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Ind. Eng. Chem. Res. 2007, 46, 4426-4434
Reaction Kinetics of CO2 in Aqueous Ethylenediamine, Ethyl Ethanolamine, and Diethyl Monoethanolamine Solutions in the Temperature Range of 298-313 K, Using the Stopped-Flow Technique Juelin Li, Amr Henni,* and Paitoon Tontiwachwuthikul International Test Center for CO2 Capture, Faculty of Engineering UniVersity of Regina, Saskatchewan, Canada S4S 0A2
The observed pseudo-first-order rate constants (k0) for the reactions between CO2 and ethylenediamine (EDA), ethyl ethanolamine (EEA), and diethyl monoethanolamine (DEMEA) have been studied using the stoppedflow technique in an aqueous solution at 298, 303, 308, and 313 K. The amine concentrations ranged from 26.2 mol/m3 to 67.6 mol/m3 for EDA, 28.2 mol/m3 to 81.9 mol/m3 for EEA, and 196.5 mol/m3 to 997.4 mol/m3 for DEMEA. The zwitterion mechanism was used to correlate the experimentally obtained rate constants. Both the zwitterion formation step and the proton removal step had a significant role for the primary and secondary amines (EDA and EEA). The reaction rate of CO2 in an aqueous EDA solution was observed to be much faster than that in aqueous MEA solution. The rate in aqueous EEA was much faster than in aqueous DEA, under the conditions studied. Finally, the reaction rate constant of CO2 in an aqueous tertiary amine (DEMEA) solution was observed to be much faster than that in methyl diethanolamine (MDEA). Only the zwitterion formation step had a significant role in the overall reaction. The base catalysis of the CO2 hydration mechanism could explain the reaction between CO2 and the tertiary amine. Therefore, the three selected amines are considered to be of interest to the gas sweetening industry. 1. Introduction As a way to mitigate rising greenhouse gas emissions, governments, industries, and researchers are becoming increasingly interested in CO2 capture from gas streams. Energy demand is projected to increase by 60% by the year 2030; CO2 emissions will increase by 63% from today’s level, 90% higher than 1990 levels. Gas sweetening by alkanolamines has been used on a commercial scale for more than 75 years. Monoethanolamine (MEA) as a primary amine, diethanolamine (DEA) as a secondary amine, and methyl diethanolamine (MDEA) as a tertiary amine are the most widely used alkanolamines. Research is ongoing to find new alkanolamines (or amines) with a higher capacity for CO2 absorption that react faster and require less energy to regenerate than what is commercially used today. In earlier publications, Jensen and Christensen,1 Weiland and Trass,2 and Hikita et al.3 published limited reaction kinetics of CO2 in aqueous ethylenediamine (EDA), using competitive reaction, liquid jet, and rapid mixing methodologies, respectively. They concluded that this primary diamine reacted rapidly with CO2. However, the techniques used for the kinetic measurements were all indirect methods, and, therefore, it appeared prudent to undertake an additional study of this amine using the stopped-flow apparatus, which is a direct technique. For diamines such as EDA, at equilibrium and high pressures of CO2, the two amines interact with CO2 molecules and, therefore, should have a higher absorption capacity with CO2 than ethanolamines. CO2 loadings of R ≈ 3 were reported by Alem in 20054 for a diamine (5 wt % aqueous piperazine) at 6000 kPa and 40 °C. * To whom correspondence should be addressed. Tel.: 306 585 4960. Fax: 306 585 4855. E-mail address: amr.henni@ uregina.ca.
Ethyl ethanolamine (EEA) is a linear secondary alkanolamine. It is a well-established fact that the basicity is provided by the amino site, whereas the presence of the -OH group results in increased solubility in water.4 However, in EEA molecule, an ethyl group replaces a H atom of an amino group in MEA. The ethyl group is supposed to increase the basicity of the amine, but it might also increase the hindrance around the N atom. Only Bavbek and Alper5 studied the kinetics of CO2 in aqueous EEA solutions using the stopped-flow method at one temperature (303 K) and in a narrower range of concentration. Tertiary amines react with CO2 at a finite and slow rate, but react instantaneously with H2S involving only a proton-transfer mechanism; therefore, such compounds are commonly used in selective gas sweetening.6 Tertiary alkanolamines require a very small stripping energy, when compared to other conventional amines. In a comprehensive literature review, Versteeg et al.7 concluded that the kinetic results of diethyl monoethanolamine (DEMEA) have large scattered results among different workers and, therefore, were considered to be not wellestablished. It was then thought interesting to study the reaction kinetics of a primary diamine (EDA), a secondary alkanolamine (EEA), and a tertiary alkanolamine (DEMEA), using the stopped-flow technique, which is a direct method. The chemical structures of the amines studied in this work and MEA, DEA, and MDEA are shown in Figure 1. 2. Reaction Fundamentals The zwitterion mechanism is well-established for the description of the reaction between CO2 and primary or secondary amines. It was originally proposed by Caplow,8 and then reintroduced by Danckwerts.9 The steps are shown below.
10.1021/ie0614982 CCC: $37.00 © 2007 American Chemical Society Published on Web 05/18/2007
Ind. Eng. Chem. Res., Vol. 46, No. 13, 2007 4427
Under pseudo-first-order conditions, the concentration of the amine ([Am]) was always in great excess of that of CO2, so eq 6 could be written as
rCO2 ) k0[CO2]
(8)
The observed pseudo-first-order reaction rate constant ko then becomes
k0 )
[Am] (1/k2) + (1/
∑ kB[B])
(9)
Two limiting cases exist for eq 9. If the inverse of k2 is much larger than the inverse of ΣkB[B], which means that the zwitterion is deprotonated relatively fast, eq 9 simplifies to
k0 ) k2[Am]
(10)
Figure 1. Chemical structures of the amine solvents.
