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Kinetics and mechanism study of homogeneous reaction of CO2 and blends of diethanolamine and monoethanolamine using the stopped-flow technique
Accepted Manuscript Kinetics and mechanism study of homogeneous reaction of CO2 and blends of diethanolamine and monoethanolamine using the stopped-flow technique Sini Xiao, Helei Liu, Hongxia Gao, Min Xiao, Xiao Luo, Raphael Idem, Paitton Tontiwachwuthikul, Zhiwu Liang PII: DOI: Reference:
S1385-8947(17)30110-9 http://dx.doi.org/10.1016/j.cej.2017.01.100 CEJ 16402
To appear in:
Chemical Engineering Journal
Received Date: Revised Date: Accepted Date:
10 November 2016 21 January 2017 25 January 2017
Please cite this article as: S. Xiao, H. Liu, H. Gao, M. Xiao, X. Luo, R. Idem, P. Tontiwachwuthikul, Z. Liang, Kinetics and mechanism study of homogeneous reaction of CO2 and blends of diethanolamine and monoethanolamine using the stopped-flow technique, Chemical Engineering Journal (2017), doi: http://dx.doi.org/ 10.1016/j.cej.2017.01.100
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Kinetics and mechanism study of homogeneous reaction of CO2 and blends of diethanolamine and monoethanolamine using the stopped-flow technique Sini Xiao1, Helei Liu2*, Hongxia Gao1, Min Xiao1, Xiao Luo1, Raphael Idem1,2, Paitton Tontiwachwuthikul1,2, Zhiwu Liang1,2* 1
Joint International Center for CO2 Capture and Storage (iCCS), Provincial Hunan
Key Laboratory for Cost-effective Utilization of Fossil Fuel Aimed at Reducing Carbon-dioxide Emissions, College ofChemistry and Chemical Engineering, Hunan University, Changsha, 410082, P.R. China 2
Clean Energy Technologies Research Institute(CETRI), Faculty of Engineering and
Applied Science, University of Regina, Regina, Saskatchewan, S4S 0A2, Canada
*CORRESPONDING AUTHOR: E-mail address: [email protected](H. Liu)[email protected] (Z. Liang)
Abstract The stopped-flow technique was applied on the measurement of the kinetics of carbon dioxide (CO2) absorption into aqueous blends of monoethanolamine (MEA) and diethanolamine (DEA) over the temperature range of 293-313K. The investigation of blended MEA+DEA with various molar ratios of DEA to MEA revealed that the reaction mechanism between CO2 and blended absorbents could be different at different ratio of DEA to MEA. Consequently, the kinetic data obtained in this work was split into two groups with respect to the different molar ratios of DEA to MEA in order to study the different mechanisms. In group A, the concentration of MEA was in the range of 5-15 mol/m3 and the concentration of DEA was low (varied between 5-15 mol/m3), while in group B, the DEA concentration was high (varied between 140-240 mol/m3) and the concentration of MEA remained in the range of 515 mol/m3. Modified models based on the termolecular mechanism were developed 1
for each group and used to interpret the experimental kinetic data. Results showed that the models could explain the data well with an AAD of 4.71% for group A at low concentration of DEA and 3.33% for group B at high concentration of DEA. It is interesting to point out that DEA barely reacts with CO2 at low molar ratios (i.e. group A) whereas at high DEA concentration (i.e. group B), both MEA and DEA reacted with CO2. Keywords: carbon capture, mixed amine, kinetics, absorption mechanism, stoppedflow technique
1. Introduction It is generally accepted that continuous CO2 emissions are posing significant global warming issues. Recently, more than 150 countries committed to take responsibility to voluntarily mitigate climate change with the signing of the Paris Agreement. The adopted consensus was to limit the increase of the earth’s surface temperature within 2℃. With this goal, the Intergovernmental Panel on Climate Change (IPCC) studied and evaluated a number of potential scenarios with various combinations of Carbon Capture and Storage (CCS), as well as other approaches including nuclear, solar and integrating fossil energy systems. They concluded that CCS is a robust technology that can cut down carbon emissions and assist in the management of the carbon system. Another conclusion was that without CCS as an option, the costs of limiting global warming to within 2℃ would be significantly higher [1, 2]. Generally, there are three major techniques for carbon capture: pre-combustion, post-combustion, and oxyfuel combustion. So far, post-combustion capture (PCC) is the only CO2 capture technique that has been demonstrated at a full commercial scale. Among the solvents used in PCC (i.e. amines, amino acid and ionic liquids, etc.), amines are considered to be the most widely used and developed due to their flexibility, good performance and moderate toxicity [3]. Some of the industrially important amines are monoethanolamine (MEA) and diethanolamine (DEA) for regenerative chemical absorption processes. MEA is one of the most used primary amines, and it is known for its fast reactivity and low cost, even though it suffers from high corrosiveness and a high regeneration energy requirement [4]. DEA is a 2
secondary amine. Although its reaction rate with CO2 is relatively slow compared with MEA, it has a higher CO2 loading at absorption temperature, lower CO2 loading at regeneration temperature and lower regeneration energy requirement [5, 6]. Recently, a great deal of research has been focused on the use of mixed amine as an alternative absorbent in order to combine the advantages of each amine in a gas treatment process. Gorji et al. [7] studied the selective separation of CO2 from CO2/H2 mixture through facilitated transport membranes with blends of MEA and DEA. The CO2 permeance of 2M aqueous MEA+DEA solution was greater than that of 2M MEA indicating that the selectivity of MEA+DEA mixture was superior to MEA. As demonstrated by Sumon et al. [8] by quantum chemistry computations, primary amines do not have an extra CO2 affinity as do secondary amines. Ismail [9] investigated the CO2 absorption efficiency as well as foaming in mixtures of MEA and DEA amines (15wt% MEA+5wt% DEA, 20wt% MEA+ 10wt% DEA, 25wt % MEA+10wt% DEA), and concluded that 20wt% MEA+10wt% DEA was the optimum blend ratio. Reaction kinetics is one of the vital parameters for CO2 capture. Accurate kinetics data are important for simulation and modeling of both the absorption and desorption processes. A fast reaction rate shortens the column height significantly thereby providing large savings on capital cost. Consequently, it is essential for precise information on reaction kinetics constants as well as a thorough knowledge of reaction mechanisms for the reaction of mixed amines with CO2. Various experimental techniques have been adopted for the investigation of the kinetics of amine-CO2 systems at a laboratory scale. One of the simple and straightforward methods to obtain kinetic data uses the stopped-flow apparatus, which generates the observed pseudo first order reaction rate constant ( ) automatically and instantly with good reproducibility from a program in the microcomputer [10]. It is a direct method that only involves a liquid-liquid process, thus there is no need to evaluate the influence of mass transfer but only to rely on the reliability and authenticity of physical-chemical data, including diffusivity of amine and CO2, and solubility of CO2 in aqueous amine solution, which are sensitive to changes in process variables like composition and temperature [11]. It can also avoid the effect of the reversibility of reactions because the condition of pseudo first order irreversible reaction can always be met [12, 13]. Alper [14] and Liu et al. [15] found the results obtained by the stopped-flow technique agreed well with that obtained from gas absorption studies, 3
including the stirred cell reactor, the wetted wall column, and the laminar jet absorber. Due to its advantages, the stopped-flow technique has been applied successfully to study new absorbent systems such as alkanolamines, amino acid and potassium carbonate [16-18]. This paper investigates the homogeneous reaction of CO2 absorption in aqueous blends of MEA and DEA using the stopped-flow apparatus. The mechanism of CO2 reactions with blends of MEA and DEA was found to be different at the low and high ratios of the two constituent amines. Thus the kinetic data was categorized into two groups according to the different molar ratio of DEA to MEA in order to identify possible mechanisms in different situations. A similar study with DEA was also carried out to confirm the validity of this direct stopped-flow technique.
