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Characterization and kinetics of CO2 absorption into aqueous 1-ethyl-3-methylimidazolium glutamate solution
Journal of Environmental Chemical Engineering 4 (2016) 1137–1147
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Characterization and kinetics of CO2 absorption into aqueous 1-ethyl -3-methylimidazolium glutamate solution Jie Ma, Zaikun Wu, Wenkang Lei, Ping Yu, Yunbai Luo* College of Chemistry and Molecular Sciences, Wuhan University, Wuhan 430072, China
A R T I C L E I N F O
A B S T R A C T
Article history: Received 26 August 2015 Received in revised form 5 December 2015 Accepted 12 January 2016 Available online 14 January 2016
The kinetic of carbon dioxide (CO2) into aqueous 1-ethyl-3-methylimidazolium glutamate ([Emim]2[Glu]) solution was investigated by a double stirred-cell absorber. The Density and viscosity of [Emim]2[Glu] solution were measured with concentrations ranging from 0.3 to 1.5 mol L1 at temperatures from 298.15 to 323.15 K. The diffusivity of CO2 in [Emim]2[Glu] solution were estimated at the same temperatures and concentrations. The effects of the [Emim]2[Glu] concentration, absorption temperature and CO2 concentration were studied. The results showed that absorption CO2 into aqueous [Emim]2[Glu] solution was the fast pseudo-first order reaction regime. The enhancement factor E, the overall reaction kinetic constant kOV and the second-order rate constant k2 were calculated by the two-
Keywords: Kinetic 1-Ethyl-3-methylimidazolium glutamate Physicochemical data Carbon dioxide Absorption
1=2
film theory. The result showed that E was linear with C ILs , and the second-order rate constant could be expressed as following equation: k2 ¼ 9:253 10 expð3665=T Þ. ã 2016 Elsevier Ltd. All rights reserved.
1. Introduction Carbon dioxide, the main greenhouse gas in the atmosphere, causes global warming. There is no doubt that fossil fuels and petroleum industries will still act as an important role for a long time, which makes CO2 capture an essential process to our society. Aqueous alkanolamine solutions, such as monoethanolamine (MEA), diethanolamine (DEA) and N-methyldiethanolamine (MDEA) [1], is the most widespread absorbent for the removal of CO2 in flue gas because of the advantage of high reactivity and low solvent cost. However, aqueous alkanolamine solutions have several serious drawbacks: solvent loss, corrosion, regeneration fee and oxidative degradation [2–4]. Recently, ionic liquids with low vapor pressure, high thermal stability, low melting point, low energy consumption and nontoxic, have been considered as new absorbents for CO2 removal [5,6]. As many literatures reported [7–9], CO2 had been demonstrated to have remarkable solubility in aqueous imidazolium-based ionic liquids solution. However, carbon dioxide dissolution in imidazolium-based ILs not only are a physical process but also have low absorption rate. In view of these reasons, a multitude kinds of ionic liquids which are amino-functionalized were regarded as a kind of novel absorbent for CO2 capture due to their large absorbing capacity, good physical and chemical properties at atmospheric
* Corresponding author. Fax: +86 2768752511. E-mail address: [email protected] (Y. Luo). http://dx.doi.org/10.1016/j.jece.2016.01.016 2213-3437/ ã 2016 Elsevier Ltd. All rights reserved.
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pressure had been reported. Bates et al. [10] reported that the molar uptake of CO2 per mole of [NH2p-bim][BF4] was closed to the theoretical maximum loading by MEA during 3 h. Wu et al. [11] reported that the solubility of CO2 in 45% [NH2p-mim][Br] aqueous solution was more than 0.444 mol CO2 per mol IL at 313 K and 106 kPa. However, the viscosity of the ILs was relatively high compared to MEA aqueous, which resulted in increased interfacial mass transfer resistance and decreased diffusion coefficient of gases [12]. In recent years, because of their high solubility in water, high absorption ability and low viscosity, amino acid ILs have been paid widespread attention by researchers [13]. Guo et al. [14] studied the characterization of carbon dioxide absorption in great amount of aqueous [Hmim][Gly] solution at low gas flow. Jing et al. [15] studied the characterization and kinetics of CO2 absorption in the 15% aqueous [N1111][Gly] solution, and the result showed that it had similar absorption ability compared to the aqueous TEA and DEA solution at the same experimental condition. Wu et al. [16] studied the kinetics of CO2 absorption into aqueous [C2mim][Gly] solution and found that the activity energy of [C2mim][Gly] solution absorbing CO2 was lower than that of MEA and DEA solution. Moreover, Zhang et al. [7] discussed the absorption of CO2 by traditional dicationic ionic liquids, and the results showed that the performance of CO2 capture by aqueous DILs solution was superior to the use of DILs only. Shiflett et al. [17] and his partners had modeled an ionic liquid that could reduce the energy losses by 16% compared to an industrial MEA process. Also, engineering
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Nomenclature N A C D E E1 G Y H H T Z M VL t x R Ha kG kL P n kOV k2
Absorption rate of CO2 (mol m2 s1) Gas–liquid interfacial area (m2) Concentration (mol m3) Diffusion coefficient (m2 s1) Enhancement factor Infinite enhancement factor Gas flow (mol s1) The mole percentage of CO2 in mix gas Henry’s constant (atm) Solubility constant (mol m3 Pa1) Temperature (K) The charge of ions Molar weight (g mol1) Volume of liquid phase (mL) Time (min) Mole fraction in the mix solution Ideal gas law constant (kmol m3 s1) Hatta number Gas-film mass transfer coefficient (mol m2 s1 Pa1) Liquid-film mass transfer coefficient (m s1) Pressure (Pa) Rotating speed (rpm) Overall kinetic constant (s1) Second-order reaction rate constant (L mol1 s1)
Greek letters m Viscosity (mPa s1) r Density (g cm3) d Solubility parameter (J cm3)1/2 Subscripts 1 Ionic liquid 2 CO2 or the second reaction order 3 Water G Gas phase L Liquid phase i Gas–liquid interface IL Ionic liquid
factor E, the second-order rate constant k2 and the overall reaction kinetic constant kOV were calculated by the two-film model. 2. Experimental 2.1. Chemicals CO2 (99.9%) and N2 (99.9%) were purchased from Xiangyun Gas Co., Ltd., China. All chemicals were obtained from Sinopharm Chemical Reagent Co., Ltd., China (AR degree). 1-Ethyl-3-methylimidazolium glutamate ([Emim]2[Glu]) was prepared as the similar procedures in literature [22]. 2.2. Experimental setups and procedures The experiment was carried out in the stirred-cell absorber at atmosphere pressure. The absorption system was mainly consisted of four parts: a absorber, two gas cylinders, recycled water system and an infrared analysis system (LT-21A). It was quite similar to the setup mentioned in literature [15,16,23]. The absorber was a glass cylinder (the inner diameter was 5 cm, length was 16 cm). The experimental setup is shown in Fig. 1. In order to form the required gas concentration, the two kinds of gases, CO2 and N2, went into a gas mixer before the absorber. Then the mixed gas was fed into the stirred-cell tank reactor and reacted with the [Emim]2[Glu] solution. The absorption temperature was controlled by the constant-temperature circulating bath. The value of volumetric flow rate was kept 120 5 mL min1 in the experiment. The mole percentage of CO2 in mixed gas was detected by an infrared Gas analyzer (LT-21A). The absorbent was injected into the absorber when the mole percentage of CO2 in mixed gas was stable. The absorption rate of CO2 into [Emim]2[Glu] aqueous solution could be expressed as follows: N¼
ðG1 Y 1 G2 Y 2 Þ A
ð1Þ
where G1 ,G2 are the flow rate (mol s1) of inlet and outlet, Y 1 , Y 2 are the mole percentage of CO2 at inlet and outlet, respectively. A is the gas–liquid interfacial area (m2). The values of G1 and G2 could be calculated by ideal gas law. 2.3. Physicochemical data measurement
design estimates of his study indicated that the investment of the ionic liquid process would be 11% lower than the amine based process and provided a 12% reduction in equipment footprint. From the above presentation, it could be found that amino acid ILs solution had excellent absorption ability of CO2. For the purpose of seeking for better absorbent of CO2, there is a current to study those imidazolium-based ionic liquids with high absorption rate and ability in aqueous solution. At present, the studies on CO2 absorption with diation ILs in the previous literatures were scarce, whether the diation ILs had good performance of CO2 capture. Moreover, there was no more information about physicochemical characters and CO2 absorption kinetics by [Emim]2[Glu] aqueous solutions. Thus, the kinetics of CO2 absorption by [Emim]2[Glu] aqueous solution, with molar concentration ranging from 0.3 to 1.5 mol L1, were studied at normal pressure in this paper. The density and viscosity of the [Emim]2[Glu] solution were measured. The solubility and diffusivity of CO2 in [Emim]2[Glu] solution were calculated using the model which had been reported [18–21] with the temperature from 298.15 K to 323.15 K. The effects of the [Emim]2[Glu] concentration, absorption temperature and CO2 concentration on the absorption were studied. The enhancement
2.3.1. Density and viscosity The density of the [Emim]2[Glu] aqueous solutions was measured by pycnometer method. The viscosity was measured by Ubbelohde viscosity meter. The Ubbelohde viscosity meter is shown in Fig. 2. Viscosity of ionic liquid under corresponding condition could be calculated according to following equation based on the time for ionic liquid of different concentration and pure water to decrease from Ubbelohde viscometer scale A to scale B at different temperature:
m¼
r mw t rw tw IL
ð2Þ
where m, mw are the viscosity of [Emim]2[Glu] aqueous solution and pure water, r, rw are the density of the [Emim]2[Glu] aqueous solution and pure water, respectively. t, tw are the residence time of [Emim]2[Glu] and pure water between scale A and B, respectively. rw and mw could be obtained through the literatures [14–16]. 2.3.2. The solubility of CO2 in [Emim]2[Glu] aqueous solutions According to the regular solution theory (RST) reported in the literatures [18–20]. there was a linear relationship between the natural log of the Henry’s constant and the square of the solubility parameters at low pressure. For a certain solution, when the
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Fig. 1. The experimental setup for CO2 absorption (1 and 2: CO2 and N2 cylinder; 3, 4, 13 and 14: gas/liquid controller; 5, 6: rotameter; 7: gas mixer; 8: stirred-cell absorber; 9 and 10: magnetic gearing; 11 and 12: variable-speed motor; 15: constant-temperature circulating bath; 16: stirrer-speed controller; 17: drying bottle; 18: three throw tap; 19: soap-film meter; 20: gas analyzer).