However, if the inverse of k2 is much smaller than 1/ΣkB[B]), the value of k0 then would be
Zwitterion formation: k2(k-1)
CO2 + RNH2 798 RN+H2CO2-
(1)
k0 ) [Am](ΣkB[B])
(2)
As can be shown by eq 11, at higher concentrations where the relative value of the different terms can change, a shift in the order of the reaction is possible. Very short reaction times are considered here; therefore, the results of the experiments involve most probably, on average, primarily the fastest reacting of the two amines in a diamine (i.e., the primary amine).
Deprotonation by a base (B): kBase(k-Base)
RN+H2CO2- + B 798RNHCO2- + BH+
The suffix B represents any base, which can be an amine, a hydroxyl ion and water, or any other solvent present other than water (such as methanol or ethanol). In aqueous solution, the deprotonation proceeded mainly via water and the amine, according to Versteeg and van Swaajj,6 because the OH- ion contribution generally was negligible in the zwitterion dissociation. Under the steady-state approximation assumption, the zwitterion concentration can be expressed as
k2[CO2][Am]
[zwitterion] )
k-1 +
∑ kBase[B]
(3)
The overall CO2 reaction rate was equal to
-rCO2 ) k2[CO2][Am] - k-1[zwitterion]
(4)
After substituting the concentration of the active intermediate [zwitterion] into eq 4, the overall forward reaction rate of CO2 can be derived with the assumption of a pseudo-steady state for the zwitterion concentration as
rCO2 )
[CO2][Am] (1/k2) + (k-1/{k2(
∑ kBase[B])})
(5)
Equation 5 can be simplified to
rCO2 )
[CO2][Am] (1/k2) + (1/
∑
(6) kB[B])
where
kB )
k2kBase k-1
(7)
(11)
3. Experimental Apparatus and Procedure The experimental technique considered in this study was a direct method, using the standard model SF-51 stopped-flow equipment manufactured by Hi-Tech Scientific, Ltd. (U.K.). It consisted of four major parts: a sample handling unit, a conductivity-detection cell, A/D converter, and a microprocessor from an Apple II microcomputer. Figure 2 shows a schematic drawing of the experimental stopped-flow equipment. The sample-handling unit was largely constructed of stainless steel, which provided the support and the enclosure of the sample flow circuit. In the front panel of this unit, a mode selector was provided, which included a temperature reading, an air pressure display, and pneumatic push plates. These plates control the stop/waste valve from the software, which allows the stop syringe to be emptied automatically before starting the run. When running the sample measurement, the valve automatically moved to the waste position and emptied the stop syringe. The valve then moved back to the drive position and the pneumatic drive plate pushed the fresh solution to the observation cell, replacing the old solution (from the previous run). A conductivity-detection system was used to measure the intrinsic rate of rapid homogeneous reaction directly. The sample flow circuits were well-thermostated, and the temperature control was within (0.1 K. It is important to mention that the amine and the CO2 solutions were placed in sealed drive syringes and pushed in equal volumes to the observation cell of the apparatus. The conductivity cell monitors the ion formation, as a function of time, as the ion formation initiates a voltage change inside the cell. The conductivity change was measured by a circuit, as described by Knipe et al.,10 which gave an output voltage that was directly proportional to the solution conductivity. The microcomputer automatically generated the observed pseudofirst-order constant k0, based on the output voltage values.
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Ind. Eng. Chem. Res., Vol. 46, No. 13, 2007
Figure 2. Schematic drawing of the experimental stopped-flow equipment.
Experiments were repeated at least seven times at each temperature for all concentrations. The other parts of the equipment were responsible for controlling the stepper motor on each drive of the sample handling unit, and they also provided a central control to power all internal units and the sample handling unit and controlled the auto emptying waste cycle and air drive control circuitry. The stopped-flow technique did not involve gas-phase absorption, as opposed to most of the other kinetic devices. Therefore, the results of the direct method used here corresponded to the intrinsic homogeneous reaction rate of solutions. This direct method was not affected by the reversibility of reaction and avoided the possible experimental errors caused by the depletion of the amine in the gas-liquid interface. In addition, each run only needed a small amount of reactants (∼0.1 mL of each reactant) and the equipment was very easy to handle. Reagent-grade EDA with a mass purity of 99%, EEA with a purity of 98%, and DEMEA (99.5% purity) were obtained from Sigma-Aldrich. All the chemicals were used without further purification. Deionized water was used when needed. For each run, freshly saturated CO2 solution was obtained by mixing the gas through the water in a jacketed glass stirred reactor (used otherwise as a phase behavior cell (Autoclave Engineers)), and then diluted with the water, to ensure an appropriate ratio (>10) of amine molar concentration to CO2 molar concentration for each experiment. In this work, a set of runs was performed to measure the kinetics of DEA with CO2 at different temperatures, and the results were compared with corresponding published data. Based on the comparison of the results for aqueous MEA and DEA systems obtained from the literature using a similar method, we estimated the data obtained has an uncertainty of 5%. Figure 3 shows that the values obtained in this work compared favorably with those obtained by Ali,11 using the same technique for the aqueous DEA system. Furthermore, a comparison of the k0 values of the reaction between MEA and CO2 at 298 K obtained in this study is shown in Figure 4. The figure shows good agreement with the values published by Ali.11 Ali further verified that the values he obtained for DEA and MEA were in accordance with those of other workers, using the direct method, as well as other indirect techniques (such as the stirred cell, wetted wall, or rapid mixing). Therefore, the experimental technique followed by this work was deemed valid. 4. Results and Discussion 4.1. EDA in Aqueous Solution. Aqueous EDA concentrations ranging from 26.2 mol/m3 to 67.6 mol/m3 were studied
Figure 3. Comparison of the k0 values for DEA obtained in this work and from Ali11 at different temperatures.
