m3 s−1 ), respectively. All the absorption rates in this manuscript are instantaneous absorption rates corresponding to a certain time. A bubbling absorber was used for the saturated absorption measurements. The schematic diagram of the apparatus was shown in Fig. S2. The absorption chamber was a glass cylinder. The feeding gas was the same as that used in the double stirred-cell absorber. The temperature of the absorption was controlled by a water baths with a deviation of ±0.2 K. When Vin and Vout were equal, the absorption rate was close to zero. At that time, the CO2 absorption reached equilibrium and the [Hmim][Gly] solution was saturated with CO2 . To measure the optimal regeneration temperature and regeneration time, the absorbed CO2 was extruded from the CO2 loaded [Hmim][Gly] solutions upon the desired temperature under 0.5 kPa vacuum by using a rotary evaporator. All the saturated absorption experiments were conducted with 100 mL of 1.0 M [Hmim][Gly] aqueous solution at 303 K. After absorption, the CO2 loaded solution was regenerated for next cycle of run. To improve the accuracy and the comparability of the results in multiple-cycle experiments, the regenerated [Hmim][Gly] solution was adjusted with some water to 100 mL. CO2 absorption amount was obtained by calculating the integral characteristic value of absorption rate to time, using Origin 8.0 (SR4, 2008, OriginLab company). CO2 loading (L) is the ratio of CO2 absorption amount to the molar of IL. The regeneration efficiency can be calculated as:

2.3. Physicochemical date measurement and calculation 2.3.1. Density, viscosity and pH The density, viscosity and pH of the [Hmim][Gly] aqueous solution were measured by using acidity meter (pHs-25), Anton Paar Density meter DMA 4500 and digit display viscometer (NDJ-5S), respectively.

2.3.2. Solubility of CO2 in aqueous solution of [Hmim][Gly] The solubility of CO2 in pure [Hmim][Gly] was modeled by using regular solution theory (Camper et al., 2005). There was a linear relationship between the natural log of the Henry’s constant H* and the square of the solubility parameters of a gas. For a certain absorbent, H* decreased as the solubility parameter of a gas increased. Regular solution equation may be applied to calculate H* as follows: b(ı1 − ı2 ) T

2

ı1 =

U V1L

M1 1

(5)

M represents the molecular weight (g mol−1 ).  represents the density (g cm−3 ). According to the Kapustinskii (Takamatsu, 1974b) equation: -1

U (J mol-1 ) = 2.40 × 106 (J Å mol )

 ×

1−

1/2 (4)

0.345 Å rc + ra

 zz  c a



rc + ra (6)

c and a represent the cation and anion, respectively. r represents the ionic radius, and z represents the ionic charge. The sum of rc and ra is estimated from the molar volume of the IL using Eq. (7) (Takamatsu, 1974c): V1L = 2NA (rc + ra )3

(7)

NA represents Avogadro’s constant. An empirical formula has been deduced to estimate the solubility parameter for IL as follows (Camper et al., 2005):



−3 1/2

)

= 2.56 × 106 (J/mol)

 ×

1−

z1 z2 (cm3 mol−1 ) (V1L )

0.367 (cm3 mol−1 ) (V1L )

1/3



4/3



(8)

As determined from the heat of vaporization data and actual solubility parameters, Eq. (9) describes how the solubility parameter of CO2 varied with temperature (Camper et al., 2004): ı2 = −0.0535T + 28.26

(9)

According to Eqs. (3)–(9), the H* for CO2 in pure [Hmim][Gly] can be computed. O’Connell and Prausnitz (1964) deduced the thermodynamic of gas solubility in mixed solvents based on regular solution theory. They considered H* in the mixed solvent is the function of the ideality constants in the pure solvents and the nonideality constants in the solute-free solvent mixture. ∗ ∗ ∗ ln H2,m = x1 ln H2,1 + x3 ln H2,3 − ˛13 x1 x3

(10)

x stands for the mole fraction, subscript m stands for the mixture of aqueous [Hmim][Gly] solution, and 3 stands for the other component in the liquid solution. In this paper, 3 represents water. ı3 is reported in the literature (Barton, 1983). ˛ describes the interaction coefficient between the two solvents, and ˛13 is given by a simplified form of Hildebrand’s equation:

(3)

In this equation, 1 represents IL, 2 represents CO2 , ı represents the solubility parameter [(J cm−3 )1/2 ] and H* represents the Henry’s constant (atm). a and b [(J cm−3 )−1 ] are constants, depending only on the gas, as reported in literatures (Camper et al., 2005; Barton, 1983). Since the IL has an unknown energy of vaporization, the solubility parameter for IL is defined by the square root of the lattice energy density (Takamatsu, 1974a):



V1L =

(2)

 represents the regeneration efficiency (%). L represents the CO2 absorption loading (mol CO2 mol−1 IL). 1 represents the first cycle. n represents the nth cycle.