solubility parameter increased, the solubility increased, but Henry’s constant decreased. This theory can be expressed as Eq. (3): 2 ð d2 d1 Þ ln H2;1 ðatmÞ ¼ a þ b T
ð3Þ
where H is the Henry’s constant. d1 , d2 are the solubility parameter for IL and gas solute, respectively. The constants a and b [(J cm3)1] are only depending on the gas. Their values are 3.26 and 0.00084, respectively [18]. Solubility parameters are related with energy of vaporization, and their relation could be expressed a function. However, the energy of vaporization of ionic liquid was scarce in the previous scientific literatures. Fortunately, Camper et al. [18] deduced an empirical formula to estimate the solubility parameter of pure ionic liquid:
"
d1 ¼
1=3 ! 2:56 106 ðJ=molÞZ 1 Z 2 cm3 =mol ðM1 =r1 Þ4=3
1
2
V L1 þ V L3
2RT
ð5Þ
ð6Þ
ð7Þ
The value of H2;L could be calculated by Eqs. (3)–(7). According to Henry’s law [27], it could be converted into the solubility constant of CO2 (HAi ) in ionic liquid.
ð4Þ
ðM1 =r1 Þ1=3
In order to obtain the gas solubility of mixed solution, an empirical formula reported by O’Connell and Prausnitz [25] could be used. According to this formula, the gas solubility of the mixed solution was the function of the Henry’s constants in the pure solvents and the nonideality constants of the solute-free solvent mixture, their relation could be expressed as follow: lnH2;L ¼ x1 lnH2;1 þ x3 lnH2;3 a1;3 x1 x3
a1;3
ðd1 þ d3 Þ
1=3 !#1=2 0:367 cm3 =mol
where Z 1 , Z 2 are the charge of cation and anion, respectively. M represents the molar weight of ionic liquid (g mol1). r represents the density of pure ionic liquid (g cm3). An empirical formula had been reported to indicate the relation between the solubility parameter of CO2 and temperature [24]. expressed as Eq. (5):
d2 ¼ 0:0535T þ 28:26
where 1, 2, 3 stand for the ionic liquid, gas solute and water in the mix solution, respectively. x represents the mole fraction of the mixture. H2;L represents the Henry’s constant of CO2 in the mix solution. d3 has been reported in literature [26]. a1;3 could be estimated by a optimized form of Hildebrand’s equation, expressed as Eq. (7):
2.3.3. Diffusivity of CO2 in aqueous solutions of [Emim]2[Glu] According to the literatures reported by Barrett [28] and Danckwerts [21], the diffusivity of CO2 in [Emim]2[Glu] aqueous solution could be estimated by Eq. (8) owing to the aqueous [Emim]2[Glu] solution is the electrolyte solutions. ¼ D2;3 m0:8 ð8Þ D2;1 m0:8 1 3 T T
logD2;3 ¼ 8:1764 þ
712:5 2:591 105 T T2
ð9Þ
where D2;1 , D2;3 stand for the diffusion coefficient of CO2 in salt solvent and pure water, respectively, m1 , m3 is the viscosity of liquid and water, T is the temperature.
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absorption rate of CO2 into [Emim]2[Glu] solution could be calculated by Eq. (10). Also, the absorption rate was a function of liquid side mass transfer coefficient (kL ) and chemical reaction enhancement factors, described by Eq. (11) [15,32]: N ¼ kG ðPA PAi Þ
ð10Þ
N ¼ EkL C Ai
ð11Þ
where PAi is the CO2 partial pressure in the gas–liquid interface. PA is the CO2 initial partial pressure in the bulk. E and C Ai stand for the enhancement factor and the liquid phase concentration of CO2, respectively. According to Henry’s law, the value of C Ai could be calculated by Eq. (12): C Ai ¼ PAi HAi
ð12Þ
Thus, according to Eqs. (10)–(12), C Ai could be reduced to: P N C Ai ¼ HAi A kG
ð13Þ
The gas-film mass transfer coefficients (kG ) and the liquid-film mass transfer coefficients (kL ) at different temperature could be calculated by the following correlations [15,23]: DAN2 2=3 ð14Þ kG ¼ kGSO2 DSO2 N2 kL ¼ kWA
DAL DAW
2=3 ð15Þ
The values of DAN2 and DSO2 N2 could be estimated by the formula of Fuller–Schettler–Giddings [33] of Eq. (16). In this formula, Ma or Mb is the every component mole mass in mixed gas. P P P is the pressure of mix gas. V a or V b is the diffusion volume of gas molecular: D¼
1:013 106 T 1:75 ð1=Ma þ 1=Mb Þ1=2 h P i2 P P ð V a Þ1=3 þ ð V b Þ1=3
ð16Þ
Then, the liquid-film overall mass transfer (K L ) could be expressed by following correlation [32]: Fig. 2. Ubbelohde viscosity meter.
3. Theory 3.1. Mechanism of CO2 absorption in aqueous [Emim]2[Glu] solution In this paper, [Emim]+ was as cation, and glutamate as anion in the aqueous solutions of [Emim]2[Glu]. Since the aqueous alkaline salts of glutamate exhibits a similar reactivity toward CO2 as primary alkanolamines [15,29], the mechanism of CO2 absorption into [Emim]2[Glu] aqueous solution could be described by the Zwitterion theory reported by Caplow [30] and Danckwerts [31]. According to the zwitterion mechanism, the [Emim]2[Glu] solutions reacted with CO2 could be described as the formation of a zwitterion, and deprotonated by a base present in solution subsequently. 3.2. Mass transfer The chemical absorption of CO2 into [Emim]2[Glu] solution gas– liquid interface could be described as the processes of diffusion and chemical reaction. According to the two-film model, the
KL ¼
HAi 1 þ kL kG
ð17Þ
It was reported that CO2 absorption into the primary alkanolamines gas–liquid interface was the fast pseudo-first order reaction regime [22,28]. According to a pseudo first order reaction, when the Hatta number (Ha) could fulfill the relation Eq. (19), the value of Ha was equal to the enhancement factors (E). In Eq. (19), kOV is the overall reaction kinetic constant and DA is the diffusion coefficient of CO2 in solution. In this case, the overall reaction kinetic constant kOV could be calculated: 3 < Ha E1
Ha ¼
pffiffiffiffiffiffiffiffiffiffiffiffiffiffi DA kOV kL
ð18Þ
ð19Þ
Moreover, the rate of CO2 absorption (N) for a pseudo first order reaction could be also expressed by Eq. (20). According to Eq. (20), the second-order rate constant (k2 ) could be calculated: pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ð20Þ N ¼ C Ai k2 DA C IL
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Table 1 Density ðrÞ, viscosity ðmÞ, diffusivity ðDA Þ and solubility HAi of aqueous [Emim]2[Glu] solution for CO2 capture. T (K)
CIL (mol L1) 0.3
0.6
0.9
1.2
1.5
293.15 303.15 313.15 323.15
1.0453 0.0002 1.0423 0.0002 1.0386 0.0001 1.0307 0.0003
1.0631 0.0004 1.0601 0.0002 1.0560 0.0002 1.0520 0.0002
1.0828 0.0002 1.0782 0.0002 1.0746 0.0004 1.0694 0.0001
1.1034 0.0004 1.0989 0.0002 1.0943 0.0001 1.0895 0.0003
1.1225 0.0001 1.1210 0.0002 1.1159 0.0002 1.1108 0.0004
m (mPa S) 293.15 303.15 313.15 323.15
1.337 0.003 1.042 0.003 0.841 0.004 0.693 0.001
1.803 0.003 1.379 0.003 1.095 0.002 0.894 0.004
2.512 0.002 1.877 0.004 1.475 0.003 1.182 0.002
3.701 0.001 2.697 0.005 2.075 0.005 1.637 0.002
5.919 0.004 4.257 0.003 3.157 0.001 2.446 0.001
DA ( 109 m2 s1) 293.15 303.15 313.15 323.15
1.380 0.001 1.832 0.002 2.346 0.003 2.928 0.005
1.087 0.004 1.464 0.005 1.900 0.002 2.388 0.004
0.833 0.003 1.144 0.006 1.497 0.003 1.910 0.005
0.611 0.002 0.856 0.0004 1.139 0.005 1.472 0.003
0.420 0.005 0.594 0.004 0.814 0.002 1.068 0.001
H*Ai (104 mol m3 Pa1) 3.716 0.003 293.15 303.15 2.844 0.001 313.15 2.258 0.005 323.15 1.843 0.003
3.469 0.005 2.659 0.003 2.113 0.001 1.734 0.005
3.234 0.004 2.479 0.005 1.974 0.003 1.620 0.001
3.006 0.003 2.310 0.004 1.841 0.003 1.514 0.001
2.780 0.005 2.152 0.004 1.719 0.003 1.416 0.002
r (g cm1)
Fig. 3. The absorption rate of CO2 into different absorbents (absorption temperature: 303 K; CO2 concentration: 15%; gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