Figure 4. Comparison of the k0 values for MEA obtained in this work and from Ali11 at 298 K.
at temperatures of 298, 303, 308, and 313 K. The observed pseudo-first-order rate constants (k0) for EDA in aqueous solution are shown in Figure 5, as a function of the EDA concentration. The rate constants exhibited the expected trend of an increase in the reaction rate with an increase in EDA concentration and in temperature. By fitting the empirical power-law kinetics to the data of experimentally observed pseudo-first-order constants for CO2 in Figure 5, the reaction orders were determined to be 1.01, 1.07, 1.08, and 1.18, with respect to [EDA] for 298, 303, 308, and 313 K, respectively.
Ind. Eng. Chem. Res., Vol. 46, No. 13, 2007 4429
Figure 5. Effect of the EDA concentration on the k0 values at different temperatures in aqueous solution. Table 1. Literature Data on the Reaction between CO2 and Aqueous EDA at 303 K investigators
k2 (m3 mol-1 s-1)
method
Jensen and Christensen1 Weiland and Trass2 Hikita et al.3 this work
2.23 × 101 13.8 × 101 1.79 × 101 1.77 × 101
competitive reaction liquid jet rapid mixing stopped-flow
All values were practically very close to 1, in agreement with the work published by Hikita et al.3 The second-order rate constants (k2) of the reaction have been reported by several researchers, as shown in Table 1. The k2 values of this work at 303 K were determined to compare favorably with the k2 values reported by Hikita et al.3 Katchalski et al.12 have reported that a side reaction,
2CO2 + 2H2N(CH2)2NH2 f +H3N(CH2)2NH3+ + -
OOCNH(CH2)2NHCOO- (12)
yielded an intramolecular salt. However, Jensen and Christensen1 did not consider this side reaction and only accounted for the monocarbamate in the overall reaction rate. Therefore, their results were slightly higher than the data obtained in this work. The experimental values of the observed pseudo-first-order constants, as a function of temperature and concentration, are presented in Table 2. The reaction orders increased as the temperature increased. Equation 9 was fitted to the experimental data for k0, using a nonlinear minimization technique in the Sigmaplot software. The generated rate constants k2, ka, and kw of the different reactions leading to carbamate ion formation are shown in Table 3. Figure 6 is the parity plot of the experimentally obtained results for k0 against the correlated values. This figure clearly indicates that the zwitterion intermediate mechanism describes the experimental data very satisfactorily. The mean relative deviation (MRD) was determined to be 2% for all the data of the given system and was calculated as follows:
MRD ) Experimental(k0) - Correlated(k0) 1 abs × 100 (13) n n Experimental(k0)
∑
(
)
where n is the number of experimental data. The increase in the temperature caused the expected effect of increasing the k2 values. In addition, ka has the same systematic trend as k2, as can be observed in Table 3 (but not
Figure 6. Parity plot for the reaction rate constants of CO2 in aqueous EDA solution.
Table 2. Pseudo-First-Order Rate Constants for CO2 Reaction in Aqueous EDA [EDA] (mol/m3)
temperature (K)
k0 (s-1)
26.21 26.21 26.21 26.21
298 303 308 313
258 ( 3 328 ( 11 410 ( 18 513 ( 16
31.00 31.00 31.00 31.00
298 303 308 313
304 ( 9 396 ( 7 477 ( 15 561 ( 9
33.28 33.28 33.28 33.28
298 303 308 313
334 ( 10 419 ( 10 520 ( 49 664 ( 15
41.68 41.68 41.68 41.68
298 303 308 313
402 ( 11 525 ( 8 648 ( 26 863 ( 43
49.92 49.92 49.92 49.92
298 303 308 313
516 ( 9 648 ( 19 792 ( 40 1107 ( 133
59.62 59.62 59.62 59.62
298 303 308 313
591 ( 24 812 ( 73 958 ( 130 1244 ( 114
67.56 67.56 67.56 67.56
298 303 308 313
663 ( 27 887 ( 53 1150 ( 74 1535 ( 177
the kw values). The second-order reaction rate constant (k2) was fitted to an Arrhenius expression, as shown below:
k2 ) A exp
( ) -Eact RT
(14)
where A is the Arrhenius constant (given in units of m3 mol-1 s-1), Ea the activation energy (given in units of kJ/mol), and R the universal gas constant (0.008315 kJ mol-1 K-1). Figure 7 displays the corresponding Arrhenius plot from which the activation energy (Eact) for the zwitterion formation step was determined to be 50.56 kJ/mol. The value is close to that obtained by Hikita et al.3 (53.6 kJ/mol) for EDA in aqueous
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Table 3. Reaction Rate Constants in EDA Aqueous Solution temperature (K) 298 303 308 313