∗ ln H2,1 =a+

U represents the lattice energy (J mol−1 ). VL represents the molar volume (cm3 mol−1 ), which can be calculated from the relations:

ı1 (J cm

Ln = × 100 L1

199

2

˛13 ≈

(ı1 + ı3 ) (V1L + V3L )

(11)

2RT

The solubility of CO2 H can be computed using H* in aqueous [Hmim][Gly] solution according to Henry’s law (Zhong et al., 2001) as follows: H (mol m−3 Pa−1 ) ≈

10 ·  (g cm−3 ) M (g mol−1 )H ∗ (atm)

(12)

¯ denote the average density and average molecular  and M weight in the solution, respectively. Therefore, the solubility of CO2 in the aqueous solvents mixture can be obtained combining Eqs. (3)–(12).

a = −0.0252 b = 8.2672 R2 = 0.9973

a = −0.0242 b = 7.8104 R2 = 0.9975

a = −0.0234 b = 7.4014 R2 = 0.9979

a = −0.0005 b = 0.0437 R2 = 0.9991 0.00 0.00 0.00 0.00 ± ± ± ± 0.82 0.97 1.30 1.68 0.00 0.00 0.00 0.00 ± ± ± ± 2.92 2.58 2.06 1.69 0.00 0.00 0.00 0.00 ± ± ± ± 2.16 1.86 1.58 1.51 2.08E−04 5.03E−05 1.00E−05 1.00E−04 ± ± ± ± 0.03 0.06 0.07 0.05 ± ± ± ± 11.43 11.3 11.10 10.98 298 303 313 323 1.2

1.03097 1.02845 1.02328 1.01757

a = −0.0005 b = 0.0378 R2 = 0.9986 0.00 0.00 0.00 0.00 ± ± ± ± 0.97 1.14 1.51 1.94 0.00 0.00 0.00 0.00 ± ± ± ± 3.06 2.70 2.15 1.77 0.00 0.00 0.00 0.00 ± ± ± ± 1.83 1.58 1.36 1.30 5.57E−05 4.51E−05 4.04E−05 4.51E−05 ± ± ± ± 0.02 0.04 0.06 0.03 ± ± ± ± 11.45 11.27 11.11 10.91 298 303 313 323 1.0

1.02561 1.0234 1.01848 1.0131

a = −0.0005 b = 0.0315 R2 = 0.9985 0.00 0.00 0.00 0.00 ± ± ± ± 1.13 1.32 1.74 2.22 0.00 0.00 0.00 0.00 ± ± ± ± 2.97 2.7 2.09 1.71 0.00 0.00 0.00 0.00 ± ± ± ± 1.57 1.36 1.18 1.22 5.77E−06 1.00E−05 1.53E−05 2.08E−05 ± ± ± ± 0.00 0.07 0.05 0.07 ± ± ± ± 11.45 11.3 11.12 10.97

1.01989 1.0178 1.01319 1.00808

ln m = a(T − 273.15) + b

a = −0.0004 b = 1.1288 R2 = 0.982 0.00 0.00 0.00 0.00 ± ± ± ± 1.43 1.65 2.14 2.69 0.00 0.01 0.00 0.00 ± ± ± ± 3.16 2.79 2.22 1.82

298 303 313 323

3.2.1. Comparison with the conventional absorbents Fig. 1 compares the CO2 absorption rate in the absorbent of 1.0 M [Hmim][Gly], MEA, DEA and AMP solution. The results showed that

0.8

3.2. Absorption of CO2 into aqueous [Hmim][Gly] solutions

0.01 0.03 0.03 0.02

the thermal expansion coefficient () was the negative value of the straight-line slope, so that  = 0.0006 K−1 . The solubility of CO2 in the pure [Hmim][Gly] was 57.13 atm at 303 K under atmospheric pressure, and the values of [Bmim][Tf2 N], [Bmim][BF4 ] were 42 and 63 atm, respectively (Hou and Baltus, 2007). The physicochemical data of aqueous [Hmim][Gly] solution, against the [Hmim][Gly] concentration at various temperatures, were presented in Table 1. The results showed that the density, viscosity and pH of the aqueous [Hmim][Gly] solution increased as the concentration of [Hmim][Gly] increased, but decreased as the temperature increased. For a certain concentration, if ln  was graphed out with T, a linear relationship can be obtained at a temperature range from 298 to 323 K. The results were in good agreement with the references (Mota-martinez et al., 2012). At 303 K, the viscosities of the CO2 -loaded MEA, DEA and [Hmim][Gly] solutions with a concentration of 1.0 M were 0.92, 1.10, 1.59 mPa s, respectively. It was possible to gain some insight into the viscosities that there were small differences. At some sort of fixed [Hmim][Gly] concentration, the diffusivity of CO2 in aqueous [Hmim][Gly] solution increased while the solubility of CO2 in aqueous [Hmim][Gly] solution decreased with the increase of temperature. Moreover, in a constant temperature, the solubility of CO2 decreased significantly as the concentration of [Hmim][Gly] increased, which was in accord with “salting-out” effect (Weisenberger and Schumpe, 1996) as described in our previous work (Jing et al., 2012). The diffusivity of CO2 in aqueous [Hmim][Gly] solution also decreased with the increase of [Hmim][Gly] concentration.