2. Mechanism 2.1. Reaction Rate for Carbamate Formation The termolecular mechanism was originally proposed by Crooks and Donellan [19] and recently represented by da Silva and Svendsen [20]. Vaidya and Kenig [21] recently reviewed the termolecular kinetic model for various individual and mixed amine systems and suggested that the reaction takes place in a single step, where one molecule of amine reacts simultaneously with one molecule of CO2 and one molecule of base to form carbamate. The intermediate is a loosely-bound encounter complex (Eq. (1)): CO2 + R1 R2 NH ···B → R1 R2 NCOO - ···BH + → R1 R2 NCOO - + BH +
(1)
where R1 is an alkyl while R2 is H for primary amine and an alkyl or an alkanol for secondary amine, and B is a base for deprotonating the amine in solution. The forward reaction rate in aqueous alkanolamine solution can be expressed by:
rCO2 - Am = ∑ k B [ B][ Am][CO2 ]
(2)
where is the reaction rate constant for the reaction between B, amine and CO2, and [B] represents the concentration of the dominating base, mainly water and amine in aqueous solutions [22]. The termolecular mechanism has been successfully applied for prediction of the kinetics of both single and blended amine systems [21]. It covers the situation when the order of the reaction between different amine and CO2 is varying from two to 4
three. When the dominant base is mainly H2O, Eq. (2) can be simplified to:
rCO2 - Am = kW [ H 2O ][ Am][CO2 ]
(3)
where is the reaction rate constant between water, amine and CO2. The concentration of water can be considered to be a constant in Eq. (3) since it is always greater by 3-4 orders of magnitude as compared to the concentration of amine and CO2. The overall reaction is then second order. When the dominant base is mainly the amine, Eq. (1) can be simplified to:
rCO2 - Am = k Am [ Am]2 [CO2 ]
(4)
where is the reaction rate constant between two molecule of amines and one molecule of CO2. The reaction is second order with respect to the amine and first order with respect to CO2. In the MEA-DEA-CO2-H2O system, the overall reaction rate in terms of CO2 and amine, which consists of the contributions of two reactions of CO2 with MEA and DEA, can be represented by:
rCO2 - MEA- DEA = ∑ kB ,M [ BM ][ MEA][CO2 ] + ∑ k B ,D [ BD ][ DEA][CO2 ]
(5)
where BM and BD represent the base for deprotonation of MEA and the base for deprotonation of DEA, and , and , represent the reaction rate constant between base and CO2 with MEA and DEA respectively. 2.2. Parallel Reactions
Besides the reactions for carbamate formation in the chemical absorption process, other reactions take place in parallel, mainly the hydration of CO2 (Eq. (6)) and the direct reaction with hydroxyl ions (Eq. (7)). CO2 + 2 H 2O → HCO3− + H 3O +
(6)
CO2 + OH − → HCO3−
(7)
The value of the rate constant for reaction (6) is given by Pinsent et al. [23] who concluded that compared to the overall reaction rate, its contribution can be negligible without loss of accuracy in aqueous amine solutions. However, the influence of reaction (7) is much more complicated; Astarita et al. [24] gave an estimate of the concentration of the hydroxyl ion for low CO2 loading as:
5
[OH - ] =
KW [ Am] K Am
(8)
and are the dissociation constants of water and amine, respectively. It
was found that hydroxyl ions are consumed at the initial stage, and due to the fact that the amount of hydroxyl ion is small, the contribution of reaction (7) to the overall reaction kinetic is less than 2% [13],[25].Thus based on the termolecular mechanism, the overall reaction rate constant can be given by Eq. (9).
k0 = ∑ kB ,M [ BM ][MEA] + ∑ kB , D [ BD ][ DEA]
(9)
3. Materials and methods 3.1. Materials
Reagent grade MEA and DEA with mass purity of 99% were purchased from Aladdin Industrial Inc. and were used without further purification. Aqueous solutions of DEA (which was used to validate the stopped-flow apparatus) and blends of DEAMEA were prepared to the desired concentrations by adding a weighed quantities amines into deionized water. Commercial-grade CO2 (with a purity of ≥99%) was supplied by Changsha Rizhen Gas Co., China. For each experiment, fresh solution saturated with CO2 was obtained by bubbling the CO2 through deionized water in a jacketed glass stirred reactor. In order to obtain the condition of the pseudo-first order reaction, the saturated CO2 solution was diluted with water to ensure that the molar ratio of amine concentration to CO2 was greater than 10. The concentration of CO2 in the liquid phase was measured with the chromatograph. [15] . 3.2. Stopped-flow technique
The kinetic data for amine-CO2-H2O system in this work were generated from a standard stopped-flow apparatus (Hi-Tech Scientific, UK, Model SF-61DX). It consists of four main parts: a sample handling unit, a conductivity cell, A/D converter, and a microprocessor. Figure 1 shows a schematic drawing of the stopped-flow apparatus. The main section of the sample handling unit not only facilitated the loading and delivery of reactants, but also provided support and enclosure of the sample flow circuit, including the drive syringes, drive valves and observation cell together with interconnecting plumbing. All sample flow circuit components were connected to an external circulator that was thermostatted, and the temperature was 6
controlled within ±0.1 K. For each run, equal volumes of amine solution and CO2 solution were pushed into the observation cell, where the two fresh solutions mixed and started to react. The formation of ion species causes a conductivity change in the cell. The conductivity change as a function of time was measured by a special circuit that gave an output voltage directly proportional to the solution conductivity, as described by Knipe et al. [26]. One sample result (conductance vs. time) from the stopped-flow technique is given in Figure 2. The curve was then fitted to Eq. (10) to estimate the overall reaction rate constant. In Eq. (10), G0 and G is the value of the conductance at the start and the end of the observed reaction respectively, and t (s) is the time. Variation of conductance is then related to CO2 consumption and to k0. G(t) = (G − G )exp (− × t) + G
(10) “Kinetasyst” software installed on the microcomputer can automatically generate the observed pseudo-first-order reaction rate constant ( ). Experiments were repeated at least six times at each temperature for all concentrations. Given that solutions used in the stopped-flow apparatus are usually at relatively low concentrations, they are close to infinite dilution and can be treated as ideal solutions. Under such low concentration and low ionic strength conditions, the effect of ionic strength can be considered to be negligible without resulting in a loss of accuracy [16]. In stopped-flow apparatus, a liquid with dissolved CO2 and another liquid with dissolved amine will mix simultaneously and instantly in a batch reactor (which is the conductivity cell) where the homogeneous reaction of CO2 and amine takes place. Therefore it is not necessary to study the residence time distribution for the fluid in the stopped-flow apparatus due to the fact that the residence times of all the fluid elements are the same in a batch reactor. In addition, the reactor cell is enclosed and filled with liquid, so there is no gas phase involved in the reactor when using stoppedflow technique as in this work. All these factors help to make the investigation of the intrinsic kinetics straightforward.
Figure 1 Schematic drawing of the experimental stopped-flow equipment.
Figure 2 A typical result of the stopped-flow investigation in six repeated 7
experiments. (T=303 K, [DEA]=7.5 mol/m3 and [MEA]=5 mol/m3 ).
4. Results and discussion
4.1. CO2-Aqueous DEA
The kinetics of aqueous DEA solutions (over a concentration range of 140-500 mol/m3) reacting with CO2 were measured at temperatures in the range of 293-313 K, and the values were then compared to corresponding published data from the literature. As Figure 3 shows, results obtained in this work were in good agreement with those presented by Siemieniec et al. [27] and Ali [28], using the same technique for aqueous DEA system. Furthermore, Siemieniec et al. [27] and Ali [28] also verified that values obtained for DEA were in accordance with those of other workers using the stopped-flow technique, as well as rapid mixing, continuously stirred cell reactor and wetted wall methods. Therefore, results obtained in this work were deemed reliable and valid. The expression of can be given by Eq. (11), which is derived from Eq. (2). A nonlinear fitting method was used for estimation of parameters at each temperature. A subroutine called NLINFIT in MatLab was applied for finding the optimum value for each parameter (similarly hereinafter). The obtained values of and are listed in Table 1, representing DEA and H2O as Base in the deprotonation process, respectively. The correlated were found to be in good accordance with the experimental , with an average absolute deviation (AAD%) of 2.32% indicating the termolecular mechanism can explain the CO2-aqueous DEA system well. Parity plot of predicted and experimental in CO2-aqueous DEA system is shown in Figure 4. k0 = (k D [ DEA] + kW [ H 2O])[ DEA]
(11)
Table 1 Fitted value of and based on the termolecular mechanism. Figure 3 Comparison of the values for DEA obtained in this work and from literature.
8
Figure 4 Parity plot of predicted and experimental in CO2-aqueous DEA system.