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4. Results and discussion 4.1. Physicochemical data The physicochemical date of the aqueous [Emim]2[Glu] solution with the mole concentrations from 0.3 to 1.5 mol L1 for CO2 capture at 293.15, 303.15, 313.15 and 323.15 K are shown in Table 1. The results showed that the density and viscosity of [Emim]2[Glu] aqueous solution increased as the concentration increased, decreased as the temperature increased. This tendency was in conformity with other similar amino acid ILs reported in literatures [14–16]. However, compared to the monoanion ILs in literature [14,16], such as [C2mim][Gly] and [Hmim][Gly] solution, aqueous [Emim]2[Glu] solution had greater density, viscosity, solubility and diffusivity of CO2 at the same mole concentration and temperature. Moreover, at a fixed temperature, the solubility and diffusivity of CO2 in [Emim]2[Glu] aqueous solution decreased as the mole concentration of IL increased, which was in agreement with the “salting-out” effect [34]. At some sort of fixed [Emim]2[Glu] concentration, the diffusivity of CO2 in [Emim]2[Glu] aqueous solution increased while the solubility of CO2 decreased as the temperature increased, which were in accord with the result in similar literatures [14,15]. 4.2. Absorption of CO2 into aqueous [Emim]2[Glu] solutions 4.2.1. Comparison with different absorbent In order to compare the CO2 capture properties of aqueous [Emim]2[Glu] solutions with different alkanolamine aqueous solutions, the absorption rates of CO2 at 303.15 K were measured in 0.9 mol L1 aqueous solution of [Emim]2[Glu], MEA and DEA, respectively, the results are shown in Fig. 3. According to Fig. 3, the absorption rate in aqueous [Emim]2[Glu] solution is lower than MEA aqueous solution, but higher than DEA aqueous solution at first several minutes. This result suggested that [Emim]2[Glu] is promising to be the candidate for CO2 absorption. 4.2.2. Effect of [Emim]2[Glu] mole concentration on the absorption of CO2 In this part, the absorption rates of CO2 into [Emim]2[Glu] aqueous solution were measured within first 30 min in concentration range from 0.3 to 1.5 mol L1 at 303.15 K. As shown in Fig. 4(a), the absorption rates increased as the [Emim]2[Glu] concentration increased at the same absorption time, and decreased as the absorption time increased. Furthermore, for the lower concentration, the absorption rates of CO2 increased sharply as the concentration increased, but for the higher concentration of [Emim]2[Glu] in solution, the increase tendency of absorption rate becomes slower due to their higher viscosity than that of lower concentration. However, for the low concentration of [Emim]2[Glu] solution, the absorption reached equilibrium quicker owing to the less molecules of [Emim]2[Glu] in the liquid. The relationship between the absorption rate and time could reflect that the absorption rate was controlled by the chemical reaction. The absorption amount obtained by calculating the integral characteristic value of absorption rate to time is shown in Fig. 4(b). The result showed that the absorption amount increased as the [Emim]2[Glu] concentration and the absorption time increased in the first 30 min. The variation tendency of high concentration was quicker than that of low concentration. The absorption load which is equal to the mole ratio of the absorption amount and [Emim]2[Glu] is shown in Fig. 4(c). The results indicated that the CO2 loading decreased as the concentration increased. Therefore, it could be found that increasing [Emim]2[Glu] concentration was conducive to increasing the absorption rate and amount of CO2. However, high concentration which resulted in low CO2 loading will waste
Fig. 4. Absorption rate (a), absorption mount (b) and absorption load (c) of CO2 into aqueous [Emim]2[Glu] with different concentrations. (T = 303 K; PCO2 = 15.20 kPa, gas mixture: 15% CO2–85% N2, gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
absorbent and energy in industry. So, it is necessary to choose an appropriate concentration for CO2 capture. The concentration of 0.9 mol L1 was chosen to detect the kinetic parameters in this paper.
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Fig. 5. Absorption rate (a) and absorption amount (b) of CO2 into 0.9 mol L1 [Emim]2[Glu] aqueous solution with different temperatures. (PCO2 = 15.20 kPa, gas mixture: 15% CO2: 85% N2, gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
4.2.3. Effect of temperature on the absorption In order to obtain the effect of temperature on CO2 absorption, the absorption rate and amount of CO2 into 0.9 mol L1 aqueous [Emim]2[Glu] solution were measured in different temperature, as shown in Fig. 5(a) and (b), respectively. From Fig. 5(a), obviously, the absorption rate of CO2 is influenced by temperature greatly in the investigated temperature range. At the same absorption time, the absorption rate of CO2 increased obviously as the temperature increased. This phenomenon could be explained by their changed
viscosity. However, the absorption rate of CO2 decreased slower at low temperature than that of high temperatures. Fig. 5(b) shows that the absorption amount of CO2 varied slightly at 303.15– 323.15 K, but their values were all more than that at 293.15 K in the first 30 min. This result could be explained that the absorption rate was faster at higher temperature at the beginning. This process was similar to the system on CO2 absorption, such as 15% [N1111][Gly] [15] and 1.0 mol L 1 [C2mim][Gly] [16].
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Fig. 6. Effect of CO2 concentration on absorption rate (a) and absorption amount (b) of CO2 into 0.9 mol L1 [Emim]2[Glu] aqueous solution. (T = 303 K; gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
Table 2 Mass-transfer coefficient of the gas and liquid films. T (K)
kG;CO2 (106 mol m2 s1 Pa1)
kL;CO2 (105 m s1)
KL (105 m s1)
293.15 303.15 313.15 323.15
6.15 6.14 6.13 6.12
1.29 1.59 1.89 2.12
1.29 1.59 1.89 2.12
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Table 3 Effect of liquid phase volume and mass transfer coefficient of liquid phase on absorption rate at 303.15 K. t (min)
VL = 20 mL, N (103 mol m2 s1)
VL = 40 mL, N (103 mol m2 s1)
nL = 130 rpm, N (103 mol m2 s1)
nL = 250 rpm, N (103 mol m2 s1)
0 1 2 5 10 20 30
2.792 0.028 2.755 0.012 2.739 0.021 2.620 0.011 2.264 0.009 1.773 0.015 1.282 0.017
2.811 0.028 2.795 0.024 2.791 0.021 2.774 0.013 2.684 0.017 2.299 0.009 1.898 0.012
2.828 0.023 2.808 0.017 2.784 0.013 2.772 0.008 2.691 0.016 2.318 0.031 1.907 0.017
2.811 0.028 2.795 0.024 2.791 0.021 2.774 0.013 2.684 0.017 2.299 0.009 1.898 0.012
4.2.4. Effect of CO2 concentration on the absorption The CO2 content of exhaust gas was almost among 8 to 15% in industry, thus the effect of CO2 content on the absorption of CO2 into 0.9 mol L1 [Emim]2[Glu] aqueous solution was measured at 303.15 K with CO2 concentration at inlet ranging from 8% (v%) to 15% (v%). As shown in Fig. 6(a) and (b), the absorption rates and amount increased as the CO2 concentration increased in first 30 min. It could be explained that the higher CO2 content had greater driving force between gas phase and gas–liquid interface. 4.3. Kinetics analysis 4.3.1. Mass-transfer coefficient The mass-transfer coefficient of gas and liquid films were needed to check the conditions for different controlling regimes, to calculate the reaction rate parameters. The values of gas and liquid films mass-transfer coefficient were calculated from the experiment data at different temperature. According to the Eqs. (12)– (14), the gas-film and liquid-film mass-transfer coefficients of CO2 in aqueous [Emim]2[Glu] solution at 293.15–323.15 K were obtained, the result is presented in Table 2. According to Table 2, as the temperature increased, the value of gas film mass-transfer coefficient varied slightly, but the liquid film and overall masstransfer coefficient increased greatly. It could be explained that the viscosity of [Emim]2[Glu] aqueous solution decreased as the temperature increased. Moreover, the liquid film mass-transfer coefficient was equal to the overall mass-transfer coefficient. This phenomenon could demonstrate that the reaction was controlled by liquid film. 4.3.2. Absorption kinetics According to the method suggested by Jing and Wu [14–16], the experiments were carried out with gas mixture of 15% CO2 and 85% N2 and the stirring speed for gas phase of 250 rpm at 303.15 K. The absorption rates of CO2 into 0.9 mol L1 [Emim]2[Glu] solution were measured in different liquid phase volumes and liquid stirring speed, the results are presented in Table 3. Obviously, the absorption rate of CO2 increased apparently as the liquid phase volume increased after 5 min for CO2 absorption, but the effect of mass transfer coefficients on the CO2 absorption rates could be