0.00 0.00 0.00 0.00

(15) p

± ± ± ±



± ± ± ±

=− p

∂ ln  ∂T

11.42 11.28 11.08 10.88



298 303 313 323



0.5

∂V ∂T

m (mPa s)



Table 1 Physicochemical data of [Hmim][Gly] solution for CO2 capture.

1 = V

H2,m (×10−4 mol m−3 Pa−1 )

The densities of the pure [Hmim][Gly] at 298, 303, 313, 323 K were 1.091, 1.088, 1.082, 1.075 g cm−3 , respectively. Comparing the densities with that of 1-propyl-3-methylimidazolium glycine ([Pmim][Gly]) and [Bmim][Gly] reported in the literature (Tong et al., 2012), it showed that the IL with a longer alkyl chain exhibited a smaller density and a larger ‘free volume’. An approximately linear relationship was found between ln  and (T-273.15). According to the following definition (Tong et al., 2012)

1.24 1.09 0.96 0.94

3.1. The physicochemical properties of the aqueous [Hmim][Gly] solution

1.15E−05 1.15E−04 1.27E−04 1.22E−04

D2,m (×10−9 m2 s−1 )

3. Results and discussion

± ± ± ±

Here, D denotes the diffusion coefficient (m2 s−1 ).  denotes viscosity (mPa s). According to the values of 3 at different temperatures reported in the literature (Korson et al., 1969), the diffusion coefficient of CO2 in [Hmim][Gly] solutions can be calculated, according to Eqs. (13) and (14).

1.01036 1.00962 1.00547 1.00076

(14)

m (g cm−3 )

712.5 2.591 × 105 − T T2

pH

log D2,3 = −8.1764 +

(13)

T (K)

0.8 (D2,m 0.8 m )T = (D2,3 3 )T = const

ln m = aT + b

2.3.3. Diffusivity of CO2 in the aqueous solution of [Hmim][Gly] Since the [Hmim][Gly] solutions are electrolyte solutions, the diffusivity of CO2 in aqueous electrolyte solutions is calculated from the relations (Barrett, 1966; Danckwerts, 1970):

a = −0.0218 b = 6.7 R2 = 0.9976

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205

CIL,0 (M)

200

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205

201

(a) 2.1

0.5 M 0.8 M 1.0 M 1.2 M

-2

1.8

-3

1.5

Absorption rate ( 10-3mol.m-2.s-1)

AMP

-1

Absorption rate ( 10 mol.m .s )

2.1 DEA MEA IL

1.2

0.9 0.6 0.3

0

30

60

90

120

150

180

1.8

1.5

1.2

0.9 0.6

0.3

210

0

30

60

Absorption time (min)

3.2.2. Effect of [Hmim][Gly] concentration on the absorption The absorption rates of CO2 in 0.5–1.2 M [Hmim][Gly] aqueous solution at 303 K were illustrated in Fig. 2(a). The results showed that the CO2 absorption rate decreased as the absorption time increased, especially in the first 30 min. Noticeably, with a further increase in the absorption time, the tendency of the decrease changed slowly. Furthermore, at the same absorption time, the CO2 absorption rate increased as the concentration increased. But the concentration is higher; the increase tendency of the absorption rate is slower. This phenomenon could be explained by the high viscosity caused by the increase of [Hmim][Gly] concentration, which hindered the diffusion of CO2 (Zhang et al., 2012). Fig. 2(b) shows that the absorption amount increased as the absorption time and the [Hmim][Gly] concentration increased. The CO2 loading into the [Hmim][Gly] aqueous solution was shown in Fig. 2(c). The CO2 loading decreased as the concentration increased, and the tendency was contrary to that of the absorption amount. Therefore, increase [Hmim][Gly] concentration would have a conflicting influence on CO2 loading and diffusion. On the other hand, a larger concentration was favor to increasing the absorption rate and the absorption amount. So, it is necessary to choose an appropriate concentration for CO2 capture. The concentration of 1.0 M was considered to be the optimal within the range of this study. 3.2.3. Effect of temperature on the absorption The absorption rates of CO2 into aqueous [Hmim][Gly] solution were measured in different temperature, and the results were presented in Fig. 3. Obviously, in the investigated range, temperature had little influence on the absorption rate. At the same detection time, the absorption rate of CO2 increased slightly as the temperature increased. The appearance was similar to other absorbents, just like potassium glycinate (Portugal et al., 2007) and [N1111 ][Gly]

120

150

180

210

180

210

180

210

0.07

0.5 M 0.8 M 1.0 M 1.2 M

Absorption amount (mol)

0.06 0.05 0.04 0.03 0.02 0.01 0.00

0

30

60

90

120

150

Absorption time (min)

(c)

0.40

0.5 M 0.8 M 1.0 M 1.2 M

0.35 -1

the absorption rates of CO2 in all absorbents were quite fast at the beginning. After 20 min, the order of the CO2 absorption rate was [Hmim][Gly] > MEA > DEA > AMP. [Hmim][Gly] had the fastest absorption rate, which can be explained by the following reasons. Since the reaction between [Hmim][Gly] and CO2 was similar to the reaction between MEA and CO2 , the primary amine reacted with CO2 , and created COO− . Then, the COO− resided in the [Gly]− , are hydrolyzed to produce COOH and OH− . The produced OH− took part in the CO2 absorption, which assured that [Hmim][Gly] was able to absorb CO2 at a considerable rate.