4.2. CO2-Aqueous MEA+DEA
The kinetics of CO2 and mixed alkanolamines, MEA and DEA, were studied using the stopped-flow technique at 293-313 K, with intervals of 5 K. The concentration of MEA was in the range of 5-15 mol/m3 and the concentration of DEA was in the range of 5-200 mol/m3, concentrations which were chosen in order to not only satisfy the pseudo-first regime, but also to cover various concentration ratios of DEA to MEA. As Figure 5and Figure 6 show, the experimentally obtained were plotted versus concentrations of MEA and DEA for all temperatures (the original data and standard deviation are given in Supportive Information). For all runs, the CO2 reaction rate in aqueous blend of MEA and DEA increased disproportionally as the concentration of one of the amines increased. Moreover, temperature accelerated the overall rate constant in a nonlinear form. The obtained for different temperatures and concentrations were fitted to the rate equation using Eq. (9), as Bosch et al. [29] suggested. However, no meaningful values were found. This result implies a different reaction mechanism in the MEADEA-CO2-H2O system as compared to that reported in the literature on blended amine-CO2 systems, in which all constituent alkanolamines reacting with CO2 were taken into account. According to Li et al. [30], there is a certain critical point in the reaction process below which the DEA carbamate cannot be detected. As well, the critical loading can decrease as the molar ratio of DEA to MEA increases. As shown in Figure 7, when the molar ratio of MEA and DEA is 1:1, the critical loading is 0.17 (mol CO2/mol amine).Thus, the kinetic data were categorized into two groups, low molar ratio and high molar ratio, as shown in Figure 5 and Figure 6, to understand the possible mechanism of reactions between CO2 and blended amines. For the low molar ratio, named as Group A, the concentration of MEA and DEA was 5-15 mol/m3. For the high molar ratio, named as Group B, the concentration of DEA was 140-240 mol/m3 while MEA remained at 10-20 mol/m3. Thus, the molar ratio of DEA to MEA in Group A was around 1:1 while that in Group B was greater than 10:1.
9
Figure 5 of CO2 absorption into the MEA-DEA-H2O system of low concentration of DEA (Group A).
Figure 6 of CO2 absorption into the MEA-DEA-H2O system of high concentration of DEA (Group B).
Figure 7 Ion speciation plots of MEA-DEA-CO2-H2O system at 298 K [30]. 4.2.1. Group A All experimental in Group A were fitted to Eq. (9). However, the experimental could not be fitted successfully to Eq. (9) whether or not the MEA for deprotonation term was included. The parameters in Eq. (9) were found to be meaningless at any investigated temperature. Therefore, a new model is proposed in this work assuming that when the CO2 loading was below 0.1 and the molar ratio of DEA to MEA was around 1:1, only MEA carbamate was formed and DEA simply participated in the deprotonation process, as presented by Eq. (12):
k 0A = (kMA , M [ MEA] + k DA, M [ DEA] + kWA ,M [ H 2O])[ MEA]
(12)
where , , , and , represent the reaction rate constant for one molecule of
MEA, DEA and water depronating one molecule of MEA in Group A respectively. Furthermore, during the fitting of the experimental data, it was observed that MEA had a very small effect on the deprotonation process, and thus, was ignored for model simplification. Then Eq. (12) became:
k 0A = (kDA,M [ DEA] + kWA ,M [ H 2O])[ MEA]
(13)
The , and , in Eq. (13) were obtained and plotted in Figure 8. They can
be expressed as a function of temperature as shown in Eq. (14) and Eq. (15), respectively. As can be seen in Figure 9, the predicted values of calculated from Eq. (13) were observed to be in good agreement with the experimental results, with an AAD of 4.71%. 5650 k DA, M = (1.72 ×107 ) exp − T
(14)
4214 kWA ,M = 59.2exp − T
(15)
Based on the results of the two models, it can be stated that the assumption for 10
Eq. (13) was reasonable. It can be attributed to the nature of the amine. MEA is a primary amine, and DEA is a secondary amine, which can be affected by the detrimental effect of the steric factor. Another reason can be the stronger basicity of MEA than DEA (the values of pKa for MEA and DEA are 9.40 and 8.88 respectively [31]), leading to a stronger possibility of the occurrence of the reaction between MEA and CO2 as compared to the reaction between DEA and CO2. Therefore, at the molar ratio of DEA to MEA around 1:1 (i.e. Group A), the relationship of MEA and DEA appeared generally to be competitive in terms of their direct reactions with CO2 even though DEA also performs as base in MEA deprotonation. Figure 8 Fitted value of , and , in Eq. (13).
Figure 9 Parity plot of predicted and experimental in Group A. 4.2.2. Group B The experimental in group B were fitted to Eq. (13) at first so as to verify if the situation in this group was the same as that in group A. The fitting values obtained were unreasonable because they were either negative or infinite whether or not the term involving MEA for deprotonation was included in the fitting. This result showed that this mechanism was not suitable for conditions in group B. Also, Eq. (9) was used to fit the experimental . It is interesting to find that the inclusion of MEA as base for deprotonation and H2O as base for deprotonation of DEA in Eq. (9) did not result in any meaningful results. Hence, it was assumed that for the case of MEA, its concentration was too small, leading to miniscule probability of collision of MEA with amine-CO2 complex. This is similar to group A, since the same assumption was made. The possible reason of the case of DEA is that the deprotonation reaction rate of DEA by H2O is far slower given the short reaction time used in the stopped-flow technique, in which a slow reaction is not expected to have an effect on the overall reaction rate [32, 33]. Consequently, Eq. (9) was modified to obtain the case where MEA for deprotonation and H2O for deprotonation of DEA were left out, as shown in Eq.(16).
k 0B = (kDB,M [ DEA] + kWB ,M [ H 2O])[MEA] + kDB, D [ DEA]2
(16)
The value of , , , and , in Eq. (16) were obtained and were plotted as
a function of the temperature (Figure 10), which can be expressed using Eq.(17)-(19). 11
As can be seen in the Figure 11, the predicted values of calculated from Eq. (16) were found to be in a good agreement with experimental results, with an AAD of 3.33%. In group B, MEA and DEA both react with CO2 on CO2 absorption in blended amine as described by other works [25, 34, 35]. 5276 k DB,M = (5.82 × 105 ) exp − T
(17)
3590 kWB ,M = 6.79exp − T
(18)
3165 k DB, D = 43.268exp − T
(19)
Figure 10 Fitted value of , , , and , in Eq. (16).
Figure 11 Parity plot of predicted and experimental in Group B. The fitted model (Eq. (13) and Eq. (16)) shows that the reaction mechanism was different between Groups A and B, and implies that it was MEA which reacted with CO2 first, or at the initial region the DEA carbamate formation was insignificant, and so, could be ignored. When the amount of CO2 or the molar ratio of DEA to MEA reaches a certain level, DEA starts to react with CO2 and a more complicated situation in terms of the amine mixture of the CO2 absorption process exists, giving rise to the occurrence of both reactions in solution. The condition of Group A was below the critical point, where the competition of MEA and DEA was obvious given MEA had an overwhelming inhibitory effect to DEA; nevertheless the inferior position of DEA was made up by the high concentration of DEA, i.e. Group B, where both MEA and DEA generally played significant roles. It should be noted that the competition and cooperation of MEA and DEA were present in both regions. This mechanism compares favorably with the NMR results and explains experimental kinetic results in this work successfully as discussed above.
4.3. Comparison of Reaction Kinetics of Aqueous MEA+DEA and Pure Aqueous DEA Due to DEA having little influence in the reactions of Group A, the comparison 12
of aqueous MEA+DEA and pure aqueous DEA here was restricted to the comparison of the addition of MEA to DEA (i.e. Group B) and its blank system (pure DEA). As the experimental data showed, the addition of small amounts of MEA to DEA resulted in a significant enhancement of the CO2 reaction rate. For instance, the overall reaction rate constant of aqueous system increased by 131% with addition of 0.01 M MEA in 0.16 M DEA, at 298 K ( at 0.16 M DEA: 29.31 s-1, and at 0.16 M DEA + 0.01 M MEA: 67.34 s-1). It should be noted that , compares favorably to , as Figure 10 shows, since
these two parameters represent the same reaction step in different environments, the former being kinetic constant for reaction of two molecules of DEA and one molecule of CO2 in blends of high ratio of DEA to MEA while the latter being that in pure DEA solution. 5. Conclusions
This paper has presented the CO2 absorption kinetics of mixed alkanolamines, MEA and DEA, using the stopped-flow technique at the temperature range of 293-313 K. The kinetic data were categorized into two groups, Group A and Group B. In Group A, the concentrations of both MEA and DEA were in the range of 5-15 mol/m3. In Group B, the concentration of DEA was in the range of 140-240 mol/m3, while the concentration of MEA remained the same at 10-20 mol/m3. The possible reaction mechanism between CO2 and blended amines was proposed. Two models based on termolecular mechanism were adopted to fit the experimental in both groups in order to explain the kinetic data. The results showed that the predicted exhibited good agreement with the experimental data with an AAD of 4.71% for Group A and an AAD of 3.33% for Group B. Comparison of each optimal model for Group A and Group B suggests that below a certain critical point (i.e. Group A), MEA had an overwhelming inhibitory effect to DEA and the contribution to was mainly from the reaction of MEA and CO2, since the DEA carbamate formation was nonexistent or insignificant; while above this point (i.e. Group B), the overall reaction of CO2 with aqueous blends of MEA and DEA could be regarded as a reaction between CO2 and MEA in parallel with the reaction of CO2 with DEA.