reaction order, n
k2 (m3 mol-1 s-1)
1.01 1.07 1.08 1.18
1.20 × 1.77 × 101 2.22 × 101 3.30 × 101
(-6080 T )
kw (m6 mol-2 s-1) 9.83 × 10-4 6.32 × 10-4 7.35 × 10-4 5.05 × 10-4
6.14 × 2.83 × 10-1 3.44 × 10-1 6.23 × 10-1
Table 4. Contribution of the Different Kinetic Constants on the Overall Observed Rate Constant for CO2 Reaction in Aqueous EDA
solution. The Arrhenius expression is
k2 ) 8.7 × 109 exp
ka (m6 mol-2 s-1) 10-2
101
Contribution
(15)
To compare the relative significance of each different reaction rate constant (k2, ka, and kw) on the overall observed pseudofirst-order rate constant k0, the values of terms 1/k2, 1/(ka[EDA] + kw[H2O]), ka[EDA], and kw[H2O] have been calculated separately and presented in Table 3. For this system, it was clear that neither the first term (1/k2) nor the second term 1/(ka[EDA] + kw[H2O]) were negligible by comparison. This meant that both the zwitterion formation step and the zwitterion deprotonation step affected the reaction rate. Moreover, the ka[EDA] terms were determined to compare favorably with the corresponding kw[H2O] terms, so both EDA and water had a dominating role and competed with each other in the zwitterion deprotonation step, under the conditions of the concentrations and temperatures studied, as shown in Table 4. The high dilution (high water concentration) of the amine solution allowed the term kw[H2O] to be larger that ka[EDA], although ka was orders of magnitude larger than the value of kw (see Table 3). For a given temperature, the term kw[H2O] was virtually constant. The reaction kinetics were best expressed by eq 9. Ethylenediamine (EDA) reacted faster with CO2 than did MEA solutions at the same concentration (see Figures 4 and 5). At equilibrium, the capacity of EDA for CO2 is expected to also be higher, because of the presence of the two amino groups. Unfortunately, the boiling point of EDA was close to 118 °C, which is much lower than that of MEA (170 °C). 4.2. EEA in Aqueous Solution. Aqueous EEA concentrations ranging from 28.2 mol/m3 to 81.9 mol/m3 were studied at temperatures of 298, 303, 308, and 313 K. A plot of the observed pseudo-first-order rate constants (k0) versus [EEA] (Figure 8) gave the reaction orders of 2.08, 2.06, 1.91, and 1.93, with respect to [EEA] for 298, 303, 308, and 313 K, respectively. Bavbek and Alper4 estimated the reaction order of EEA at 303 K to be equal to 2.0, which is consistent with our measurements.
Figure 7. Arrhenius plots for the reaction of CO2 in aqueous EDA, EEA, and DEMEA solutions.
[EDA] (mol/m3)
298 K
303 K
308 K
313 K
8.34 × 10-2
1/k2 Term 5.64 × 10-2
4.51 × 10-2
3.03 × 10-2
26.21 31.00 33.28 41.68 49.92 59.62 67.56
1/(ka[EDA] + kw[H2O]) Term 1.78 × 10-2 2.36 × 10-2 2.01 × 10-2 1.78 × 10-2 2.29 × 10-2 1.95 × 10-2 1.77 × 10-2 2.25 × 10-2 1.92 × 10-2 1.76 × 10-2 2.14 × 10-2 1.82 × 10-2 1.75 × 10-2 2.04 × 10-2 1.73 × 10-2 1.73 × 10-2 1.93 × 10-2 1.64 × 10-2 1.72 × 10-2 1.86 × 10-2 1.57 × 10-2
2.26 × 10-2 2.12 × 10-2 2.06 × 10-2 1.86 × 10-2 1.70 × 10-2 1.54 × 10-2 1.43 × 10-2
26.21 31.00 33.28 41.68 49.92 59.62 67.56
1.61 × 100 1.90 × 100 2.04 × 100 2.56 × 100 3.06 × 100 3.66 × 100 4.15 × 100
ka[EDA] Term 7.41 × 100 8.77 × 100 9.41 × 100 1.18 × 101 1.41 × 101 1.69 × 101 1.91 × 101
9.00 × 100 1.07 × 101 1.14 × 101 1.43 × 101 1.71 × 101 2.05 × 101 2.32 × 101
1.63 × 101 1.93 × 101 2.07 × 101 2.60 × 101 3.11 × 101 3.71 × 101 4.21 × 101
26.21 31.00 33.28 41.68 49.92 59.62 67.56
5.44 × 101 5.44 × 101 5.43 × 101 5.43 × 101 5.42 × 101 5.41 × 101 5.41 × 101
kw[EDA] Term 3.50 × 101 4.07 × 101 3.50 × 101 4.06 × 101 3.50 × 101 4.06 × 101 3.49 × 101 4.06 × 101 3.49 × 101 4.05 × 101 3.48 × 101 4.05 × 101 3.48 × 101 4.04 × 101
2.80 × 101 2.79 × 101 2.79 × 101 2.79 × 101 2.79 × 101 2.78 × 101 2.78 × 101
Figure 9 exhibited the expected trend of increasing the reaction rate with increases in both the concentration and temperature, as evidenced by the k0 values (see Table 5). The experimental data for k0 were fitted to eq 9. The rate constants k2, ka, and kw are shown in Table 6. Figure 9 is a plot of the experimentally obtained results k0 against the correlated values. These two sets of values were determined to have a mean relative deviation (MRD) of 8%, which indicated an acceptable overall parity. The values of k2, ka and kw all increased with increasing temperature (see Table 6).