(b)

CO2-loading (mol CO2.mol IL)

Fig. 1. Absorption rate of CO2 into different aqueous absorbents (absorption temperature: 303 K; gas flow: 120 mL min−1 ; CO2 concentration: 15% (v/v); solution volume: 200 mL; rG : 250 rpm; rL : 130 rpm; data were the mean values of triplicate experiments).

90

Absorption time (min)

0.30 0.25 0.20 0.15 0.10 0.05 0.00

0

30

60

90

120

150

Absorption time (min) Fig. 2. Absorption of CO2 into aqueous [Hmim][Gly] solution with different concentrations. (a) Absorption rate; (b) absorption amount and (c) CO2 loading (absorption temperature: 303 K; gas flow: 120 mL min−1 ; CO2 concentration: 15% (v/v); solution volume: 200 mL; rG : 250 rpm; rL : 130 rpm; data were the mean values of triplicate experiments).

202

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205

0.5

298 K 303 K 313 K 323 K

1.5 1.2 0.9 0.6 0.3 0.0

0

30

60

90

120

150

180

0.4

-1

1.8

CO2 loading (mol CO2.mol IL)

Absorption rate (10 -3 mol.m -2 .s -1 )

2.1

210

Absorption time (min)

0.3

0.2

Initial solution 120 min 90 min 60 min 30min

0.1

0.0

0

20

40

60

80

100

Absorption time (min) Fig. 3. Absorption rate of CO2 into aqueous [Hmim][Gly] at different temperature (IL concentration: 1.0 M; gas flow: 120 mL min−1 ; CO2 concentration: 15% (v/v); solution volume: 200 mL; rG : 250 rpm; rL : 130 rpm; data were the mean values of triplicate experiments).

(Jing et al., 2012). To compare the results with that reported by other researchers under the same condition, 303 K was chosen as the absorption temperature in this paper. 3.3. Regeneration of CO2 saturated [Hmim][Gly] solutions 3.3.1. Effect of regeneration temperature on regeneration efficiency To explore the optimal regeneration temperature, the regeneration experiments were run at 323, 328, 333, 338 and 343 K for 90 min, respectively. Fig. 4 shows that, no matter the initial [Hmim][Gly] solution or the regenerated barren solution was used, all the absorptions were almost completed within 60 min. This rapid absorption was related to their low viscosities. The CO2 loading of the initial solution was up to 0.45 mol CO2 mol−1 IL. The CO2 loading of the barren solution regenerated at 343 K in the second cycle was very close to that of the initial solution, and decreased slightly with the decrease of the regeneration temperature. The regeneration efficiency was 86.63%, 92.31%, 94.64%, 96.64% and 99.11% as the regeneration temperature was 323, 328, 333, 338

0.4

-1

CO2 loading (mol CO2,mol IL)

0.5

0.3

Initial solution 343 K 338 K 333 K 328 K 323 K

0.2

0.1

0.0

0

20

40

60

80

100

Absorption time (min) Fig. 4. The CO2 loading into the initial [Hmim][Gly] solution and the regenerated barren solution at vary regeneration temperature (absorption temperature: 303 K; gas flow: 120 mL min−1 ; CO2 concentration: 15% (v/v); solution volume: 100 mL; regeneration time: 90 min).

Fig. 5. The CO2 loading into the initial [Hmim][Gly] solution and the regenerated barren solution at vary regeneration time (absorption temperature: 303 K; gas flow: 120 mL min−1 ; CO2 concentration: 15% (v/v); solution volume: 100 mL; regeneration temperature: 338 K).

and 343 K, respectively. At lower temperatures, this increase tendency was more obviously. However, the increase rose very slightly above 338 K. Therefore, 338 K was the most appropriate regeneration temperature. 3.3.2. Effect of regeneration time on regeneration efficiency According to the above experience, the effect of regeneration time aggrandized from 30 to 120 min under 338 K was studied. As shown in Fig. 5, the CO2 loading in the second absorption cycle was only 0.34 mol CO2 mol−1 IL for the barren solution regenerated with 30 min, and the value was up to 0.44 mol CO2 mol−1 IL when the regeneration time was increased to 120 min. The regeneration efficiency was exceedingly sensitive to the changes in regeneration time. When the regeneration time was 30, 60, 90 and 120 min, the regeneration efficiency was 74.92%, 91.97%, 96.64% and 98.33%, respectively. It was of interest to note that there was 17.06% increase in the regeneration efficiency between the 30 and the 60 min. While increase the time from 90 to 120 min, only led to a 1.69% increase in regeneration efficiency. As a result, the optimum regeneration time was 90 min. 3.3.3. Effect of absorption/regeneration cycles on regeneration efficiency To investigate the effect of absorption/regeneration cycles on regeneration efficiency, the experiments were run for four cycles with a regeneration temperature of 338 K and a regeneration time of 90 min. Fig. 6 shows that the CO2 absorption loadings changed slightly as the cycles increased, which revealed that [Hmim][Gly] solution for CO2 capture can be recycled with high CO2 absorption loading and rapid absorption rate. The CO2 loading of the first absorption cycle was 0.45 mol CO2 mol−1 IL, however, the value of the fourth absorption cycle was 0.40 mol CO2 mol−1 IL. The regeneration efficiency in turn was 96.7%, 95.0%, 88.8% for the first, the second, and the third cycle of regeneration, respectively. The results were inconsistent with that of the other imidazolium-based ILs from literatures (Zhang et al., 2006; Lu et al., 2011), in which they reported that the recovered ILs recycled for CO2 uptake (four cycles) with no observed loss of CO2 absorption rate and absorption loading. One reason was that the reaction of [Hmim][Gly] with CO2 produced a number of heat-stable salts, which were unregenerate under the regeneration conditions (Winyu et al., 2006).