Acknowledgements 13
The financial support from the National Natural Science Foundation of China (NSFCNos. 21476064, U1362112 and 21376067), National Key Technology R&D Program (MOST-Nos.2012BAC26B01 and 2014BAC18B04), Innovative Research Team Development Plan (MOE-No.IRT1238), Specialized Research Fund for the Doctoral Program of Higher Education (MOE-No. 20130161110025), China’s State “Project 985” in Hunan University-Novel Technology Research & Development for CO2 Capture, Key Project of International & Regional Cooperation of Hunan Provincial Science and Technology plan (2014WK2037), and China Outstanding Engineer Training Plan for Students of Chemical Engineering & Technology in Hunan University (MOE-No.2011-40) is gratefully acknowledged. We also greatly appreciate Mr. Wilfred Olson for his great contribution to help us correct any grammar mistakes and insightful inputs to our research work.
14
Nomenclature
AAD
absolute average deviation
Am
amine
B
base (amine, water or hydroxyl ion)
DEA
diethanolamine
observed pseudo first order reaction rate constant (s-1)
, , , , ,
second order kinetic rate constant between amine and CO2 with amine as base (m6/mol2 s2) second order kinetic rate constant between base, amine and CO2 (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with DEA as base (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with DEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with MEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between MEA and CO2 with DEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between MEA and CO2 with MEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with water as base (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with water as base in blended amine (m6/mol2 s2)
dissociation constant of amine (mol/m3)
dissociation constant of water (mol/m3)
M
molarity (mol/m3)
MEA
monoethanolamine
T
temperature (K)
[]
concentration (mol/m3)
15
Reference [1] D.C. Miller, J.T. Litynski, L.A. Brickett, B.D. Morreale, Toward transformational carbon capture systems, AIChE J. 62 (2016) 2-10. [2] M.G. Plaza, S. García, F. Rubiera, J.J. Pis, C. Pevida, Post-combustion CO2 capture with a commercial activated carbon: Comparison of different regeneration strategies, Chem. Eng. J. 163 (2010) 41-47. [3] Z. Liang, W. Rongwong, H. Liu, K. Fu, H. Gao, F. Cao, R. Zhang, T. Sema, A. Henni, K. Sumon, D. Nath, D. Gelowitz, W. Srisang, C. Saiwan, A. Benamor, M. Al-Marri, H. Shi, T. Supap, C. Chan, Q. Zhou, M. Abu-Zahra, M. Wilson, W. Olson, R. Idem, P. Tontiwachwuthikul, Recent progress and new developments in post-combustion carboncapture technology with amine based solvents, Int. J. Greenhouse Gas Control 40 (2015) 2654. [4] L. Raynal, P.-A. Bouillon, A. Gomez, P. Broutin, From MEA to demixing solvents and future steps, a roadmap for lowering the cost of post-combustion carbon capture, Chem. Eng. J. 171 (2011) 742-752. [5] W. Fan, Y. Liu, K. Wang, Detailed experimental study on the performance of monoethanolamine, diethanolamine, and diethylenetriamine at absorption/regeneration conditions, Journal of Cleaner Production, 125 (2016) 296-308. [6] Y.E. Kim, A.L. Jin, S.K. Jeong, Y.I. Yoon, S.T. Bae, S.C. Nam, Y.E. Kim, A.L. Jin, S.K. Jeong, Y.I. Yoon, Comparison of carbon dioxide absorption in aqueous MEA, DEA, TEA, and AMP solutions, Bulletin- Korean Chem. Soc.34 (2013) 783-787. [7] A.H. Gorji, T. Kaghazchi, A. Kargari, Selective removal of carbon dioxide from wt CO2 /H2 mxtures via facilitated transport membranes containing amine blends as carriers, Chem. Eng. Technol. 32 (2008) 120-128. [8] K.Z. Sumon, C.H. Bains, D.J. Markewich, A. Henni, A.L. East, Semicontinuum aolvation modeling improves predictions of carbamate stability in the CO2 +aqueous amine reaction, J. Phys. Chem. B, 119 (2015) 12256-12264. [9] M. Ismail, Study of carbon dioxide (CO2) absorption in single MEA and mixed MEA&DEA, University Teknologi Petronas, (2005). [10] J. Li, H. Liu, Z. Liang, X. Luo, H. Liao, R. Idem, P. Tontiwachwuthikul, Experimental study of the kinetics of the homogenous reaction of CO2 into a novel aqueous 3-diethylamino-1,2propanediol solution using the stopped-flow technique, Chem. Eng. J. 270 (2015) 485-495. [11] P.D. Vaidya, E.Y. Kenig, A study on CO2 absorption kinetics by aqueous sSolutions of N,NDiethylethanolamine and N-Ethylethanolamine, Chem. Eng. Technol. 32 (2009) 556-563. [12] G.F. Versteeg, L.A.J. Van Dijck, W.P.M. Van Swaaij, On the kinetics between CO2 and alkanolamines both in aqueous and non-aqueous solutions. An overview, Chem. Eng. Commun. 144 (1996) 113-158. [13] H. Kierzkowska-Pawlak, M. Siemieniec, A. Chacuk, Reaction kinetics of CO2 in aqueous 16
methyldiethanolamine solutions using the stopped-flow technique, Chem. Process Eng. 33 (2012) 7-18. [14] E. Alper, Reaction mechanism and kinetics of aqueous solutions of 2-amino-2-methyl-1propanol and carbon dioxide, Ind. Eng. Chem. Res. 29 (2002) 1725-1728. [15] H. Liu, T. Sema, Z. Liang, K. Fu, R. Idem, Y. Na, P. Tontiwachwuthikul, CO2 absorption kinetics of 4-diethylamine-2-butanol solvent using stopped-flow technique, Sep. Purif. Technol. 136 (2014) 81-87. [16] G. Hu, K.H. Smith, L. Liu, S.E. Kentish, G.W. Stevens, Reaction kinetics and mechanism between histidine and carbon dioxide, Chem. Eng. J. 307 (2017) 56-62. [17] G. Hu, K.H. Smith, N.J. Nicholas, J. Yong, S.E. Kentish, G.W. Stevens, Enzymatic carbon dioxide capture using a thermally stable carbonic anhydrase as a promoter in potassium carbonate solvents, Chem. Eng. J. 307 (2017) 49-55. [18] H. Liu, M. Xiao, Z. Liang, W. Rongwong, J. Li, P. Tontiwachwuthikul, Analysis of reaction kinetics of CO2 absorption into a novel 1-(2-hydroxyethyl)-piperidine solvent using stoppedflow technique, Ind. Eng. Chem. Res. 54 (2015) 12525-12533. [19] J.E. Crooks, J.P. Donnellan, Kinetics and mechanism of the reaction between carbon dioxide and amines in aqueous solution, J. Chem. Soc. Perkin Trans. 2 (1989) 331-333. [20] E.F. da Silva, H.F. Svendsen, Ab initio study of the reaction of carbamate formation from CO2 and alkanolamines, Ind. Eng. Chem. Res. 43 (2004) 3413-3418. [21] P.D. Vaidya, E.Y. Kenig, Termolecular Kinetic Model for CO2-Alkanolamine Reactions: An Overview, Chem. Eng. Technol. 33 (2010) 1577-1581. [22] G. Versteeg, W. Van Swaaij, On the kinetics between CO2 and alkanolamines both in aqueous and non-aqueous solutions -----Ⅰ. Primary and secondary amines, Chem.Eng. Sci. 43 (1988) 573-585. [23] B.R.W. Pinsent, L. Pearson, F.J.W. Roughton, The kinetics of combination of carbon dioxide with hydroxide ions, Trans. Faraday Soc.52 (1956) 1512-1520. [24] G. Astarita, D. Savage, A. Bisio, Gas treatment with chemical solvents, John Wiley & Sons, New York, 1983. [25] C.-H. Liao, M.-H. Li, Kinetics of absorption of carbon dioxide into aqueous solutions of monoethanolamine+N-methyldiethanolamine, Chem. Eng. Sci. 57 (2002) 4569-4582. [26] A. Knipe, D. McLean, R. Tranter, A fast response conductivity amplifier for chemical kinetics, J. Phys. E 7 (1974) 586-590. [27] M. Siemieniec, H. Kierzkowska-Pawlak, A. Chacuk, Reaction kinetics of carbon dioxide in aqueous diethanolamine solutions using the stopped-flow technique, Ecol. Chem. Eng. S 19 (2012) 55-66. [28] S.H. Ali, Kinetic study of the reaction of diethanolamine with carbon dioxide in aqueous and mixed solvent systems—application to acid gas cleaning, Sep. Purif. Technol. 38 (2004) 281296. 17
[29] H. Bosch, G.F. Versteeg, W.P.M.V. Swaaij, Gas-liquid mass transfer with parallel reversible reactions——III. Absorption of CO2 into solutions of blends of amines, Chem. Eng. Sci. 44 (1989) 2723-2734. [30] M. Li, H. Liu, Z. Liang, R. Idem, P. Tontiwachwuthikul, The research of coordinative and competitive relationship of CO2 absorption into MEA and DEA in blended aqueous amines, http://www3.aiche.org/proceedings/Abstract.aspx?PaperID=426569, (2016). [31] M.R. Simond, K. Ballerat-Busserolles, Y. Coulier, L. Rodier, J.Y. Coxam, Dissociation Constants of Protonated Amines in Water at Temperatures from 293.15 K to 343.15 K, J. Soln. Chem. 41 (2012) 130-142. [32] N. Ramachandran, A. Aboudheir, R. Idem, P. Tontiwachwuthikul, Kinetics of the absorption of CO2 into mixed aqueous loaded solutions of monoethanolamine and methyldiethanolamine, Ind. Eng. Chem. Res. 45 (2006) 2608-2616. [33] E.B. Rinker, S.A. Sami, O.C. Sandall, Kinetics and modeling of carbon dioxide absorption into aqueous solutions of diethanolamine, Ind. Eng. Chem. Res. 35 (1996) 1107-1114. [34] S.H. Ali, O. Al-Rashed, S.Q. Merchant, Opportunities for faster carbon dioxide removal: A kinetic study on the blending of methyl monoethanolamine and morpholine with 2-amino-2methyl-1-propanol, Sep. Purif. Technol. 74 (2010) 64-72. [35] S. Paul, A.K. Ghoshal, B. Mandal, Theoretical studies on separation of CO2 by single and blended aqueous alkanolamine solvents in flat sheet membrane contactor (FSMC), Chem. Eng. J. 144 (2008) 352-360.