Figure 8. Effect of EEA concentration on the k0 values at different temperatures in aqueous solution.
Ind. Eng. Chem. Res., Vol. 46, No. 13, 2007 4431 Table 7. Contribution of the Different Kinetic Constants to the Overall Observed Rate Constant for CO2 Reaction in Aqueous EEA Solution Contribution [EEA] (mol/m3)
Figure 9. Parity plot for the reaction rate constants of CO2 in aqueous EEA solution. Table 5. Pseudo-First-Order Rate Constants for CO2 Reaction in Aqueous EEA [EEA] (mol/m3)
temperature (K)
k0 (s-1)
28.21 28.21 28.21 28.21
298 303 308 313
24 ( 1 31 ( 2 46 ( 3 54 ( 1
33.60 33.60 33.60 33.60
298 303 308 313
45 ( 5 49 ( 3 64 ( 2 75 ( 1
45.21 45.21 45.21 45.21
298 303 308 313
87 ( 1 100 ( 4 113 ( 4 132 ( 5
308 K
313 K
1.25 × 10-1
1/k2 Term 1.06 × 10-1
7.73 × 10-2
6.89 × 10-2
28.21 33.60 45.21 57.44 66.20 71.71 81.90
1/(ka[EEA] + kw[H2O]) Term 6.97 × 10-1 5.79 × 10-1 4.80 × 10-1 5.86 × 10-1 4.86 × 10-1 4.03 × 10-1 4.35 × 10-1 3.61 × 10-1 2.99 × 10-1 3.43 × 10-1 2.84 × 10-1 2.36 × 10-1 2.97 × 10-1 2.47 × 10-1 2.05 × 10-1 2.74 × 10-1 2.28 × 10-1 1.89 × 10-1 2.40 × 10-1 1.99 × 10-1 1.65 × 10-1
4.04 × 10-1 3.39 × 10-1 2.52 × 10-1 1.99 × 10-1 1.72 × 10-1 1.59 × 10-1 1.39 × 10-1
28.21 33.60 45.21 57.44 66.20 71.71 81.90
1.43 × 100 1.71 × 100 2.30 × 100 2.92 × 100 3.36 × 100 3.64 × 100 4.16 × 100
ka[EEA] Term 1.73 × 100 2.06 × 100 2.77 × 100 3.52 × 100 4.06 × 100 4.39 × 100 5.02 × 100
2.08 × 100 2.48 × 100 3.34 × 100 4.24 × 100 4.89 × 100 5.30 × 100 6.05 × 100
2.47 × 100 2.95 × 100 3.97 × 100 5.04 × 100 5.81 × 100 6.29 × 100 7.18 × 100
28.21 33.60 45.21 57.44 66.20 71.71 81.90
5.53 × 10-7 5.52 × 10-7 5.50 × 10-7 5.49 × 10-7 5.47 × 10-7 5.48 × 10-7 5.47 × 10-7
kw[H2O] Term 1.11 × 10-6 1.10 × 10-6 1.10 × 10-6 1.10 × 10-6 1.09 × 10-6 1.10 × 10-6 1.09 × 10-6
2.38 × 10-6 2.37 × 10-6 2.37 × 10-6 2.36 × 10-6 2.35 × 10-6 2.36 × 10-6 2.35 × 10-6
2.82 × 10-6 2.82 × 10-6 2.81 × 10-6 2.80 × 10-6 2.79 × 10-6 2.79 × 10-6 2.79 × 10-6
Value
298 303 308 313
128 ( 3 160 ( 5 184 ( 4 218 ( 10
66.20 66.20 66.20 66.20
298 303 308 313
157 ( 8 192 ( 8 232 ( 5 281 ( 8
71.71 71.71 71.71 71.71
298 303 308 313
194 ( 8 215 ( 8 274 ( 9 315 ( 5
81.90 81.90 81.90 81.90
298 303 308 313
251 ( 6 299 ( 14 348 ( 12 432 12
parameter
@ 298 K
@ 303 K
@ 308 K
@ 313 K
k2 (m3 mol-1 s-1) 8.05 × 100 9.35 × 100 1.24 × 101 1.45 × 101 ka (m3 mol-1 s-1) 5.07 × 10-2 6.15 × 10-2 6.90 × 10-2 8.77 × 10-2
Figure 7 shows the corresponding Arrhenius plot (ln k2 versus 1/T) of this system. The regression of the k2 values gave the following Arrhenius equation:
(-3926.2 T )
303 K
Table 8. Rate Constants for CO2 Reaction with Aqueous EEA Obtained by Fitting k0 to eq 17
57.44 57.44 57.44 57.44
k2 ) 4.16 × 106 exp
298 K
The calculated values of 1/k2, 1/(ka[EEA] + kw[H2O]), ka[EEA], and kw[H2O] were presented separately in Table 7 at four different temperatures and at different EEA concentrations. The relative significance of each different reaction rate constant (k2, ka and kw) on the overall observed pseudo-first-order rate constant k0 can be clearly noted. Table 7 shows that the second term, 1/(ka[EEA] + kw[H2O]), was not negligible, in comparison with the first term (1/k2). Therefore, both the zwitterion formation step and the deprotonation step were very important in the overall reaction. It is worth mentioning that the ka[EEA] term was much larger than the corresponding kw[H2O] term (see Table 7). The major contribution to the proton removal step was provided by the amine (EEA). The term kw[H2O] could be ignored, and the reaction kinetics could therefore be expressed as
(16)
From the slope of the plot, the activation energy for the zwitterion formation step of this work was determined to be 32.65 kJ/mol.