ignored. This phenomenon could be explained that the [Emim]2[Glu] concentration in lower liquid phase volume solution reduced very quickly and the absorption was controlled by gas-film. And according to the discussion in Section 4.2.2, the absorption rates increased as the [Emim]2[Glu] concentration increased, which implied the reaction evidently associated with the concentration of [Emim]2[Glu] other than related to the liquid phase volume and the liquid phase mass transfer coefficient. Consequently, the kinetic region of CO2 absorption into aqueous [Emim]2[Glu] solution was the fast pseudo-first order reaction regime. According to the two-film model and the criterion of the fast pseudo-first order reaction regime described in Section 3.2, the values of enhancement factor E, Hatta number Ha, overall reaction kinetic constant kOV and second-order rate constant k2 could be calculated with the concentration of [Emim]2[Glu] solution ranging from 0.3 to 1.5 mol L1 at 303.15 K by Eqs. (10)–(18), as shown in Table 4. The results showed that the second-order rate constant k2 and the overall reaction kinetic constant kOV increased significantly in 0.3–1.5 mol L1 aqueous [Emim]2[Glu] solution as concentration increased. The enhancement factor E plotted of the square root of concentration is shown in Fig. 7. The result showed that enhancement factor increased as the concentration of [Emim]2[Glu] solution increased, and there was a linear relationship between the enhancement factor and the square root of the [Emim]2[Glu] solution concentration. Moreover, the values of E, Ha, kOV and k2 of CO2 absorption in 0.9 mol L1 aqueous [Emim]2[Glu] solution were also calculated at 293.15–303.15 K and the results are presented in Table 5. Obviously, the second-order reaction rate and overall reaction kinetic constant kOV increased as the temperature increased. The kinetic data can be calculated from an Arrhenius plot of k2 1=T, represented as following: 3665 ð21Þ k2 ¼ 9:253 107 exp T From this equation, the value of activation energy could be calculated to be 30.48 kJ mol1. This value was a little higher than some monoanion ionic liquid, such as [C2mim][Gly] (23.29 kJ mol1) and [Hmim][Gly] (25.36 kJ mol1) reported by Wu
Table 4 The kinetic parameters for CO2 absorption into [Emim]2[Glu] aqueous solution at 303.15 K. CIL (mol L1)
N 103 (mol m2 s1)
PG (Pa)
CAi (mol m3)
E
Ha
kOV (s1)
K2 (L mol1 s1)
0.3 0.6 0.9 1.2 1.5
1.461 0.011 2.153 0.027 2.811 0.028 3.321 0.036 3.557 0.041
12331.253 30.323 10801.245 21.152 9088.853 22.647 7487.918 13.301 7133.28 15.48
4.367 0.011 4.078 0.007 3.752 0.009 3.416 0.004 3.220 0.005
21.043 0.223 33.207 0.486 47.124 0.542 61.137 0.781 69.484 0.897
21.043 0.223 33.207 0.486 47.124 0.542 61.137 0.781 69.484 0.897
61.093 3.163 190.36 6.441 490.59 11.446 1103.53 34.557 2053.59 55.736
67.881 3.745 211.514 7.883 545.093 13.471 1226.142 37.992 2281.769 58.337
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Fig. 7. Enhancement factor of CO2 absorption into aqueous [Emim]2[Glu] solution with different concentrations.
Table 5 The kinetic parameters for CO2 absorption into 0.9 mol L1 [Emim]2[Glu] aqueous solution at different temperature. T (K)
N 103 (mol m2 s1)
PG (Pa)
CAi (mol m3)
E
Ha
kOV (s1)
k2 (L mol1 s1)
293.15 303.15 313.15 323.15
2.425 0.021 2.811 0.028 3.129 0.033 3.227 0.031
10223.692 24.325 9088.853 22.647 8014.808 17.693 7437.255 13.968
4.912 0.012 3.752 0.009 2.976 0.006 2.432 0.005
38.272 0.413 47.124 0.542 55.634 0.624 62.593 0.719
38.272 0.413 47.124 0.542 55.634 0.624 62.593 0.719
292.500 8.764 490.588 11.446 738.644 23.683 921.804 29.543
325.000 11.583 545.097 13.471 820.715 17.679 1024.226 25.994
et al. [16] and Jing et al. [15], respectively. But the activation energy in this paper was much lower than the value of sodium glycinate (63.8 kJ mol1) and MEA (45 kJ mol1) calculated by Lee et al. [35] and Versteeg et al. [36], respectively.
0.9 mol L1 [Emim]2[Glu] was equal to 30.48 kJ mol1, lower than some absorbents reported in literatures.
5. Conclusions
The authors are grateful for the assistance offered by the National Science and Technology Support Program, grant no. 2012BAC02B04.
The density and viscosity of [Emim]2[Glu] aqueous solutions increased as [Emim]2[Glu] concentration increased and decreased as the temperature increased at all the concentrations and temperature investigated. As temperature increased, the diffusivity of [Emim]2[Glu] solutions increased, and the solubility decreased. Furthermore, as the [Emim]2[Glu] concentration increased, the diffusivity and solubility both decreased. The absorption rate and amount of CO2 all increased significantly as the increase of the [Emim]2[Glu] concentration and temperature, but the CO2 load decreased as the increase of the [Emim]2[Glu] concentration. Compared with other absorbents, the absorption rate of CO2 into 0.9 mol L1 [Emim]2[Glu] solution was lower than MEA, but higher than DEA at the same experimental condition. According to the judge method reported in literatures, It was proved that kinetics region of CO2 absorption into aqueous [Emim]2[Glu] was the fast pseudo-first order reaction regime. 1=2
The enhancement factor was linear with C ILs , and the reaction rate could be expressed by the following constant equal:k2 ¼ 9:253 107 expð3665=T Þ. The activation energy of
Acknowledgment
References [1] J.M. Navaza, D. Gomez-Diaz, M.D. La Rubia, Removal process of CO2 using MDEA aqueous solutions in a bubble column reactor, Chem. Eng. J. 146 (2009) 184–188. [2] D.J. Fauth, E.A. Frommell, J.S. Hoffman, R.P. Reasbeck, H.W. Pennline, Eutectic salt promoted lithium zirconate: novel high temperature sorbent for CO2 capture, Fuel Process. Technol. 86 (2005) 1503–1521. [3] B.R. Strazisar, R.R. Anderson, C.M. White, Degradation pathways for monoethanolamine in a CO2 capture facility, Energy Fuels 17 (2003) 1034– 1039. [4] J.T. Yeh, K.P. Resnik, K. Rygle, H.W. Pennline, Semi-batch absorption and regeneration studies for CO2 capture by aqueous ammonia, Fuel Process. Technol. 86 (2005) 1533–1546. [5] J.L. Anthony, J.L. Anderson, E.J. Maginn, J.F. Brennecke, Anion effects on gas solubility in ionic liquids, J. Phys. Chem. B109 (2005) 6366–6374. [6] C. Cadena, J.L. Anthony, J.K. Shah, T.I. Morrow, J.F. Brennecke, E.J. Maginn, Why is CO2 so soluble in imidazolium-based ionic liquids? J Am. Chem. Soc. 126 (2004) 5300–5308. [7] Y. Zhang, P. Yu, Y.B. Luo, Absorption of CO2 by amino acid-functionalized and traditional dicationic ionic liquids: properties, Henry’s law constants and mechanisms, Chem. Eng. J. 214 (2013) 355–363. [8] R.K. Rogers, K.R. Seddon, Ionic Liquids as Green Solvents: Progress and Oxford Prospects, University Press, Washington, D.C, 2003.
J. Ma et al. / Journal of Environmental Chemical Engineering 4 (2016) 1137–1147 [9] Y. Zhang, Z. Wu, S. Chen, P. Yu, Y. Luo, CO2 captue by imidazolate-based ionic liquids: effect of functionalized cation and dication, Ind. Eng. Chem. Res. 52 (2013) 6069–6075. [10] E.D. Bates, R.D. Mayton, I. Ntai, J.H. Davis, CO2 capture by a task-specific ionic liquid, J. Am. Chem. Soc. 124 (2002) 926–927. [11] Y.L. Wu, Z. Jiao, G.N. Wang, Y.T. Wu, Z.B. Zhang, Synthesis, characterization and absorption efficiency of an ionic liquid for the absorption of CO2, Fine Chem. 24 (2007) 324–327. [12] K.E. Gutowski, E.J. Maginn, Amine-functionalized task-specific ionic liquids: a mechanistic explanation for the dramatic increase in viscosity upon complexation with CO2 from molecular simulation, J. Am. Chem. Soc. 130 (2008) 14690–14704. [13] C.-C. Wei, G. Puxty, P. Feron, Amino acid salts for CO2 capture at flue gas temperatures, Chem. Eng. Sci. 107 (2014) 218–226. [14] H. Guo, Z. Zhou, G. Jing, Kinetics of carbon dioxide absorption into aqueous [Hmim][Gly] solution, Int. J. Greenh. Gas Control 16 (2013) 197–205. [15] G. Jing, L. Zhou, Z. Zhou, Characterization and kinetics of carbon dioxide absorption into aqueous tetramethylammonium glycinate solution, Chem. Eng. J. 181–182 (2012) 85–92. [16] Z.K. Wu, Y. Zhang, W.K. Lei, P. Yu, Y.B. Luo, Kinetics of CO2 absorption into aqueous 1-ethyl-3-methylimidazolium glycinate solution, Chem. Eng. J. 264 (2015) 744–752. [17] Mark B. Shiflett, David W. Drew, Robert A. Cantini, A. Yokozeki, Carbon dioxide capture using ionic liquid 1-butyl-3-methylimidazolium acetate, Energy Fuels 24 (2010) 5781–5789. [18] D. Camper, C. Becker, C. Koval, R. Noble, Low pressure hydrocarbon solubility in room temperature ionic liquids containing imidazolium rings interpreted using regular solution theory, Ind. Eng. Chem. Res. 44 (2005) 1928–1933. [19] P. Scovazzo, D. Camper, J. Kieft, J. Poshusta, C. Koval, R. Noble, Regular solution theory and CO2 gas solubility in room-temperature ionic liquids, Ind. Eng. Chem. Res. 43 (2004) 6855–6860. [20] A. Finotello, J.E. Bara, D. Camper, R.D. Noble, Room-temperature ionic liquids: temperature dependence of gas solubility selectivity, Ind. Eng. Chem. Res. 47 (2007) 3453–3459. [21] P.V. Danckwerts, Gas–Liquid Reactions, McGraw-Hill, New York, 1970.