18
Figure Captions Figure 1 Schematic drawing of the experimental stopped-flow equipment.
Figure 2 A typical result of the stopped-flow investigation in six repeated experiments. (T=303 K, [DEA]=7.5 mol/m3 and [MEA]=5 mol/m3 ) Figure 3 Comparison of the values for DEA obtained in this work and from literature. Figure 4 Parity plot of predicted and experimental in CO2-aqueous DEA system. Figure 5 of CO2 absorption into the MEA-DEA-H2O system of low concentration of DEA (Group A). Figure 6 of CO2 absorption into the MEA-DEA-H2O system of high concentration of DEA (Group B). Figure 7 Ion speciation plots of MEA-DEA-CO2-H2O system at 298 K [30]. Figure 8 Fitted value of , and , in Eq. (13).
Figure 9 Parity plot of predicted and experimental in Group A. Figure 10 Fitted value of , , , and , in Eq. (16).
Figure 11 Parity plot of predicted and experimental in Group B.
19
Figure 1 Schematic drawing of the experimental stopped-flow equipment.
Figure 2 A typical result of the stopped-flow investigation in six repeated experiments. (T=303 K, [DEA]=7.5 mol/m3 and [MEA]=5 mol/m3 )
Figure 3 Comparison of the values for DEA obtained in this work and from literature. 20
Figure 4 Parity plot of predicted and experimental in CO2-aqueous DEA system.
Figure 5 of CO2 absorption into the MEA-DEA-H2O system of low concentration of DEA (Group A).
Figure 6 of CO2 absorption into the MEA-DEA-H2O system of high concentration of DEA (Group B). 21
Figure 7 Ion speciation plots of MEA-DEA-CO2-H2O system at 298 K [30].
Figure 8 Fitted value of , and , in Eq. (13).
Figure 9 Parity plot of predicted and experimental in Group A.
22
Figure 10 Fitted value of , , , and , in Eq. (16).
Figure 11 Parity plot of predicted and experimental in Group B.
23
Table Captions
Table 1 Fitted value of and based on the termolecular mechanism.
24
Table 1 Fitted value of and based on the termolecular mechanism. kD
kW
(104 m6/mol2 s2)
(106 m6/mol2 s2)
293
5.44±0.15
1.95±0.11
298
6.31±0.27
2.70±0.20
303
7.91±0.25
3.20±0.19
308
9.52±0.44
3.94±0.33
313
12.16±0.37
4.69±0.27
T (K)
25
Highlights: •
Reaction kinetics between CO2 and blends of MEA and DEA are studied.
•
Two possible situations are discussed based on the ratio of blends.
•
DEA barely reacts with CO2 in blends at low concentration.
•
The proposed mechanisms explain all experimental data very well with excellent AADs less than 5%.
26
S1385-8947(17)30110-9 http://dx.doi.org/10.1016/j.cej.2017.01.100 CEJ 16402
To appear in:
Chemical Engineering Journal
Received Date: Revised Date: Accepted Date:
10 November 2016 21 January 2017 25 January 2017
Please cite this article as: S. Xiao, H. Liu, H. Gao, M. Xiao, X. Luo, R. Idem, P. Tontiwachwuthikul, Z. Liang, Kinetics and mechanism study of homogeneous reaction of CO2 and blends of diethanolamine and monoethanolamine using the stopped-flow technique, Chemical Engineering Journal (2017), doi: http://dx.doi.org/ 10.1016/j.cej.2017.01.100
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Kinetics and mechanism study of homogeneous reaction of CO2 and blends of diethanolamine and monoethanolamine using the stopped-flow technique Sini Xiao1, Helei Liu2*, Hongxia Gao1, Min Xiao1, Xiao Luo1, Raphael Idem1,2, Paitton Tontiwachwuthikul1,2, Zhiwu Liang1,2* 1
Joint International Center for CO2 Capture and Storage (iCCS), Provincial Hunan
Key Laboratory for Cost-effective Utilization of Fossil Fuel Aimed at Reducing Carbon-dioxide Emissions, College ofChemistry and Chemical Engineering, Hunan University, Changsha, 410082, P.R. China 2
Clean Energy Technologies Research Institute(CETRI), Faculty of Engineering and
Applied Science, University of Regina, Regina, Saskatchewan, S4S 0A2, Canada
*CORRESPONDING AUTHOR: E-mail address: [email protected](H. Liu)[email protected] (Z. Liang)
Abstract The stopped-flow technique was applied on the measurement of the kinetics of carbon dioxide (CO2) absorption into aqueous blends of monoethanolamine (MEA) and diethanolamine (DEA) over the temperature range of 293-313K. The investigation of blended MEA+DEA with various molar ratios of DEA to MEA revealed that the reaction mechanism between CO2 and blended absorbents could be different at different ratio of DEA to MEA. Consequently, the kinetic data obtained in this work was split into two groups with respect to the different molar ratios of DEA to MEA in order to study the different mechanisms. In group A, the concentration of MEA was in the range of 5-15 mol/m3 and the concentration of DEA was low (varied between 5-15 mol/m3), while in group B, the DEA concentration was high (varied between 140-240 mol/m3) and the concentration of MEA remained in the range of 515 mol/m3. Modified models based on the termolecular mechanism were developed 1
for each group and used to interpret the experimental kinetic data. Results showed that the models could explain the data well with an AAD of 4.71% for group A at low concentration of DEA and 3.33% for group B at high concentration of DEA. It is interesting to point out that DEA barely reacts with CO2 at low molar ratios (i.e. group A) whereas at high DEA concentration (i.e. group B), both MEA and DEA reacted with CO2. Keywords: carbon capture, mixed amine, kinetics, absorption mechanism, stoppedflow technique
1. Introduction It is generally accepted that continuous CO2 emissions are posing significant global warming issues. Recently, more than 150 countries committed to take responsibility to voluntarily mitigate climate change with the signing of the Paris Agreement. The adopted consensus was to limit the increase of the earth’s surface temperature within 2℃. With this goal, the Intergovernmental Panel on Climate Change (IPCC) studied and evaluated a number of potential scenarios with various combinations of Carbon Capture and Storage (CCS), as well as other approaches including nuclear, solar and integrating fossil energy systems. They concluded that CCS is a robust technology that can cut down carbon emissions and assist in the management of the carbon system. Another conclusion was that without CCS as an option, the costs of limiting global warming to within 2℃ would be significantly higher [1, 2]. Generally, there are three major techniques for carbon capture: pre-combustion, post-combustion, and oxyfuel combustion. So far, post-combustion capture (PCC) is the only CO2 capture technique that has been demonstrated at a full commercial scale. Among the solvents used in PCC (i.e. amines, amino acid and ionic liquids, etc.), amines are considered to be the most widely used and developed due to their flexibility, good performance and moderate toxicity [3]. Some of the industrially important amines are monoethanolamine (MEA) and diethanolamine (DEA) for regenerative chemical absorption processes. MEA is one of the most used primary amines, and it is known for its fast reactivity and low cost, even though it suffers from high corrosiveness and a high regeneration energy requirement [4]. DEA is a 2
secondary amine. Although its reaction rate with CO2 is relatively slow compared with MEA, it has a higher CO2 loading at absorption temperature, lower CO2 loading at regeneration temperature and lower regeneration energy requirement [5, 6]. Recently, a great deal of research has been focused on the use of mixed amine as an alternative absorbent in order to combine the advantages of each amine in a gas treatment process. Gorji et al. [7] studied the selective separation of CO2 from CO2/H2 mixture through facilitated transport membranes with blends of MEA and DEA. The CO2 permeance of 2M aqueous MEA+DEA solution was greater than that of 2M MEA indicating that the selectivity of MEA+DEA mixture was superior to MEA. As demonstrated by Sumon et al. [8] by quantum chemistry computations, primary amines do not have an extra CO2 affinity as do secondary amines. Ismail [9] investigated the CO2 absorption efficiency as well as foaming in mixtures of MEA and DEA amines (15wt% MEA+5wt% DEA, 20wt% MEA+ 10wt% DEA, 25wt % MEA+10wt% DEA), and concluded that 20wt% MEA+10wt% DEA was the optimum blend ratio. Reaction kinetics is one of the vital parameters for CO2 capture. Accurate kinetics data are important for simulation and modeling of both the absorption and desorption processes. A fast reaction rate shortens the column height significantly thereby providing large savings on capital cost. Consequently, it is essential for precise information on reaction kinetics constants as well as a thorough knowledge of reaction mechanisms for the reaction of mixed amines with CO2. Various experimental techniques have been adopted for the investigation of the kinetics of amine-CO2 systems at a laboratory scale. One of the simple and straightforward methods to obtain kinetic data uses the stopped-flow apparatus, which generates the observed pseudo first order reaction rate constant ( ) automatically and instantly with good reproducibility from a program in the microcomputer [10]. It is a direct method that only involves a liquid-liquid process, thus there is no need to evaluate the influence of mass transfer but only to rely on the reliability and authenticity of physical-chemical data, including diffusivity of amine and CO2, and solubility of CO2 in aqueous amine solution, which are sensitive to changes in process variables like composition and temperature [11]. It can also avoid the effect of the reversibility of reactions because the condition of pseudo first order irreversible reaction can always be met [12, 13]. Alper [14] and Liu et al. [15] found the results obtained by the stopped-flow technique agreed well with that obtained from gas absorption studies, 3