k0 )
[RNH2] (1/k2) + (1/ka[RNH2])
Consequently, the values of k2 and ka determined by eq 17 are shown in Table 8, and were determined to be very close to
Table 6. Reaction Rate Constants of EEA in Aqueous Solution temperature (K) 298 303 308 313
(17)
reaction order, n
k2 (m3 mol-1 s-1)
2.08 2.06 1.91 1.93
8.00 × 9.44 × 100 1.29 × 101 1.45 × 101 100
ka (m6 mol-2 s-1) 10-2
5.08 × 6.13 × 10-2 7.39 × 10-2 8.77 × 10-2
kw (m6 mol-2 s-1) 1.00 × 10-11 2.00 × 10-11 4.30 × 10-11 5.10 × 10-11
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Figure 10. Effect of DEMEA concentrations on the k0 values at different temperatures in aqueous solution. Table 9. Pseudo-First-Order Rate Constants for CO2 Reaction in Aqueous DEMEA Solution [DEMEA] (mol/m3)
temperature (K)
k0 (s-1)
196.52 196.52 196.52 196.52
298 303 308 313
17 ( 1 25 ( 1 37 ( 1 52 ( 1
300.80 300.80 300.80 300.80
298 303 308 313
24 ( 1 37 ( 1 54 ( 2 74 ( 3
402.21 402.21 402.21 402.21
298 303 308 313
32 ( 1 48 ( 1 68 ( 1 94 ( 2
502.18 502.18 502.18 502.18
298 303 308 313
40 ( 2 59 ( 1 81 ( 2 109 ( 1
608.78 608.78 608.78 608.78
298 303 308 313
47 ( 1 68 ( 1 96 ( 2 128 ( 2
758.17 758.17 758.17 758.17
298 303 308 313
60 ( 2 81 ( 2 112 ( 2 158 ( 3
997.36 997.36 997.36 997.36
298 303 308 313
67 ( 2 97 ( 4 139 ( 4 192 ( 11
Figure 11. Parity plot for the reaction rate constants of CO2 in aqueous DEMEA solution.
reaction rate, relative to increases in both concentration and temperature, as evidenced by the k0 values (see Table 9). The reaction orders decreased marginally as the temperature increased. The experimental data for k0 were fitted to eq 9. The rate constants k2, ka, and kw were presented in Table 10. Figure 11 is a plot of the experimentally obtained k0 results against the correlated values placing the generated reaction constants k2, ka, and kw in eq 9. These two sets of values have an MRD of 7%, which clearly indicated that these two sets of values correspond well with each other. The corresponding Arrhenius plot for the second-order rate constant (k2) of DEMEA is displayed in Figure 7. From the slope of the plot, the activation energy for the zwitterion formation step of this work was determined to be 51.87 kJ/ mol. The regression of k2 values gave the following Arrhenius equation:
k2 ) 9.95 × 107 exp
those obtained using eq 9, as shown in Table 6. Figures 8 and 3 (extrapolation) show that the reaction rate of EEA with CO2 was faster than that of DEA (secondary amine). 4.3. DEMEA in Aqueous Solution. Aqueous DEMEA concentrations ranging from 196.5 mol/m3 to 997.4 mol/m3 were studied at temperatures of 298, 303, 308, and 313 K. The observed pseudo-first-order rate constants (k0) are plotted in Figure 10, as a function of the DEMEA concentration, and gave, as exponents, reaction orders of 0.89, 0.84, 0.81, and 0.80, with respect to [DEMEA] for 298, 303, 308, and 313 K, respectively. This figure also shows the expected trend of increasing the
(-6238.4 T )
Table 11 lists all the literature data on the reaction of CO2 in aqueous DEMEA using indirect methods. It demonstrated that a large scatter existed in the literature data (see Figure 12). The k2 value (0.41 m3 mol-1 s-1) of this work, which has been obtained by extrapolating the temperature to 323 K in the Arrhenius equation (eq 18), was slightly less than that of Kim and Savage,13 with a value of 0.54 m3 mol-1 s-1 at 323 K, but higher than that of Littel et al.,14 with a value of 0.164 m3 mol-1 s-1. Thus, the k2 values of this work were determined to be in the midrange of these literature data. Both the diffusivity and solubility of CO2 in amine solutions were indirectly calculated in the two studies, using the N2O analogy. The values obtained in this work were obtained directly without involving any gas phase absorption and, therefore, correspond to the intrinsic reaction rate. Although the k2 values were determined to increase and the ka values were observed to decrease with increasing temperature,