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205

203

The first cycle The third cycle

0.5

K,k2

The second cycle The forth cycle

H2 N − CH2 − COO− R+ + CO2  − OOC+ H2 N − CH2 − COO− R+ (16)

-1

CO2 loading ( mol CO2.mol IL)

0.6

0.4 −

OOC+ H2 N − CH2 − COO− R+ + Bi 

−

OOCHN − CH2 − COO− R+ + H2 O 

−

0.3

OOCHN − CH2 − COO− R+ + Bi H+

(17)

H2 N − CH2 − COO− R+ + HCO− 3

0.2 +

0.1

H3 N − CH2 − COOR+  H2 N − CH2 − COO− R+ + H+

CO2 + H2 O  H 0.0 0

30

60

90

120

150

180

210

240

HCO− 3

H

4. Absorption kinetics The kinetics region of [Hmim][Gly] aqueous solution absorbing CO2 was determined according to the judge method (Levenspiel and Godfrey, 1974). The absorption rates of CO2 into [Hmim][Gly] solution were measured with a stirring speed for gas phase of 250 rpm at 303 K. The results listed in Table 2 show that liquid-phase volumes and liquid stirring speed had no effect on absorption rates of CO2 , while the absorption rates increased as the [Hmim][Gly] solution concentration increased. The changes of the liquid stirring speed reflected the changes of the liquid-phase mass transfer coefficient. Hence, the kinetics region of absorption CO2 into the aqueous was the fast pseudo-first order reaction regime. Since the reaction kinetics can be well described using the zwitterion mechanism proposed originally (Caplow, 1968), the equilibrium reactions in the liquid phase were suggested as follows:

E=

0 3 5 10 15 20 25 30 40 50 60

1.96 1.73 1.62 1.66 1.58 1.50 1.43 1.38 1.34 1.31 1.27

± ± ± ± ± ± ± ± ± ± ±

0.01 0.00 0.02 0.05 0.01 0.00 0.00 0.02 0.00 0.02 0.00

1.97 1.73 1.62 1.66 1.58 1.49 1.44 1.37 1.34 1.30 1.27

± ± ± ± ± ± ± ± ± ± ±

0.00 0.01 0.01 0.01 0.02 0.06 0.01 0.04 0.00 0.01 0.05

VL = 200 mL nL = 180 rpm 1.96 1.74 1.63 1.66 1.58 1.49 1.43 1.37 1.34 1.30 1.27

(22)

N kL .C2,i

(23)

According to the phase balance in the gas–liquid interface, the concentration of CO2 at interface, C2,i , can be expressed as follows:



C2,i = H2 P2 −

N kG



(24)

Here, kG and kL are the gas-side and liquid-side mass transfer coefficients, respectively. The values of which have been calculated in our previous work (Jing et al., 2012). P2 was equaled to Pout , the partial pressure of outlet CO2 . where Ha is:



D2,m kov

(25)

kL

N = C2,i VL = 200 mL nL = 130 rpm

(21)

−

Moreover, the CO2 absorption rate is also given by Eq. (26)

N (10−3 mol m−2 s−1 ) VL = 100 mL nL = 130 rpm

(20)

CO2 reacted with the [Hmim][Gly] via the formation of a zwitterion, subsequently deprotonated by a base present in solution, where Bi were the bases present in the solution able to deprotonate the zwitterions, including H2 O, OH− and [Hmim][Gly], and the last one played a leading role here. According to the two-film model described in detail in our previous work (Jing et al., 2012), for the fast pseudo-first order reaction regime, the Hatta number Ha can be considered equal to the enhancement factors E.

Ha = Table 2 Effects of liquid-phase volume and liquid stirring speed on absorption rate.