including the stirred cell reactor, the wetted wall column, and the laminar jet absorber. Due to its advantages, the stopped-flow technique has been applied successfully to study new absorbent systems such as alkanolamines, amino acid and potassium carbonate [16-18]. This paper investigates the homogeneous reaction of CO2 absorption in aqueous blends of MEA and DEA using the stopped-flow apparatus. The mechanism of CO2 reactions with blends of MEA and DEA was found to be different at the low and high ratios of the two constituent amines. Thus the kinetic data was categorized into two groups according to the different molar ratio of DEA to MEA in order to identify possible mechanisms in different situations. A similar study with DEA was also carried out to confirm the validity of this direct stopped-flow technique.
2. Mechanism 2.1. Reaction Rate for Carbamate Formation The termolecular mechanism was originally proposed by Crooks and Donellan [19] and recently represented by da Silva and Svendsen [20]. Vaidya and Kenig [21] recently reviewed the termolecular kinetic model for various individual and mixed amine systems and suggested that the reaction takes place in a single step, where one molecule of amine reacts simultaneously with one molecule of CO2 and one molecule of base to form carbamate. The intermediate is a loosely-bound encounter complex (Eq. (1)): CO2 + R1 R2 NH ···B → R1 R2 NCOO - ···BH + → R1 R2 NCOO - + BH +
(1)
where R1 is an alkyl while R2 is H for primary amine and an alkyl or an alkanol for secondary amine, and B is a base for deprotonating the amine in solution. The forward reaction rate in aqueous alkanolamine solution can be expressed by:
rCO2 - Am = ∑ k B [ B][ Am][CO2 ]
(2)
where is the reaction rate constant for the reaction between B, amine and CO2, and [B] represents the concentration of the dominating base, mainly water and amine in aqueous solutions [22]. The termolecular mechanism has been successfully applied for prediction of the kinetics of both single and blended amine systems [21]. It covers the situation when the order of the reaction between different amine and CO2 is varying from two to 4
three. When the dominant base is mainly H2O, Eq. (2) can be simplified to:
rCO2 - Am = kW [ H 2O ][ Am][CO2 ]
(3)
where is the reaction rate constant between water, amine and CO2. The concentration of water can be considered to be a constant in Eq. (3) since it is always greater by 3-4 orders of magnitude as compared to the concentration of amine and CO2. The overall reaction is then second order. When the dominant base is mainly the amine, Eq. (1) can be simplified to:
rCO2 - Am = k Am [ Am]2 [CO2 ]
(4)
where is the reaction rate constant between two molecule of amines and one molecule of CO2. The reaction is second order with respect to the amine and first order with respect to CO2. In the MEA-DEA-CO2-H2O system, the overall reaction rate in terms of CO2 and amine, which consists of the contributions of two reactions of CO2 with MEA and DEA, can be represented by:
rCO2 - MEA- DEA = ∑ kB ,M [ BM ][ MEA][CO2 ] + ∑ k B ,D [ BD ][ DEA][CO2 ]
(5)
where BM and BD represent the base for deprotonation of MEA and the base for deprotonation of DEA, and , and , represent the reaction rate constant between base and CO2 with MEA and DEA respectively. 2.2. Parallel Reactions
Besides the reactions for carbamate formation in the chemical absorption process, other reactions take place in parallel, mainly the hydration of CO2 (Eq. (6)) and the direct reaction with hydroxyl ions (Eq. (7)). CO2 + 2 H 2O → HCO3− + H 3O +
(6)
CO2 + OH − → HCO3−
(7)
The value of the rate constant for reaction (6) is given by Pinsent et al. [23] who concluded that compared to the overall reaction rate, its contribution can be negligible without loss of accuracy in aqueous amine solutions. However, the influence of reaction (7) is much more complicated; Astarita et al. [24] gave an estimate of the concentration of the hydroxyl ion for low CO2 loading as:
5
[OH - ] =
KW [ Am] K Am
(8)
and are the dissociation constants of water and amine, respectively. It
was found that hydroxyl ions are consumed at the initial stage, and due to the fact that the amount of hydroxyl ion is small, the contribution of reaction (7) to the overall reaction kinetic is less than 2% [13],[25].Thus based on the termolecular mechanism, the overall reaction rate constant can be given by Eq. (9).
k0 = ∑ kB ,M [ BM ][MEA] + ∑ kB , D [ BD ][ DEA]
(9)
3. Materials and methods 3.1. Materials
Reagent grade MEA and DEA with mass purity of 99% were purchased from Aladdin Industrial Inc. and were used without further purification. Aqueous solutions of DEA (which was used to validate the stopped-flow apparatus) and blends of DEAMEA were prepared to the desired concentrations by adding a weighed quantities amines into deionized water. Commercial-grade CO2 (with a purity of ≥99%) was supplied by Changsha Rizhen Gas Co., China. For each experiment, fresh solution saturated with CO2 was obtained by bubbling the CO2 through deionized water in a jacketed glass stirred reactor. In order to obtain the condition of the pseudo-first order reaction, the saturated CO2 solution was diluted with water to ensure that the molar ratio of amine concentration to CO2 was greater than 10. The concentration of CO2 in the liquid phase was measured with the chromatograph. [15] . 3.2. Stopped-flow technique
The kinetic data for amine-CO2-H2O system in this work were generated from a standard stopped-flow apparatus (Hi-Tech Scientific, UK, Model SF-61DX). It consists of four main parts: a sample handling unit, a conductivity cell, A/D converter, and a microprocessor. Figure 1 shows a schematic drawing of the stopped-flow apparatus. The main section of the sample handling unit not only facilitated the loading and delivery of reactants, but also provided support and enclosure of the sample flow circuit, including the drive syringes, drive valves and observation cell together with interconnecting plumbing. All sample flow circuit components were connected to an external circulator that was thermostatted, and the temperature was 6
controlled within ±0.1 K. For each run, equal volumes of amine solution and CO2 solution were pushed into the observation cell, where the two fresh solutions mixed and started to react. The formation of ion species causes a conductivity change in the cell. The conductivity change as a function of time was measured by a special circuit that gave an output voltage directly proportional to the solution conductivity, as described by Knipe et al. [26]. One sample result (conductance vs. time) from the stopped-flow technique is given in Figure 2. The curve was then fitted to Eq. (10) to estimate the overall reaction rate constant. In Eq. (10), G0 and G is the value of the conductance at the start and the end of the observed reaction respectively, and t (s) is the time. Variation of conductance is then related to CO2 consumption and to k0. G(t) = (G − G )exp (− × t) + G
(10) “Kinetasyst” software installed on the microcomputer can automatically generate the observed pseudo-first-order reaction rate constant ( ). Experiments were repeated at least six times at each temperature for all concentrations. Given that solutions used in the stopped-flow apparatus are usually at relatively low concentrations, they are close to infinite dilution and can be treated as ideal solutions. Under such low concentration and low ionic strength conditions, the effect of ionic strength can be considered to be negligible without resulting in a loss of accuracy [16]. In stopped-flow apparatus, a liquid with dissolved CO2 and another liquid with dissolved amine will mix simultaneously and instantly in a batch reactor (which is the conductivity cell) where the homogeneous reaction of CO2 and amine takes place. Therefore it is not necessary to study the residence time distribution for the fluid in the stopped-flow apparatus due to the fact that the residence times of all the fluid elements are the same in a batch reactor. In addition, the reactor cell is enclosed and filled with liquid, so there is no gas phase involved in the reactor when using stoppedflow technique as in this work. All these factors help to make the investigation of the intrinsic kinetics straightforward.