Table 10. Reaction Rate Constants for CO2 Reaction with DEMEA in Aqueous Solution temperature (K) 298 303 308 313
reaction order, n 0.87 0.84 0.81 0.80
(18)
k2 (m3 mol-1 s-1) 10-2
7.98 × 1.18 × 10-1 1.62 × 10-1 2.18 × 10-1
ka (m6 mol-2 s-1) 10-2
8.20 × 7.98 × 10-2 9.93 × 10-3 8.82 × 10-4
kw (m6 mol-2 s-1) 5.00 × 10-2 1.00 × 10-1 2.98 × 10-2 3.58 × 10-3
Ind. Eng. Chem. Res., Vol. 46, No. 13, 2007 4433 Table 11. Literature Data for the Reaction between CO2 and Aqueous DEMEA reference Savage14
Kim and Sotelo et al.18 Little et al.15 Benitez-Garcica et al.19
temperature (K)
[DEMEA] (mol/m3)
323 283-293 293-333 298
0-900 0-25 0-2200 0-1280
the variation of kw values did not show any systematic trend (Table 12). The calculated values of 1/k2, 1/(ka[DEMEA] + kw[H2O]), ka[DEMEA], and kw[H2O] are presented in Table 12, and show the relative significance of each reaction rate constant (k2, ka, and kw) on the overall observed pseudo-first-order rate constant (k0). Comparing the ka[DEMEA] term with the corresponding kw[H2O] term in Table 12, under most of the conditions studied, the ka[DEMEA] term could be neglected in the proton removal step. This observation is consistent with the fact that, for tertiary amines, the carbamate formation would not occur. Consequently, the only path was bicarbonate formation after the formation of the zwitterion, as reported by Veldman15 and shown in the following reaction:
R3N+COO-+ H2O T R3NH+ + HCO3-
(19)
Donaldson and Nguyen17 proposed the base catalysis of the CO2 hydration for the reaction between CO2 and tertiary amines. For this system, it was clear that the second term 1/(ka[DEMEA] + kw[H2O]) was small enough to be ignored in the observed pseudo-first-order reaction rate constants, when compared with the first term (1/k2). This means that only the zwitterion
k2 (m3 mol-1 s-1) 5.10 × 10-1 1.75 × 108 exp(-4749/T) 3.61 × 106 exp(-5431/T) 2.10 × 10-2
Table 12. Contribution of the Different Kinetic Constants to the Overall Observed Rate Constant for CO2 Reaction in Aqueous DEMEA Solution Contribution [DEMEA] (mol/m3)
298 K
303 K
308 K
313 K
1.25 × 101
1/k2 Term 8.47 × 100
6.16 × 100
4.58 × 100
196.52 300.80 402.21 502.18 608.78 758.17 997.36
1/(ka[DEMEA] + kw[H2O]) Term 3.76 × 10-4 1.88 × 10-4 6.35 × 10-4 3.85 × 10-4 1.93 × 10-4 6.52 × 10-4 3.95 × 10-4 1.98 × 10-4 6.69 × 10-4 4.04 × 10-4 2.03 × 10-4 6.87 × 10-4 4.20 × 10-4 2.11 × 10-4 7.16 × 10-4 4.32 × 10-4 2.18 × 10-4 7.41 × 10-4 4.60 × 10-4 2.34 × 10-4 7.95 × 10-4
196.52 300.80 402.21 502.18 608.78 758.17 997.36
1.61 × 101 2.47 × 101 3.30 × 101 4.12 × 101 4.99 × 101 6.22 × 101 8.18 × 101
196.52 300.80 402.21 502.18 608.78 758.17 997.36
2.64 × 103 2.57 × 103 2.50 × 103 2.43 × 103 2.33 × 103 2.25 × 103 2.09 × 103
ka[DEMEA] Term 1.57 × 101 1.95 × 100 2.40 × 101 2.99 × 100 3.21 × 101 3.99 × 100 4.01 × 101 4.99 × 100 4.86 × 101 6.04 × 100 6.05 × 101 7.53 × 100 7.96 × 101 9.90 × 100 kw[H2O] 5.30 × 103 5.16 × 103 5.02 × 103 4.88 × 103 4.68 × 103 4.52 × 103 4.20 × 103
1.57 × 103 1.53 × 103 1.49 × 103 1.45 × 103 1.39 × 103 1.34 × 103 1.25 × 103
5.29 × 10-3 5.43 × 10-3 5.58 × 10-3 5.73 × 10-3 5.97 × 10-3 6.18 × 10-3 6.64 × 10-3 1.73 × 10-1 2.65 × 10-1 3.55 × 10-1 4.43 × 10-1 5.37 × 10-1 6.69 × 10-1 8.80 × 10-1 1.89 × 102 1.84 × 102 1.79 × 102 1.74 × 102 1.67 × 102 1.61 × 102 1.50 × 102
formation step had a significant role in the overall reaction, under the conditions of concentration and temperature that have been studied. Therefore, the pseudo-first-order rate constant k0 can be approximated to the following:
k0 ) k2[R3NH2]
(20)
This further confirmed the validity of the reaction mechanism of CO2 in tertiary amines, as proposed by Donaldson and Nguyen.16 Figure 13 showed that the reaction rate of DEMEA with CO2 was much faster than that of MDEA.
Figure 12. Comparison of the k2 values for CO2 in aqueous DEMEA solution obtained by different researchers.
Figure 13. Comparison of the k2 values for DEMEA and MDEA.