(19)

+ HCO2− 3

+ CO2− 3

H2 O  H + OH

Fig. 6. Four consecutive cycles of CO2 absorption (absorption temperature: 303 K; gas flow: 120 mL min−1 ; CO2 concentration: 15% (v/v); solution volume: 100 mL; regeneration temperature: 338 K; regeneration time: 90 min).

t (min)

+

+

+

Absorption time (min)

(18)

± ± ± ± ± ± ± ± ± ± ±

0.01 0.00 0.03 0.00 0.01 0.00 0.00 0.02 0.00 0.01 0.02



k2 D2,m C1

(26)

Therefore, E, Ha, the overall reaction kinetic constant kov and the second-order rate constant k2 , of [Hmim][Gly] aqueous solution for a concentration range from 0.5 to 1.2 M at 303 K, were calculated and presented in Table 3. The results indicated that Ha > 2 and C1,0  C2,i , which was satisfied with the criteria for the fast pseudofirst order reaction regime. As the concentration of [Hmim][Gly] increased, E increased and varied similarly with the change of Ha, kov and k2 . Comparisons of E between the experimental values and the calculated line by Eq. (23) were shown in Fig. 7. The 1/2 results showed that E was linear with CIL . The error bars for the experimental values seemed to indicate acceptable tolerance, even

Table 3 The kinetic data for CO2 absorption in [Hmim][Gly] aqueous solution at 303 K. CIL,0 (M)

N (×10−3 mol m−2 s−1 )

0.5 0.8 1.0 1.2

1.47 1.65 1.73 1.81

± ± ± ±

0.01 0.01 0.00 0.00

C2 ,i (mol m−3 )

P2 (Pa) 6862.08 5896.7 5422.65 5029.78

± ± ± ±

43.88 42.99 21.22 21.09

1.86 1.53 1.37 1.23

± ± ± ±

0.01 0.01 0.01 0.01

E 36.02 48.86 57.50 66.75

Kov (s−1 )

Ha ± ± ± ±

0.44 0.61 0.37 0.44

36.02 48.86 57.50 66.75

± ± ± ±

0.44 0.61 0.37 0.44

379.61 873.45 1402.38 2227.18

k2 (L mol−1 s−1 ) ± ± ± ±

9.00 20.86 19.20 31.18

759.22 1091.81 1402.38 1855.99

± ± ± ±

18.01 26.08 19.20 25.99

204

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205

Table 4 The kinetic data for CO2 absorption in 1 M [Hmim][Gly] aqueous solution at different temperature. T (K)

N (×10−3 mol m−2 s−1 )

298 303 313 323

1.68 1.73 1.77 1.79

± ± ± ±

C2,i (mol.0 m−3 )

P2 (Pa)

0.00 0.01 0.00 0.01

5794.29 5647.41 5426.96 5161.11

± ± ± ±

0.00 36.86 21.21 42.29

1.65 1.42 1.09 0.85

± ± ± ±

0.00 0.01 0.00 0.01

considering the fact that the y-axis of this figure is on a logarithmic scale. Additionally, Table 4 shows the values of E, Ha, kov and k2 of 1.0 M [Hmim][Gly] aqueous solution absorbing CO2 at 298–303 K. It was observed that k2 increased as the temperature increased. Fig. 8 shows the natural log of k2 versus the difference of the reciprocal of the temperatures, and a linear relationship was obtained. Therefore, the kinetic data can be calculated from an Arrhenius plot of ln k2 ∼ 1/T, represented by k2 = 3.04 × 107 exp

 −3050 

(27)

T

The activation energy for k2 was calculated to be 25.36 kJ mol−1 from Eq. (27). This value was lower than those reported in literatures. For the reaction between CO2 and aqueous MEA, the value 70

y=-19.83+78.12x 2 R =0.9901

65 60

E

55 50 45 40 35 0.7

0.8

0.9

1.0

1.1

CIL

1/2

Fig. 7. E vs CIL .

± ± ± ±

0.01 0.38 0.25 0.62

50.38 55.36 63.32 71.52

± ± ± ±

0.01 0.38 0.25 0.62

1066.60 1300.28 1752.7 2300.21

k2 (L mol−1 s−1 ) ± ± ± ±

1.13 18.13 12.49 38.94

1066.60 1300.28 1752.77 2300.21

± ± ± ±

1.13 18.13 12.49 38.94

was 45 kJ mol−1 (Versteeg et al., 1996). Blauwhoff et al. (1983) summarized the values for the reaction of CO2 and DEA was about 42.0 kJ mol−1 . Lee and Paul (2007) found a value of 63.8 kJ mol−1 for sodium glycinate absorbing CO2 . Simons et al. (2010) obtained the activation energy of 26.0 kJ mol−1 for CO2 absorbed in the sarcosine salt solutions. The calculated k2 were considerably smaller about 5, than the reaction constant reported in literature for the reaction of CO2 with aqueous MEA (Alper, 1990). However, the same magnitude of the reaction constant indicated that the reactivity of [Hmim][Gly] with CO2 was comparable or could be even higher than that of MEA with CO2 (Galán Sánchez et al., 2011). 5. Conclusions Density, viscosity of [Hmim][Gly] aqueous solutions were small, and increased as [Hmim][Gly] concentration increased ranging from 0.5 to 1.2 M, but decreased as the temperature increased from 298 to 323 K. The solubility of CO2 in [Hmim][Gly] aqueous solution decreased while the diffusion coefficient increased with the increase of temperature. The solubility increased as the concentration increased, and also, likewise with the diffusion coefficient. Compared to MEA, DEA, AMP, the absorption rate of CO2 into [Hmim][Gly] solution was the fastest, and increased as the concentration increased. Temperature had a slightly influence on the absorption rate between 298 and 323 K. The regeneration efficiency of the CO2 -loaded solution was high, and increased with the increase of the regeneration temperature between 323 and 343 K and the regeneration time from 30 to 120 min under 0.095 MPa. On the contrary, it decreased slightly with the increase of regeneration cycles. The kinetics region of absorption CO2 into aqueous [Hmim][Gly] was the fast pseudo first order reaction regime. At the absorption temperature of 303 K, the enhancement factor E was 1/2 linear with CIL . The second-order reaction rate was determined as