Figure 1 Schematic drawing of the experimental stopped-flow equipment.
Figure 2 A typical result of the stopped-flow investigation in six repeated 7
experiments. (T=303 K, [DEA]=7.5 mol/m3 and [MEA]=5 mol/m3 ).
4. Results and discussion
4.1. CO2-Aqueous DEA
The kinetics of aqueous DEA solutions (over a concentration range of 140-500 mol/m3) reacting with CO2 were measured at temperatures in the range of 293-313 K, and the values were then compared to corresponding published data from the literature. As Figure 3 shows, results obtained in this work were in good agreement with those presented by Siemieniec et al. [27] and Ali [28], using the same technique for aqueous DEA system. Furthermore, Siemieniec et al. [27] and Ali [28] also verified that values obtained for DEA were in accordance with those of other workers using the stopped-flow technique, as well as rapid mixing, continuously stirred cell reactor and wetted wall methods. Therefore, results obtained in this work were deemed reliable and valid. The expression of can be given by Eq. (11), which is derived from Eq. (2). A nonlinear fitting method was used for estimation of parameters at each temperature. A subroutine called NLINFIT in MatLab was applied for finding the optimum value for each parameter (similarly hereinafter). The obtained values of and are listed in Table 1, representing DEA and H2O as Base in the deprotonation process, respectively. The correlated were found to be in good accordance with the experimental , with an average absolute deviation (AAD%) of 2.32% indicating the termolecular mechanism can explain the CO2-aqueous DEA system well. Parity plot of predicted and experimental in CO2-aqueous DEA system is shown in Figure 4. k0 = (k D [ DEA] + kW [ H 2O])[ DEA]
(11)
Table 1 Fitted value of and based on the termolecular mechanism. Figure 3 Comparison of the values for DEA obtained in this work and from literature.
8
Figure 4 Parity plot of predicted and experimental in CO2-aqueous DEA system.
4.2. CO2-Aqueous MEA+DEA
The kinetics of CO2 and mixed alkanolamines, MEA and DEA, were studied using the stopped-flow technique at 293-313 K, with intervals of 5 K. The concentration of MEA was in the range of 5-15 mol/m3 and the concentration of DEA was in the range of 5-200 mol/m3, concentrations which were chosen in order to not only satisfy the pseudo-first regime, but also to cover various concentration ratios of DEA to MEA. As Figure 5and Figure 6 show, the experimentally obtained were plotted versus concentrations of MEA and DEA for all temperatures (the original data and standard deviation are given in Supportive Information). For all runs, the CO2 reaction rate in aqueous blend of MEA and DEA increased disproportionally as the concentration of one of the amines increased. Moreover, temperature accelerated the overall rate constant in a nonlinear form. The obtained for different temperatures and concentrations were fitted to the rate equation using Eq. (9), as Bosch et al. [29] suggested. However, no meaningful values were found. This result implies a different reaction mechanism in the MEADEA-CO2-H2O system as compared to that reported in the literature on blended amine-CO2 systems, in which all constituent alkanolamines reacting with CO2 were taken into account. According to Li et al. [30], there is a certain critical point in the reaction process below which the DEA carbamate cannot be detected. As well, the critical loading can decrease as the molar ratio of DEA to MEA increases. As shown in Figure 7, when the molar ratio of MEA and DEA is 1:1, the critical loading is 0.17 (mol CO2/mol amine).Thus, the kinetic data were categorized into two groups, low molar ratio and high molar ratio, as shown in Figure 5 and Figure 6, to understand the possible mechanism of reactions between CO2 and blended amines. For the low molar ratio, named as Group A, the concentration of MEA and DEA was 5-15 mol/m3. For the high molar ratio, named as Group B, the concentration of DEA was 140-240 mol/m3 while MEA remained at 10-20 mol/m3. Thus, the molar ratio of DEA to MEA in Group A was around 1:1 while that in Group B was greater than 10:1.
9
Figure 5 of CO2 absorption into the MEA-DEA-H2O system of low concentration of DEA (Group A).
Figure 6 of CO2 absorption into the MEA-DEA-H2O system of high concentration of DEA (Group B).
Figure 7 Ion speciation plots of MEA-DEA-CO2-H2O system at 298 K [30]. 4.2.1. Group A All experimental in Group A were fitted to Eq. (9). However, the experimental could not be fitted successfully to Eq. (9) whether or not the MEA for deprotonation term was included. The parameters in Eq. (9) were found to be meaningless at any investigated temperature. Therefore, a new model is proposed in this work assuming that when the CO2 loading was below 0.1 and the molar ratio of DEA to MEA was around 1:1, only MEA carbamate was formed and DEA simply participated in the deprotonation process, as presented by Eq. (12):
k 0A = (kMA , M [ MEA] + k DA, M [ DEA] + kWA ,M [ H 2O])[ MEA]
(12)
where , , , and , represent the reaction rate constant for one molecule of
MEA, DEA and water depronating one molecule of MEA in Group A respectively. Furthermore, during the fitting of the experimental data, it was observed that MEA had a very small effect on the deprotonation process, and thus, was ignored for model simplification. Then Eq. (12) became:
k 0A = (kDA,M [ DEA] + kWA ,M [ H 2O])[ MEA]
(13)
The , and , in Eq. (13) were obtained and plotted in Figure 8. They can
be expressed as a function of temperature as shown in Eq. (14) and Eq. (15), respectively. As can be seen in Figure 9, the predicted values of calculated from Eq. (13) were observed to be in good agreement with the experimental results, with an AAD of 4.71%. 5650 k DA, M = (1.72 ×107 ) exp − T
(14)
4214 kWA ,M = 59.2exp − T
(15)
Based on the results of the two models, it can be stated that the assumption for 10
Eq. (13) was reasonable. It can be attributed to the nature of the amine. MEA is a primary amine, and DEA is a secondary amine, which can be affected by the detrimental effect of the steric factor. Another reason can be the stronger basicity of MEA than DEA (the values of pKa for MEA and DEA are 9.40 and 8.88 respectively [31]), leading to a stronger possibility of the occurrence of the reaction between MEA and CO2 as compared to the reaction between DEA and CO2. Therefore, at the molar ratio of DEA to MEA around 1:1 (i.e. Group A), the relationship of MEA and DEA appeared generally to be competitive in terms of their direct reactions with CO2 even though DEA also performs as base in MEA deprotonation. Figure 8 Fitted value of , and , in Eq. (13).