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[22] K. Fukumoto, M. Yoshizawa, H. Ohno, Room temperature ionic liquids from 20 natural amino acids, J. Am. Chem. Soc. 127 (2005) 2398–2399. [23] X. Gao, et al., Gas–liquid absorption reaction between (NH4)2SO3 solution and SO2 for ammonia-based wet flue gas desulfurization, Appl. Energy 2010 (2010) 023. http://dx.doi.org/10.1016/j.apenergy.03. [24] D. Camper, P. Scovazzo, C. Koval, R. Noble, Gas solubilities in room-temperature ionic liquids, Ind. Eng. Chem. Res. 43 (2004) 3049–3054. [25] J.P. O’Connell, J.M. Prausnitz, Thermodynamics of gas solubility in mixed solvents, I&EC Fundam. 3 (1964) 347–364. [26] A.F.M. Barton, Handbook of Solubility Parameters and Other Cohesion Parameters, CRC Press, Boca Raton, FL, 1983. [27] C. Chen, B. Liu, G. Liu, C. Chai, Chemical Principle, Tianjin Science and Technology Press, Tianjin, 2001. [28] P.V.L. Barrett, Gas Absorption on a Sieve Plate, Ph.D. Thesis, University of Cambridge, 1966, 2016. [29] P.S. Kumar, J.A. Hogendoorn, G.F. Versteeg, P.H.M. Feron, Kinetics of the reaction of CO2 with aqueous potassium salt of taurine and glycine, AIChE J. 49 (2003) 203–213. [30] M. Caplow, Kinetics of carbamate formation and breakdown, J. Am. Chem. Soc. 90 (1968) 6795–6803. [31] P.V. Danckwerts, The reaction of CO2 with ethanolamines, Chem. Eng. Sci. 34 (1979) 443. [32] Y. Shi, D. Littlejohn, S.G. Chang, Kinetics of NO absorption in aqueous iron(II) bis(2,3-dimercapto-1-propanesulfonate) solutions using a stirred reactor, Ind. Eng. Chem. Res. 35 (1996) 1668–1672. [33] Wuhan University, The Fundamentals of Chemical Engineering, Higher Education Press, 2009. [34] S. Weisenberger, A. Schumpe, Estimation of gas solubilities in salt solutions at temperatures from 273 K to 363 K, AIChE J. 42 (1996) 298–300. [35] S. Lee, H.-J. Song, S. Maken, J.-W. Park, Kinetics of CO2 absorption in aqueous sodium glycinate solutions, Ind. Eng. Chem. Res. 46 (2007) 1578–1583. [36] G.F. Versteeg, L.A.J. Van Dijck, W.P.M. Van Swaaij, On the kinetics between CO2 and alkanolamines both in aqueous and nonaqueous solutions. An overview, Chem. Eng. Commun. 144 (1996) 133.
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Journal of Environmental Chemical Engineering journal homepage: www.elsevier.com/locate/jece
Characterization and kinetics of CO2 absorption into aqueous 1-ethyl -3-methylimidazolium glutamate solution Jie Ma, Zaikun Wu, Wenkang Lei, Ping Yu, Yunbai Luo* College of Chemistry and Molecular Sciences, Wuhan University, Wuhan 430072, China
A R T I C L E I N F O
A B S T R A C T
Article history: Received 26 August 2015 Received in revised form 5 December 2015 Accepted 12 January 2016 Available online 14 January 2016
The kinetic of carbon dioxide (CO2) into aqueous 1-ethyl-3-methylimidazolium glutamate ([Emim]2[Glu]) solution was investigated by a double stirred-cell absorber. The Density and viscosity of [Emim]2[Glu] solution were measured with concentrations ranging from 0.3 to 1.5 mol L1 at temperatures from 298.15 to 323.15 K. The diffusivity of CO2 in [Emim]2[Glu] solution were estimated at the same temperatures and concentrations. The effects of the [Emim]2[Glu] concentration, absorption temperature and CO2 concentration were studied. The results showed that absorption CO2 into aqueous [Emim]2[Glu] solution was the fast pseudo-first order reaction regime. The enhancement factor E, the overall reaction kinetic constant kOV and the second-order rate constant k2 were calculated by the two-
Keywords: Kinetic 1-Ethyl-3-methylimidazolium glutamate Physicochemical data Carbon dioxide Absorption
1=2
film theory. The result showed that E was linear with C ILs , and the second-order rate constant could be expressed as following equation: k2 ¼ 9:253 10 expð3665=T Þ. ã 2016 Elsevier Ltd. All rights reserved.
1. Introduction Carbon dioxide, the main greenhouse gas in the atmosphere, causes global warming. There is no doubt that fossil fuels and petroleum industries will still act as an important role for a long time, which makes CO2 capture an essential process to our society. Aqueous alkanolamine solutions, such as monoethanolamine (MEA), diethanolamine (DEA) and N-methyldiethanolamine (MDEA) [1], is the most widespread absorbent for the removal of CO2 in flue gas because of the advantage of high reactivity and low solvent cost. However, aqueous alkanolamine solutions have several serious drawbacks: solvent loss, corrosion, regeneration fee and oxidative degradation [2–4]. Recently, ionic liquids with low vapor pressure, high thermal stability, low melting point, low energy consumption and nontoxic, have been considered as new absorbents for CO2 removal [5,6]. As many literatures reported [7–9], CO2 had been demonstrated to have remarkable solubility in aqueous imidazolium-based ionic liquids solution. However, carbon dioxide dissolution in imidazolium-based ILs not only are a physical process but also have low absorption rate. In view of these reasons, a multitude kinds of ionic liquids which are amino-functionalized were regarded as a kind of novel absorbent for CO2 capture due to their large absorbing capacity, good physical and chemical properties at atmospheric
* Corresponding author. Fax: +86 2768752511. E-mail address: [email protected] (Y. Luo). http://dx.doi.org/10.1016/j.jece.2016.01.016 2213-3437/ ã 2016 Elsevier Ltd. All rights reserved.