Acknowledgments The authors are grateful for the financial support from the National Natural Science Foundation of China (No. 21277053) and Fujian Province (No. 2011J01052), and the Program for New Century Excellent Talents in the University of China (NCET-11-0851).

7.8

lnk2

50.38 55.36 63.32 71.52

Kov (s−1 )

Ha

the following equation: k2 = 3.04 × 107 exp(−3050/T ).

1/2

y=17.23-3.05x 2 R =0.9985

7.6

E

Appendix A. Supplementary data

7.4

Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.ijggc. 2013.03.024.

7.2

References 7.0 3.10

3.15

3.20

3.25

3.30

3.35

3.40

-1

1000/T (K ) Fig. 8. Arrhenius plot of the second-order rate constant k2 for CO2 -[Hmim][Gly] reaction as a function of temperature.

Alper, E., 1990. Reaction mechanism and kinetics of aqueous solutions of 2-amino2-methyl-1-propanol and carbon dioxide. Industrial and Engineering Chemistry Research 29, 1725–1728. Study on Aparicio, S., Atilhan, M., Khraisheh, M., Alcalde, R., 2011. hydroxylammonium-based ionic liquids. I. Characterization. Journal of Physical Chemistry B 115, 12473–12486. Barrett, P.V.L., 1966. Gas absorption on a sieve plate, Ph.D. Thesis. University of Cambridge.

H. Guo et al. / International Journal of Greenhouse Gas Control 16 (2013) 197–205 Barton, A.F.M., 1983. Handbook of Solubility Parameters and Other Cohesion Parameters. CRC Press, Boca Raton, FL. Blanchard, L.A., Haneu, D., Beckman, E.J., Brennecke, J.F., 1999. Green processing using ionic liquids and CO2 . Nature 399, 28–29. Blauwhoff, P.M.M., Versteeg, G.F., Van Swaaij, W.P.M., 1983. A study on the reaction between CO2 and alkanolamines in aqueous solutions. Chemical Engineering Science 38, 1411–1429. Cadena, C., Anthony, J.L., Shah, J.K., Morrow, T.I., 2004. Why is CO2 so soluble in imidazolium-based ionic liquids? Journal of the American Chemical Society 126, 5300–5308. Camper, D., Scovazzo, P., Koval, C., Noble, R., 2004. Gas solubilities in roomtemperature ionic liquids. Industrial and Engineering Chemistry Research 43, 3049–3054. Camper, D., Becker, C., Koval, C., Noble, R., 2005. Low pressure hydrocarbon solubility in room temperature ionic liquids containing imidazolium rings interpreted using regular solution theory. Industrial and Engineering Chemistry Research 44, 1928–1933. Caplow, M., 1968. Kinetics of carbamate formation and breakdown. Journal of the American Chemical Society 90, 6795–6803. Chen, H.S., Dou, B.L., Song, Y.C., 2012a. Studies on absorption and regeneration for CO2 capture by aqueous ammonia. International Journal of Greenhouse Gas Control 6, 171–178. Chen, J.J., Li, W.W., Li, X.L., Yu, H.Q., 2012b. Carbon dioxide capture by aminoalkyl imidazolium-based ionic liquid: a computational investigation. Physical Chemistry Chemical Physics 14, 4589–4596. Danckwerts, P.V., 1970. Gas–Liquid Reactions. McGraw-Hill, New York. Fukumoto, K., Yoshizawa, M., Ohno, H., 2005. Room temperature ionic liquids from 20 natural amino acids. Journal of the American Chemical Society 127, 2398–2399. Galán Sánchez, L.M., Meindersma, G.M., de Haan, A.B., 2011. Kinetics of absorption of CO2 in amino-functionalized ionic liquids. Chemical Engineering Journal 166, 1104–1115. Hou, Y., Baltus, R.E., 2007. Experimental measurement of the solubility and diffusivity of CO2 in room-temperature ionic liquids using a transient thin-liquid-film method. Industrial and Engineering Chemistry Research 46, 8166–8175. Jing, G.H., Zhou, L.J., Zhou, Z.M., 2012. Characterization and kinetics of carbon dioxide absorption into aqueous tetramethylammonium glycinate solution. Chemical Engineering Journal 181–182, 85–92. Korson, L., Drost-Hansen, W., Millero, F.J., 1969. Viscosity of water at various temperatures. Physical Chemistry 73, 34–39. Lee, F., Paul, S., 2007. Solubility, diffusivity, and permeability of gases in phosphonium-based room temperature ionic liquids: data and correlations. Industrial and Engineering Chemistry Research 46, 1369–1374. Lee, S., Song, H.J., Maken, S., Park, J.W., 2007. Kinetics of CO2 absorption in aqueous sodium glycinate solutions. Industrial and Engineering Chemistry Research 46, 1578. Levenspiel, O., Godfrey, J.M., 1974. A gradientless contactor for experimental study of interface mass transfer with/without reaction. Chemical Engineering Science 29, 1723–1730. Lu, B.H., Jin, J.J., Zhang, L., Li, W., 2012. Absorption of carbon dioxide into aqueous blend of monoethanolamine and 1-butyl-3-methylimidazolium tetrafluoroborate. International Journal of Greenhouse Gas Control 11, 152–157.