Figure 9 Parity plot of predicted and experimental in Group A. 4.2.2. Group B The experimental in group B were fitted to Eq. (13) at first so as to verify if the situation in this group was the same as that in group A. The fitting values obtained were unreasonable because they were either negative or infinite whether or not the term involving MEA for deprotonation was included in the fitting. This result showed that this mechanism was not suitable for conditions in group B. Also, Eq. (9) was used to fit the experimental . It is interesting to find that the inclusion of MEA as base for deprotonation and H2O as base for deprotonation of DEA in Eq. (9) did not result in any meaningful results. Hence, it was assumed that for the case of MEA, its concentration was too small, leading to miniscule probability of collision of MEA with amine-CO2 complex. This is similar to group A, since the same assumption was made. The possible reason of the case of DEA is that the deprotonation reaction rate of DEA by H2O is far slower given the short reaction time used in the stopped-flow technique, in which a slow reaction is not expected to have an effect on the overall reaction rate [32, 33]. Consequently, Eq. (9) was modified to obtain the case where MEA for deprotonation and H2O for deprotonation of DEA were left out, as shown in Eq.(16).
k 0B = (kDB,M [ DEA] + kWB ,M [ H 2O])[MEA] + kDB, D [ DEA]2
(16)
The value of , , , and , in Eq. (16) were obtained and were plotted as
a function of the temperature (Figure 10), which can be expressed using Eq.(17)-(19). 11
As can be seen in the Figure 11, the predicted values of calculated from Eq. (16) were found to be in a good agreement with experimental results, with an AAD of 3.33%. In group B, MEA and DEA both react with CO2 on CO2 absorption in blended amine as described by other works [25, 34, 35]. 5276 k DB,M = (5.82 × 105 ) exp − T
(17)
3590 kWB ,M = 6.79exp − T
(18)
3165 k DB, D = 43.268exp − T
(19)
Figure 10 Fitted value of , , , and , in Eq. (16).
Figure 11 Parity plot of predicted and experimental in Group B. The fitted model (Eq. (13) and Eq. (16)) shows that the reaction mechanism was different between Groups A and B, and implies that it was MEA which reacted with CO2 first, or at the initial region the DEA carbamate formation was insignificant, and so, could be ignored. When the amount of CO2 or the molar ratio of DEA to MEA reaches a certain level, DEA starts to react with CO2 and a more complicated situation in terms of the amine mixture of the CO2 absorption process exists, giving rise to the occurrence of both reactions in solution. The condition of Group A was below the critical point, where the competition of MEA and DEA was obvious given MEA had an overwhelming inhibitory effect to DEA; nevertheless the inferior position of DEA was made up by the high concentration of DEA, i.e. Group B, where both MEA and DEA generally played significant roles. It should be noted that the competition and cooperation of MEA and DEA were present in both regions. This mechanism compares favorably with the NMR results and explains experimental kinetic results in this work successfully as discussed above.
4.3. Comparison of Reaction Kinetics of Aqueous MEA+DEA and Pure Aqueous DEA Due to DEA having little influence in the reactions of Group A, the comparison 12
of aqueous MEA+DEA and pure aqueous DEA here was restricted to the comparison of the addition of MEA to DEA (i.e. Group B) and its blank system (pure DEA). As the experimental data showed, the addition of small amounts of MEA to DEA resulted in a significant enhancement of the CO2 reaction rate. For instance, the overall reaction rate constant of aqueous system increased by 131% with addition of 0.01 M MEA in 0.16 M DEA, at 298 K ( at 0.16 M DEA: 29.31 s-1, and at 0.16 M DEA + 0.01 M MEA: 67.34 s-1). It should be noted that , compares favorably to , as Figure 10 shows, since
these two parameters represent the same reaction step in different environments, the former being kinetic constant for reaction of two molecules of DEA and one molecule of CO2 in blends of high ratio of DEA to MEA while the latter being that in pure DEA solution. 5. Conclusions
This paper has presented the CO2 absorption kinetics of mixed alkanolamines, MEA and DEA, using the stopped-flow technique at the temperature range of 293-313 K. The kinetic data were categorized into two groups, Group A and Group B. In Group A, the concentrations of both MEA and DEA were in the range of 5-15 mol/m3. In Group B, the concentration of DEA was in the range of 140-240 mol/m3, while the concentration of MEA remained the same at 10-20 mol/m3. The possible reaction mechanism between CO2 and blended amines was proposed. Two models based on termolecular mechanism were adopted to fit the experimental in both groups in order to explain the kinetic data. The results showed that the predicted exhibited good agreement with the experimental data with an AAD of 4.71% for Group A and an AAD of 3.33% for Group B. Comparison of each optimal model for Group A and Group B suggests that below a certain critical point (i.e. Group A), MEA had an overwhelming inhibitory effect to DEA and the contribution to was mainly from the reaction of MEA and CO2, since the DEA carbamate formation was nonexistent or insignificant; while above this point (i.e. Group B), the overall reaction of CO2 with aqueous blends of MEA and DEA could be regarded as a reaction between CO2 and MEA in parallel with the reaction of CO2 with DEA.
Acknowledgements 13
The financial support from the National Natural Science Foundation of China (NSFCNos. 21476064, U1362112 and 21376067), National Key Technology R&D Program (MOST-Nos.2012BAC26B01 and 2014BAC18B04), Innovative Research Team Development Plan (MOE-No.IRT1238), Specialized Research Fund for the Doctoral Program of Higher Education (MOE-No. 20130161110025), China’s State “Project 985” in Hunan University-Novel Technology Research & Development for CO2 Capture, Key Project of International & Regional Cooperation of Hunan Provincial Science and Technology plan (2014WK2037), and China Outstanding Engineer Training Plan for Students of Chemical Engineering & Technology in Hunan University (MOE-No.2011-40) is gratefully acknowledged. We also greatly appreciate Mr. Wilfred Olson for his great contribution to help us correct any grammar mistakes and insightful inputs to our research work.
14
Nomenclature
AAD
absolute average deviation
Am
amine
B
base (amine, water or hydroxyl ion)
DEA
diethanolamine
observed pseudo first order reaction rate constant (s-1)
, , , , ,
second order kinetic rate constant between amine and CO2 with amine as base (m6/mol2 s2) second order kinetic rate constant between base, amine and CO2 (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with DEA as base (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with DEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with MEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between MEA and CO2 with DEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between MEA and CO2 with MEA as base in blended amine (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with water as base (m6/mol2 s2) second order kinetic rate constant between DEA and CO2 with water as base in blended amine (m6/mol2 s2)
dissociation constant of amine (mol/m3)
dissociation constant of water (mol/m3)
M
molarity (mol/m3)
MEA
monoethanolamine
T
temperature (K)
[]
concentration (mol/m3)
15
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18
Figure Captions Figure 1 Schematic drawing of the experimental stopped-flow equipment.
Figure 2 A typical result of the stopped-flow investigation in six repeated experiments. (T=303 K, [DEA]=7.5 mol/m3 and [MEA]=5 mol/m3 ) Figure 3 Comparison of the values for DEA obtained in this work and from literature. Figure 4 Parity plot of predicted and experimental in CO2-aqueous DEA system. Figure 5 of CO2 absorption into the MEA-DEA-H2O system of low concentration of DEA (Group A). Figure 6 of CO2 absorption into the MEA-DEA-H2O system of high concentration of DEA (Group B). Figure 7 Ion speciation plots of MEA-DEA-CO2-H2O system at 298 K [30]. Figure 8 Fitted value of , and , in Eq. (13).
Figure 9 Parity plot of predicted and experimental in Group A. Figure 10 Fitted value of , , , and , in Eq. (16).
Figure 11 Parity plot of predicted and experimental in Group B.
19
Figure 1 Schematic drawing of the experimental stopped-flow equipment.
Figure 2 A typical result of the stopped-flow investigation in six repeated experiments. (T=303 K, [DEA]=7.5 mol/m3 and [MEA]=5 mol/m3 )
Figure 3 Comparison of the values for DEA obtained in this work and from literature. 20
Figure 4 Parity plot of predicted and experimental in CO2-aqueous DEA system.
Figure 5 of CO2 absorption into the MEA-DEA-H2O system of low concentration of DEA (Group A).
Figure 6 of CO2 absorption into the MEA-DEA-H2O system of high concentration of DEA (Group B). 21
Figure 7 Ion speciation plots of MEA-DEA-CO2-H2O system at 298 K [30].
Figure 8 Fitted value of , and , in Eq. (13).
Figure 9 Parity plot of predicted and experimental in Group A.
22
Figure 10 Fitted value of , , , and , in Eq. (16).
Figure 11 Parity plot of predicted and experimental in Group B.
23
Table Captions
Table 1 Fitted value of and based on the termolecular mechanism.
24
Table 1 Fitted value of and based on the termolecular mechanism. kD
kW
(104 m6/mol2 s2)
(106 m6/mol2 s2)
293
5.44±0.15
1.95±0.11
298
6.31±0.27
2.70±0.20
303
7.91±0.25
3.20±0.19
308
9.52±0.44
3.94±0.33
313
12.16±0.37
4.69±0.27
T (K)
25
Highlights: •
Reaction kinetics between CO2 and blends of MEA and DEA are studied.
•
Two possible situations are discussed based on the ratio of blends.
•
DEA barely reacts with CO2 in blends at low concentration.
•
The proposed mechanisms explain all experimental data very well with excellent AADs less than 5%.
26