7
pressure had been reported. Bates et al. [10] reported that the molar uptake of CO2 per mole of [NH2p-bim][BF4] was closed to the theoretical maximum loading by MEA during 3 h. Wu et al. [11] reported that the solubility of CO2 in 45% [NH2p-mim][Br] aqueous solution was more than 0.444 mol CO2 per mol IL at 313 K and 106 kPa. However, the viscosity of the ILs was relatively high compared to MEA aqueous, which resulted in increased interfacial mass transfer resistance and decreased diffusion coefficient of gases [12]. In recent years, because of their high solubility in water, high absorption ability and low viscosity, amino acid ILs have been paid widespread attention by researchers [13]. Guo et al. [14] studied the characterization of carbon dioxide absorption in great amount of aqueous [Hmim][Gly] solution at low gas flow. Jing et al. [15] studied the characterization and kinetics of CO2 absorption in the 15% aqueous [N1111][Gly] solution, and the result showed that it had similar absorption ability compared to the aqueous TEA and DEA solution at the same experimental condition. Wu et al. [16] studied the kinetics of CO2 absorption into aqueous [C2mim][Gly] solution and found that the activity energy of [C2mim][Gly] solution absorbing CO2 was lower than that of MEA and DEA solution. Moreover, Zhang et al. [7] discussed the absorption of CO2 by traditional dicationic ionic liquids, and the results showed that the performance of CO2 capture by aqueous DILs solution was superior to the use of DILs only. Shiflett et al. [17] and his partners had modeled an ionic liquid that could reduce the energy losses by 16% compared to an industrial MEA process. Also, engineering
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Nomenclature N A C D E E1 G Y H H T Z M VL t x R Ha kG kL P n kOV k2
Absorption rate of CO2 (mol m2 s1) Gas–liquid interfacial area (m2) Concentration (mol m3) Diffusion coefficient (m2 s1) Enhancement factor Infinite enhancement factor Gas flow (mol s1) The mole percentage of CO2 in mix gas Henry’s constant (atm) Solubility constant (mol m3 Pa1) Temperature (K) The charge of ions Molar weight (g mol1) Volume of liquid phase (mL) Time (min) Mole fraction in the mix solution Ideal gas law constant (kmol m3 s1) Hatta number Gas-film mass transfer coefficient (mol m2 s1 Pa1) Liquid-film mass transfer coefficient (m s1) Pressure (Pa) Rotating speed (rpm) Overall kinetic constant (s1) Second-order reaction rate constant (L mol1 s1)
Greek letters m Viscosity (mPa s1) r Density (g cm3) d Solubility parameter (J cm3)1/2 Subscripts 1 Ionic liquid 2 CO2 or the second reaction order 3 Water G Gas phase L Liquid phase i Gas–liquid interface IL Ionic liquid
factor E, the second-order rate constant k2 and the overall reaction kinetic constant kOV were calculated by the two-film model. 2. Experimental 2.1. Chemicals CO2 (99.9%) and N2 (99.9%) were purchased from Xiangyun Gas Co., Ltd., China. All chemicals were obtained from Sinopharm Chemical Reagent Co., Ltd., China (AR degree). 1-Ethyl-3-methylimidazolium glutamate ([Emim]2[Glu]) was prepared as the similar procedures in literature [22]. 2.2. Experimental setups and procedures The experiment was carried out in the stirred-cell absorber at atmosphere pressure. The absorption system was mainly consisted of four parts: a absorber, two gas cylinders, recycled water system and an infrared analysis system (LT-21A). It was quite similar to the setup mentioned in literature [15,16,23]. The absorber was a glass cylinder (the inner diameter was 5 cm, length was 16 cm). The experimental setup is shown in Fig. 1. In order to form the required gas concentration, the two kinds of gases, CO2 and N2, went into a gas mixer before the absorber. Then the mixed gas was fed into the stirred-cell tank reactor and reacted with the [Emim]2[Glu] solution. The absorption temperature was controlled by the constant-temperature circulating bath. The value of volumetric flow rate was kept 120 5 mL min1 in the experiment. The mole percentage of CO2 in mixed gas was detected by an infrared Gas analyzer (LT-21A). The absorbent was injected into the absorber when the mole percentage of CO2 in mixed gas was stable. The absorption rate of CO2 into [Emim]2[Glu] aqueous solution could be expressed as follows: N¼
ðG1 Y 1 G2 Y 2 Þ A
ð1Þ
where G1 ,G2 are the flow rate (mol s1) of inlet and outlet, Y 1 , Y 2 are the mole percentage of CO2 at inlet and outlet, respectively. A is the gas–liquid interfacial area (m2). The values of G1 and G2 could be calculated by ideal gas law. 2.3. Physicochemical data measurement
design estimates of his study indicated that the investment of the ionic liquid process would be 11% lower than the amine based process and provided a 12% reduction in equipment footprint. From the above presentation, it could be found that amino acid ILs solution had excellent absorption ability of CO2. For the purpose of seeking for better absorbent of CO2, there is a current to study those imidazolium-based ionic liquids with high absorption rate and ability in aqueous solution. At present, the studies on CO2 absorption with diation ILs in the previous literatures were scarce, whether the diation ILs had good performance of CO2 capture. Moreover, there was no more information about physicochemical characters and CO2 absorption kinetics by [Emim]2[Glu] aqueous solutions. Thus, the kinetics of CO2 absorption by [Emim]2[Glu] aqueous solution, with molar concentration ranging from 0.3 to 1.5 mol L1, were studied at normal pressure in this paper. The density and viscosity of the [Emim]2[Glu] solution were measured. The solubility and diffusivity of CO2 in [Emim]2[Glu] solution were calculated using the model which had been reported [18–21] with the temperature from 298.15 K to 323.15 K. The effects of the [Emim]2[Glu] concentration, absorption temperature and CO2 concentration on the absorption were studied. The enhancement
2.3.1. Density and viscosity The density of the [Emim]2[Glu] aqueous solutions was measured by pycnometer method. The viscosity was measured by Ubbelohde viscosity meter. The Ubbelohde viscosity meter is shown in Fig. 2. Viscosity of ionic liquid under corresponding condition could be calculated according to following equation based on the time for ionic liquid of different concentration and pure water to decrease from Ubbelohde viscometer scale A to scale B at different temperature:
m¼
r mw t rw tw IL
ð2Þ
where m, mw are the viscosity of [Emim]2[Glu] aqueous solution and pure water, r, rw are the density of the [Emim]2[Glu] aqueous solution and pure water, respectively. t, tw are the residence time of [Emim]2[Glu] and pure water between scale A and B, respectively. rw and mw could be obtained through the literatures [14–16]. 2.3.2. The solubility of CO2 in [Emim]2[Glu] aqueous solutions According to the regular solution theory (RST) reported in the literatures [18–20]. there was a linear relationship between the natural log of the Henry’s constant and the square of the solubility parameters at low pressure. For a certain solution, when the
J. Ma et al. / Journal of Environmental Chemical Engineering 4 (2016) 1137–1147
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Fig. 1. The experimental setup for CO2 absorption (1 and 2: CO2 and N2 cylinder; 3, 4, 13 and 14: gas/liquid controller; 5, 6: rotameter; 7: gas mixer; 8: stirred-cell absorber; 9 and 10: magnetic gearing; 11 and 12: variable-speed motor; 15: constant-temperature circulating bath; 16: stirrer-speed controller; 17: drying bottle; 18: three throw tap; 19: soap-film meter; 20: gas analyzer).
solubility parameter increased, the solubility increased, but Henry’s constant decreased. This theory can be expressed as Eq. (3): 2 ð d2 d1 Þ ln H2;1 ðatmÞ ¼ a þ b T
ð3Þ
where H is the Henry’s constant. d1 , d2 are the solubility parameter for IL and gas solute, respectively. The constants a and b [(J cm3)1] are only depending on the gas. Their values are 3.26 and 0.00084, respectively [18]. Solubility parameters are related with energy of vaporization, and their relation could be expressed a function. However, the energy of vaporization of ionic liquid was scarce in the previous scientific literatures. Fortunately, Camper et al. [18] deduced an empirical formula to estimate the solubility parameter of pure ionic liquid:
"
d1 ¼
1=3 ! 2:56 106 ðJ=molÞZ 1 Z 2 cm3 =mol ðM1 =r1 Þ4=3
1
2
V L1 þ V L3
2RT
ð5Þ
ð6Þ
ð7Þ
The value of H2;L could be calculated by Eqs. (3)–(7). According to Henry’s law [27], it could be converted into the solubility constant of CO2 (HAi ) in ionic liquid.
ð4Þ
ðM1 =r1 Þ1=3
In order to obtain the gas solubility of mixed solution, an empirical formula reported by O’Connell and Prausnitz [25] could be used. According to this formula, the gas solubility of the mixed solution was the function of the Henry’s constants in the pure solvents and the nonideality constants of the solute-free solvent mixture, their relation could be expressed as follow: lnH2;L ¼ x1 lnH2;1 þ x3 lnH2;3 a1;3 x1 x3
a1;3
ðd1 þ d3 Þ
1=3 !#1=2 0:367 cm3 =mol
where Z 1 , Z 2 are the charge of cation and anion, respectively. M represents the molar weight of ionic liquid (g mol1). r represents the density of pure ionic liquid (g cm3). An empirical formula had been reported to indicate the relation between the solubility parameter of CO2 and temperature [24]. expressed as Eq. (5):
d2 ¼ 0:0535T þ 28:26
where 1, 2, 3 stand for the ionic liquid, gas solute and water in the mix solution, respectively. x represents the mole fraction of the mixture. H2;L represents the Henry’s constant of CO2 in the mix solution. d3 has been reported in literature [26]. a1;3 could be estimated by a optimized form of Hildebrand’s equation, expressed as Eq. (7):
2.3.3. Diffusivity of CO2 in aqueous solutions of [Emim]2[Glu] According to the literatures reported by Barrett [28] and Danckwerts [21], the diffusivity of CO2 in [Emim]2[Glu] aqueous solution could be estimated by Eq. (8) owing to the aqueous [Emim]2[Glu] solution is the electrolyte solutions. ¼ D2;3 m0:8 ð8Þ D2;1 m0:8 1 3 T T
logD2;3 ¼ 8:1764 þ
712:5 2:591 105 T T2
ð9Þ
where D2;1 , D2;3 stand for the diffusion coefficient of CO2 in salt solvent and pure water, respectively, m1 , m3 is the viscosity of liquid and water, T is the temperature.
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absorption rate of CO2 into [Emim]2[Glu] solution could be calculated by Eq. (10). Also, the absorption rate was a function of liquid side mass transfer coefficient (kL ) and chemical reaction enhancement factors, described by Eq. (11) [15,32]: N ¼ kG ðPA PAi Þ
ð10Þ
N ¼ EkL C Ai
ð11Þ
where PAi is the CO2 partial pressure in the gas–liquid interface. PA is the CO2 initial partial pressure in the bulk. E and C Ai stand for the enhancement factor and the liquid phase concentration of CO2, respectively. According to Henry’s law, the value of C Ai could be calculated by Eq. (12): C Ai ¼ PAi HAi
ð12Þ
Thus, according to Eqs. (10)–(12), C Ai could be reduced to: P N C Ai ¼ HAi A kG
ð13Þ
The gas-film mass transfer coefficients (kG ) and the liquid-film mass transfer coefficients (kL ) at different temperature could be calculated by the following correlations [15,23]: DAN2 2=3 ð14Þ kG ¼ kGSO2 DSO2 N2 kL ¼ kWA
DAL DAW
2=3 ð15Þ
The values of DAN2 and DSO2 N2 could be estimated by the formula of Fuller–Schettler–Giddings [33] of Eq. (16). In this formula, Ma or Mb is the every component mole mass in mixed gas. P P P is the pressure of mix gas. V a or V b is the diffusion volume of gas molecular: D¼
1:013 106 T 1:75 ð1=Ma þ 1=Mb Þ1=2 h P i2 P P ð V a Þ1=3 þ ð V b Þ1=3
ð16Þ
Then, the liquid-film overall mass transfer (K L ) could be expressed by following correlation [32]: Fig. 2. Ubbelohde viscosity meter.