205

Lu, J.G., Ji, Y., Zhang, H., Suo, F.X., Fan, F., Liu, C., 2011. A type of ionic liquids for CO2 capture and the synthesis method. Patent 102008870, China. Mahurin, S.M., Lee, J.S., Baker, G.A., Luo, H., Dai, S., 2010. Performance of nitrilecontaining anions in task-specific ionic liquids for improved CO2 /N2 separation. Journal of Membrane Science 353, 177–183. Mahurin, S.M., Yearya, J.S., Bakerb, S.N., Jianga, D., Daia, S., Baker, G.A., 2012. Ringopened heterocycles: promising ionic liquids for gas separation and capture. Journal of Membrane Science 401–402, 61–67. Mota-martinez, M.T., Althuluth, M., Kroon, M.C., Peters, C.J., 2012. Solubility of carbon dioxide in the low-viscosity ionic liquid 1-hexyl-3-methylimidazolium tetracyanoborate. Fluid Phase Equilibria 332, 35–39. O’Connell, J.P., Prausnitz, J.M., 1964. Thermodynamics of gas solubility in mixed solvents. I&EC Fundamentals 3, 347–364. Portugal, A.F., Derks, P.W.J., Versteeg, G.F., Magalhães, F.D., Mendes, A., 2007. Characterization of potassium glycinate for carbon dioxide. Chemical Engineering Science 62, 6534–6547. Simons, K., Brilman, D.W.F., Mengers, H., Nijmeijer, K., Wessling, M., 2010. Kinetics of CO2 absorption in aqueous sarcosine salt solutions: influence of concentration, temperature, and CO2 loading. Industrial and Engineering Chemistry Research 49, 9693–9702. Takamatsu, T., 1974a. The solubility study of ion-pairs in organic solvents. Bulletin of the Chemical Society of Japan 47, 2647. Takamatsu, T., 1974b. The application of the regular solution theory to the ion-pair systems. Bulletin of the Chemical Society of Japan 47, 1287. Takamatsu, T., 1974c. Solubility studies of ion-pairs in organic solvents. Bulletin of the Chemical Society of Japan 47, 1285. Tong, J., Hong, M., Chen, Y., Wang, H., Guan, W., Yang, J.Z., 2012. The surface tension, density and refractive index of amino acid ionic liquids: [C3 mim][Gly] and [C4 mim][Gly]. Journal of Chemical Thermodynamics 54, 352–357. Versteeg, G.F., Van Dijck, L.A.J., Van Swaaij, W.P.M., 1996. On the kinetics between CO2 and alkanolamines both in aqueous and nonaqueous solutions. An overview. Chemical Engineering Communications 144, 133. Weisenberger, S., Schumpe, A., 1996. Estimation of gas solubilities in salt solutions at temperatures from 273 to 363 K. AIChE Journal 42, 298–300. Winyu, T., Amornvadee, V., Bryce, M., 2006. Electrochemical investigation on the effect of heat-stable salts on corrosion in CO2 capture plants using aqueous solution of MEA. Industrial and Engineering Chemistry Research 45, 2586–2593. Xue, Z.M., Zhang, Z.F., Han, J., Chen, Y., Mu, T.C., 2011. Carbon dioxide capture by a dual amino ionic liquid with amino-functionalized imidazolium cation and taurine anion. International Journal of Greenhouse Gas Control 5, 628–633. Zhang, F., Ma, J.W., Zhou, Z., Wu, Y.T., Zhang, Z.B., 2012. Study on the absorption of carbon dioxide in high concentrated MDEA and ILs solutions. Chemical Engineering Journal 181–182, 222–228. Zhang, J.M., Zhang, S.J., Dong, K., Zhang, Y.Q., Shen, Y.Q., Lv, X.G., 2006. Supported absorption of CO2 by tetrabutylphosphonium amino acid ionic liquids. Chemistry – A European Journal 12, 4021–4026. Zhong, Q., Wang, J., Chen, Q.Q., Qu, H.X., 2001. Chemical Principle. National Defence Industry Press, Beijing. Zhou, Z.M., Jing, G.H., Zhou, L.J., 2012. Characterization and absorption of carbon dioxide into aqueous solution of amino acid ionic liquid [N1111 ][Gly] and 2-amino-2-methyl-1-propanol. Chemical Engineering Journal 204–206, 235–324.