3. Theory 3.1. Mechanism of CO2 absorption in aqueous [Emim]2[Glu] solution In this paper, [Emim]+ was as cation, and glutamate as anion in the aqueous solutions of [Emim]2[Glu]. Since the aqueous alkaline salts of glutamate exhibits a similar reactivity toward CO2 as primary alkanolamines [15,29], the mechanism of CO2 absorption into [Emim]2[Glu] aqueous solution could be described by the Zwitterion theory reported by Caplow [30] and Danckwerts [31]. According to the zwitterion mechanism, the [Emim]2[Glu] solutions reacted with CO2 could be described as the formation of a zwitterion, and deprotonated by a base present in solution subsequently. 3.2. Mass transfer The chemical absorption of CO2 into [Emim]2[Glu] solution gas– liquid interface could be described as the processes of diffusion and chemical reaction. According to the two-film model, the
KL ¼
HAi 1 þ kL kG
ð17Þ
It was reported that CO2 absorption into the primary alkanolamines gas–liquid interface was the fast pseudo-first order reaction regime [22,28]. According to a pseudo first order reaction, when the Hatta number (Ha) could fulfill the relation Eq. (19), the value of Ha was equal to the enhancement factors (E). In Eq. (19), kOV is the overall reaction kinetic constant and DA is the diffusion coefficient of CO2 in solution. In this case, the overall reaction kinetic constant kOV could be calculated: 3 < Ha E1
Ha ¼
pffiffiffiffiffiffiffiffiffiffiffiffiffiffi DA kOV kL
ð18Þ
ð19Þ
Moreover, the rate of CO2 absorption (N) for a pseudo first order reaction could be also expressed by Eq. (20). According to Eq. (20), the second-order rate constant (k2 ) could be calculated: pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ð20Þ N ¼ C Ai k2 DA C IL
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Table 1 Density ðrÞ, viscosity ðmÞ, diffusivity ðDA Þ and solubility HAi of aqueous [Emim]2[Glu] solution for CO2 capture. T (K)
CIL (mol L1) 0.3
0.6
0.9
1.2
1.5
293.15 303.15 313.15 323.15
1.0453 0.0002 1.0423 0.0002 1.0386 0.0001 1.0307 0.0003
1.0631 0.0004 1.0601 0.0002 1.0560 0.0002 1.0520 0.0002
1.0828 0.0002 1.0782 0.0002 1.0746 0.0004 1.0694 0.0001
1.1034 0.0004 1.0989 0.0002 1.0943 0.0001 1.0895 0.0003
1.1225 0.0001 1.1210 0.0002 1.1159 0.0002 1.1108 0.0004
m (mPa S) 293.15 303.15 313.15 323.15
1.337 0.003 1.042 0.003 0.841 0.004 0.693 0.001
1.803 0.003 1.379 0.003 1.095 0.002 0.894 0.004
2.512 0.002 1.877 0.004 1.475 0.003 1.182 0.002
3.701 0.001 2.697 0.005 2.075 0.005 1.637 0.002
5.919 0.004 4.257 0.003 3.157 0.001 2.446 0.001
DA ( 109 m2 s1) 293.15 303.15 313.15 323.15
1.380 0.001 1.832 0.002 2.346 0.003 2.928 0.005
1.087 0.004 1.464 0.005 1.900 0.002 2.388 0.004
0.833 0.003 1.144 0.006 1.497 0.003 1.910 0.005
0.611 0.002 0.856 0.0004 1.139 0.005 1.472 0.003
0.420 0.005 0.594 0.004 0.814 0.002 1.068 0.001
H*Ai (104 mol m3 Pa1) 3.716 0.003 293.15 303.15 2.844 0.001 313.15 2.258 0.005 323.15 1.843 0.003
3.469 0.005 2.659 0.003 2.113 0.001 1.734 0.005
3.234 0.004 2.479 0.005 1.974 0.003 1.620 0.001
3.006 0.003 2.310 0.004 1.841 0.003 1.514 0.001
2.780 0.005 2.152 0.004 1.719 0.003 1.416 0.002
r (g cm1)
Fig. 3. The absorption rate of CO2 into different absorbents (absorption temperature: 303 K; CO2 concentration: 15%; gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
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4. Results and discussion 4.1. Physicochemical data The physicochemical date of the aqueous [Emim]2[Glu] solution with the mole concentrations from 0.3 to 1.5 mol L1 for CO2 capture at 293.15, 303.15, 313.15 and 323.15 K are shown in Table 1. The results showed that the density and viscosity of [Emim]2[Glu] aqueous solution increased as the concentration increased, decreased as the temperature increased. This tendency was in conformity with other similar amino acid ILs reported in literatures [14–16]. However, compared to the monoanion ILs in literature [14,16], such as [C2mim][Gly] and [Hmim][Gly] solution, aqueous [Emim]2[Glu] solution had greater density, viscosity, solubility and diffusivity of CO2 at the same mole concentration and temperature. Moreover, at a fixed temperature, the solubility and diffusivity of CO2 in [Emim]2[Glu] aqueous solution decreased as the mole concentration of IL increased, which was in agreement with the “salting-out” effect [34]. At some sort of fixed [Emim]2[Glu] concentration, the diffusivity of CO2 in [Emim]2[Glu] aqueous solution increased while the solubility of CO2 decreased as the temperature increased, which were in accord with the result in similar literatures [14,15]. 4.2. Absorption of CO2 into aqueous [Emim]2[Glu] solutions 4.2.1. Comparison with different absorbent In order to compare the CO2 capture properties of aqueous [Emim]2[Glu] solutions with different alkanolamine aqueous solutions, the absorption rates of CO2 at 303.15 K were measured in 0.9 mol L1 aqueous solution of [Emim]2[Glu], MEA and DEA, respectively, the results are shown in Fig. 3. According to Fig. 3, the absorption rate in aqueous [Emim]2[Glu] solution is lower than MEA aqueous solution, but higher than DEA aqueous solution at first several minutes. This result suggested that [Emim]2[Glu] is promising to be the candidate for CO2 absorption. 4.2.2. Effect of [Emim]2[Glu] mole concentration on the absorption of CO2 In this part, the absorption rates of CO2 into [Emim]2[Glu] aqueous solution were measured within first 30 min in concentration range from 0.3 to 1.5 mol L1 at 303.15 K. As shown in Fig. 4(a), the absorption rates increased as the [Emim]2[Glu] concentration increased at the same absorption time, and decreased as the absorption time increased. Furthermore, for the lower concentration, the absorption rates of CO2 increased sharply as the concentration increased, but for the higher concentration of [Emim]2[Glu] in solution, the increase tendency of absorption rate becomes slower due to their higher viscosity than that of lower concentration. However, for the low concentration of [Emim]2[Glu] solution, the absorption reached equilibrium quicker owing to the less molecules of [Emim]2[Glu] in the liquid. The relationship between the absorption rate and time could reflect that the absorption rate was controlled by the chemical reaction. The absorption amount obtained by calculating the integral characteristic value of absorption rate to time is shown in Fig. 4(b). The result showed that the absorption amount increased as the [Emim]2[Glu] concentration and the absorption time increased in the first 30 min. The variation tendency of high concentration was quicker than that of low concentration. The absorption load which is equal to the mole ratio of the absorption amount and [Emim]2[Glu] is shown in Fig. 4(c). The results indicated that the CO2 loading decreased as the concentration increased. Therefore, it could be found that increasing [Emim]2[Glu] concentration was conducive to increasing the absorption rate and amount of CO2. However, high concentration which resulted in low CO2 loading will waste
Fig. 4. Absorption rate (a), absorption mount (b) and absorption load (c) of CO2 into aqueous [Emim]2[Glu] with different concentrations. (T = 303 K; PCO2 = 15.20 kPa, gas mixture: 15% CO2–85% N2, gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
absorbent and energy in industry. So, it is necessary to choose an appropriate concentration for CO2 capture. The concentration of 0.9 mol L1 was chosen to detect the kinetic parameters in this paper.
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Fig. 5. Absorption rate (a) and absorption amount (b) of CO2 into 0.9 mol L1 [Emim]2[Glu] aqueous solution with different temperatures. (PCO2 = 15.20 kPa, gas mixture: 15% CO2: 85% N2, gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
4.2.3. Effect of temperature on the absorption In order to obtain the effect of temperature on CO2 absorption, the absorption rate and amount of CO2 into 0.9 mol L1 aqueous [Emim]2[Glu] solution were measured in different temperature, as shown in Fig. 5(a) and (b), respectively. From Fig. 5(a), obviously, the absorption rate of CO2 is influenced by temperature greatly in the investigated temperature range. At the same absorption time, the absorption rate of CO2 increased obviously as the temperature increased. This phenomenon could be explained by their changed
viscosity. However, the absorption rate of CO2 decreased slower at low temperature than that of high temperatures. Fig. 5(b) shows that the absorption amount of CO2 varied slightly at 303.15– 323.15 K, but their values were all more than that at 293.15 K in the first 30 min. This result could be explained that the absorption rate was faster at higher temperature at the beginning. This process was similar to the system on CO2 absorption, such as 15% [N1111][Gly] [15] and 1.0 mol L 1 [C2mim][Gly] [16].
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Fig. 6. Effect of CO2 concentration on absorption rate (a) and absorption amount (b) of CO2 into 0.9 mol L1 [Emim]2[Glu] aqueous solution. (T = 303 K; gas phase stirring speed of 250 rpm, liquid phase stirring speed of 250 rpm).
Table 2 Mass-transfer coefficient of the gas and liquid films. T (K)
kG;CO2 (106 mol m2 s1 Pa1)
kL;CO2 (105 m s1)
KL (105 m s1)
293.15 303.15 313.15 323.15
6.15 6.14 6.13 6.12
1.29 1.59 1.89 2.12
1.29 1.59 1.89 2.12
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Table 3 Effect of liquid phase volume and mass transfer coefficient of liquid phase on absorption rate at 303.15 K. t (min)
VL = 20 mL, N (103 mol m2 s1)
VL = 40 mL, N (103 mol m2 s1)
nL = 130 rpm, N (103 mol m2 s1)
nL = 250 rpm, N (103 mol m2 s1)
0 1 2 5 10 20 30
2.792 0.028 2.755 0.012 2.739 0.021 2.620 0.011 2.264 0.009 1.773 0.015 1.282 0.017
2.811 0.028 2.795 0.024 2.791 0.021 2.774 0.013 2.684 0.017 2.299 0.009 1.898 0.012
2.828 0.023 2.808 0.017 2.784 0.013 2.772 0.008 2.691 0.016 2.318 0.031 1.907 0.017
2.811 0.028 2.795 0.024 2.791 0.021 2.774 0.013 2.684 0.017 2.299 0.009 1.898 0.012
4.2.4. Effect of CO2 concentration on the absorption The CO2 content of exhaust gas was almost among 8 to 15% in industry, thus the effect of CO2 content on the absorption of CO2 into 0.9 mol L1 [Emim]2[Glu] aqueous solution was measured at 303.15 K with CO2 concentration at inlet ranging from 8% (v%) to 15% (v%). As shown in Fig. 6(a) and (b), the absorption rates and amount increased as the CO2 concentration increased in first 30 min. It could be explained that the higher CO2 content had greater driving force between gas phase and gas–liquid interface. 4.3. Kinetics analysis 4.3.1. Mass-transfer coefficient The mass-transfer coefficient of gas and liquid films were needed to check the conditions for different controlling regimes, to calculate the reaction rate parameters. The values of gas and liquid films mass-transfer coefficient were calculated from the experiment data at different temperature. According to the Eqs. (12)– (14), the gas-film and liquid-film mass-transfer coefficients of CO2 in aqueous [Emim]2[Glu] solution at 293.15–323.15 K were obtained, the result is presented in Table 2. According to Table 2, as the temperature increased, the value of gas film mass-transfer coefficient varied slightly, but the liquid film and overall masstransfer coefficient increased greatly. It could be explained that the viscosity of [Emim]2[Glu] aqueous solution decreased as the temperature increased. Moreover, the liquid film mass-transfer coefficient was equal to the overall mass-transfer coefficient. This phenomenon could demonstrate that the reaction was controlled by liquid film. 4.3.2. Absorption kinetics According to the method suggested by Jing and Wu [14–16], the experiments were carried out with gas mixture of 15% CO2 and 85% N2 and the stirring speed for gas phase of 250 rpm at 303.15 K. The absorption rates of CO2 into 0.9 mol L1 [Emim]2[Glu] solution were measured in different liquid phase volumes and liquid stirring speed, the results are presented in Table 3. Obviously, the absorption rate of CO2 increased apparently as the liquid phase volume increased after 5 min for CO2 absorption, but the effect of mass transfer coefficients on the CO2 absorption rates could be
ignored. This phenomenon could be explained that the [Emim]2[Glu] concentration in lower liquid phase volume solution reduced very quickly and the absorption was controlled by gas-film. And according to the discussion in Section 4.2.2, the absorption rates increased as the [Emim]2[Glu] concentration increased, which implied the reaction evidently associated with the concentration of [Emim]2[Glu] other than related to the liquid phase volume and the liquid phase mass transfer coefficient. Consequently, the kinetic region of CO2 absorption into aqueous [Emim]2[Glu] solution was the fast pseudo-first order reaction regime. According to the two-film model and the criterion of the fast pseudo-first order reaction regime described in Section 3.2, the values of enhancement factor E, Hatta number Ha, overall reaction kinetic constant kOV and second-order rate constant k2 could be calculated with the concentration of [Emim]2[Glu] solution ranging from 0.3 to 1.5 mol L1 at 303.15 K by Eqs. (10)–(18), as shown in Table 4. The results showed that the second-order rate constant k2 and the overall reaction kinetic constant kOV increased significantly in 0.3–1.5 mol L1 aqueous [Emim]2[Glu] solution as concentration increased. The enhancement factor E plotted of the square root of concentration is shown in Fig. 7. The result showed that enhancement factor increased as the concentration of [Emim]2[Glu] solution increased, and there was a linear relationship between the enhancement factor and the square root of the [Emim]2[Glu] solution concentration. Moreover, the values of E, Ha, kOV and k2 of CO2 absorption in 0.9 mol L1 aqueous [Emim]2[Glu] solution were also calculated at 293.15–303.15 K and the results are presented in Table 5. Obviously, the second-order reaction rate and overall reaction kinetic constant kOV increased as the temperature increased. The kinetic data can be calculated from an Arrhenius plot of k2 1=T, represented as following: 3665 ð21Þ k2 ¼ 9:253 107 exp T From this equation, the value of activation energy could be calculated to be 30.48 kJ mol1. This value was a little higher than some monoanion ionic liquid, such as [C2mim][Gly] (23.29 kJ mol1) and [Hmim][Gly] (25.36 kJ mol1) reported by Wu
Table 4 The kinetic parameters for CO2 absorption into [Emim]2[Glu] aqueous solution at 303.15 K. CIL (mol L1)
N 103 (mol m2 s1)
PG (Pa)
CAi (mol m3)
E
Ha
kOV (s1)
K2 (L mol1 s1)
0.3 0.6 0.9 1.2 1.5
1.461 0.011 2.153 0.027 2.811 0.028 3.321 0.036 3.557 0.041
12331.253 30.323 10801.245 21.152 9088.853 22.647 7487.918 13.301 7133.28 15.48
4.367 0.011 4.078 0.007 3.752 0.009 3.416 0.004 3.220 0.005
21.043 0.223 33.207 0.486 47.124 0.542 61.137 0.781 69.484 0.897
21.043 0.223 33.207 0.486 47.124 0.542 61.137 0.781 69.484 0.897
61.093 3.163 190.36 6.441 490.59 11.446 1103.53 34.557 2053.59 55.736
67.881 3.745 211.514 7.883 545.093 13.471 1226.142 37.992 2281.769 58.337
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Fig. 7. Enhancement factor of CO2 absorption into aqueous [Emim]2[Glu] solution with different concentrations.
Table 5 The kinetic parameters for CO2 absorption into 0.9 mol L1 [Emim]2[Glu] aqueous solution at different temperature. T (K)
N 103 (mol m2 s1)
PG (Pa)
CAi (mol m3)
E
Ha
kOV (s1)
k2 (L mol1 s1)
293.15 303.15 313.15 323.15
2.425 0.021 2.811 0.028 3.129 0.033 3.227 0.031
10223.692 24.325 9088.853 22.647 8014.808 17.693 7437.255 13.968
4.912 0.012 3.752 0.009 2.976 0.006 2.432 0.005
38.272 0.413 47.124 0.542 55.634 0.624 62.593 0.719
38.272 0.413 47.124 0.542 55.634 0.624 62.593 0.719
292.500 8.764 490.588 11.446 738.644 23.683 921.804 29.543
325.000 11.583 545.097 13.471 820.715 17.679 1024.226 25.994
et al. [16] and Jing et al. [15], respectively. But the activation energy in this paper was much lower than the value of sodium glycinate (63.8 kJ mol1) and MEA (45 kJ mol1) calculated by Lee et al. [35] and Versteeg et al. [36], respectively.
0.9 mol L1 [Emim]2[Glu] was equal to 30.48 kJ mol1, lower than some absorbents reported in literatures.
5. Conclusions
The authors are grateful for the assistance offered by the National Science and Technology Support Program, grant no. 2012BAC02B04.
The density and viscosity of [Emim]2[Glu] aqueous solutions increased as [Emim]2[Glu] concentration increased and decreased as the temperature increased at all the concentrations and temperature investigated. As temperature increased, the diffusivity of [Emim]2[Glu] solutions increased, and the solubility decreased. Furthermore, as the [Emim]2[Glu] concentration increased, the diffusivity and solubility both decreased. The absorption rate and amount of CO2 all increased significantly as the increase of the [Emim]2[Glu] concentration and temperature, but the CO2 load decreased as the increase of the [Emim]2[Glu] concentration. Compared with other absorbents, the absorption rate of CO2 into 0.9 mol L1 [Emim]2[Glu] solution was lower than MEA, but higher than DEA at the same experimental condition. According to the judge method reported in literatures, It was proved that kinetics region of CO2 absorption into aqueous [Emim]2[Glu] was the fast pseudo-first order reaction regime. 1=2
The enhancement factor was linear with C ILs , and the reaction rate could be expressed by the following constant equal:k2 ¼ 9:253 107 expð3665=T Þ. The activation energy of
Acknowledgment
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