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Ind. Eng. Chem. Res., Vol. 46, No. 13, 2007
5. Conclusions For the three amines investigated, the overall reaction rates were determined to be dependent on the amine concentration, the solvent type, and the temperature. The sequences of the systems, in terms of k0 values were EDA > MEA, EEA > DEA, and DEMEA > MDEA, under the conditions studied in this work. The reaction order was practically 1 for EDA and DEMEA and 2 for EEA for the low amine concentration used in this study. The primary and secondary amines (EDA and EEA) reacted via the zwitterion mechanism, resulting in the formation of carbamate ions. The standard reaction rate expression (eq 9) proposed by Caplow and Danckwerts was used to calculate the k0 values. Both the zwitterion formation step and the zwitterion deprotonation step had significant roles for EDA and EEA. EDA competed with water in the proton removal step. EEA provided the major contribution to the proton removal step. With a higher capacity for CO2, EDA (a diamine) could be a good substitute to MEA, because it reacted faster, from the kinetics point of view. EEA reacted faster than DEA, under the conditions studied. The reaction of CO2 with the tertiary amine (DEMEA) was faster than that with MDEA. The zwitterion formation step had a significant role in the overall reaction. The base catalysis of the CO2 hydration mechanism could explain the reaction between CO2 and this tertiary amine. Therefore, the three proposed amines were determined to be promising for the field of gas sweetening. Nomenclature A ) Arrhenius constant (m3 mol-1 s-1) B ) base (amine, water, or hydroxyl ion) Eact ) activation energy (kJ/mol) k0 ) observed pseudo-first-order reaction rate constant (s-1) k-1 ) backward first-order reaction rate constant (s-1) k2 ) forward second-order reaction rate constant for the formation of the zwitterion (m3 mol-1 s-1) kBase ) rate constant for the deprotonation of the zwitterion by a base (amine, water, or hydroxyl ion) (m3 mol-1 s-1) kB ) equal to k2kBase/k-1 (m6 mol-2 s-1) ka ) equal to k2kamine/k-1 (m6 mol-2 s-1) kw ) equal to k2kwater/k-1 (m6 mol-2 s-1) MRD ) mean relative deviation rCO2 ) rate of reaction, with respect to CO2 (mol m-3 s-1) R ) universal gas constant (0.008315 kJ mol-1 K-1) T ) temperature (K) [ ] ) concentration (mol/m3) Acknowledgment The financial support provided by the Natural Science and Engineering Research Council of Canada (NSERC) in the form
of a Discovery and a Strategic Grants and from CANMET (Energy Technology Centre, Natural Resources Canada, Ottawa) is gratefully acknowledged. Literature Cited (1) Jensen, A.; Christensen, R. Determination of Some Inorganic Substances in the Labyrinthine Fluids. Acta Chem. Scand. 1955, 9, 486. (2) Weiland, R. H.; Trass, O. Absorption of Carbon Dioxide in Ethylenediamine Solutions. Can. J. Chem. Eng. 1971, 49, 767. (3) Hikita, H.; Asai, S.; Ishikawa, H.; Honda, M. The Kinetics of Reactions of Carbon Dioxide with Monoisopropanolamine, Diglycolamine and Ethylenediamine by a Rapid Mixing Method. Chem. Eng. J. 1977, 14, 27. (4) Alem, M. Study of the Rates of Absorption and Equilibrium Solubility of Carbon Dioxide in Aqueous Amine Solutions, M.A.Sc. Thesis, University of Regina, Regina, Canada, 2005. (5) Bavbek, O.; Alper, E. Reaction Mechanism and Kinetics of Aqueous Solutions of Primary and Secondary Alkanolamines and Carbon Dioxide. Turk. J. Chem. 1999, 23, 293. (6) Versteeg, G. F.; van Swaaij, W. P. M. On the Kinetics Between CO2 and Alkanolamines both in Aqueous and Non-aqueous SolutionssI. Primary and Secondary Amines. Chem. Eng. Sci. 1988, 43, 573. (7) Versteeg, G. F.; van Dijk, L. A.; van Swaaij, W. P. M. On the Kinetics of CO2 and Alkanolamines Both in Aqueous and Non-aqueous solutions: An Overview. Chem. Eng. Commun. 1996, 144, 113. (8) Caplow, M. Kinetics of Carbamate Formation and Breakdown. J. Am. Chem. Soc. 1968, 90, 6795. (9) Danckwerts, P. V. The Reactions of CO2 with Ethanolamines. Chem. Eng. Sci. 1979, 34, 443. (10) Knipe, A. C.; McLean, D.; Tranter, N. L. J. Phys. E 1974, 7, 56. (11) Ali, S. H. Kinetic of the Reaction of Carbon Dioxide with Blends of Amines in Aqueous Media Using the Stopped-Flow Technique. Int. J. Chem. Kinetics 2005, 37(7), (391) (12) Katchalski, E.; Berliner-Klibanski, C.; Berger, A. The Chemical Structure of Some Diamine Carbamates. J. Am. Chem. Soc. 1951, 73, 1829. (13) Kim, C. J.; Savage, D. W. Kinetics of Carbon Dioxide Reaction with Diethyl-amino-ethanol in Aqueous Solutions. Chem. Eng. Sci. 1987, 42, 1481. (14) Littel, R. J.; van Swaaij, W. P. M.; Versteeg, G. F. Kinetics of Carbon Dioxide with Tertiary Amines in Aqueous Solution. AIChE J. 1990, 36 (11), 1633. (15) Veldman, R. R. Alkanolamine Solution Corrosion Mechanisms and Inhibition from Heat Stable Salts and CO2. Presented at Corrosion 2000, Houston, TX, 2000; Paper 00496. (16) Donaldson, T. L.; Nguyen, Y. N. Carbon Dioxide Reaction Kinetics and Transport in Aqueous Amine Membrane. Ind. Eng. Chem. Fundam. 1980, 19, 260. (17) Sotelo, J. L.; Benitez, F. J.; Beltran-Heredia, J.; Rodriquez, C. Cinetica de la Reaccion de Dioxido de Carbono con Caminas. II. Amina Terciaria: Dietiletonalamina. Ann. Quim. 1990, 87, 212. (18) Benitez-Garcia, J.; Ruiz-Ibanez, G.; Al-Ghawas, H. A. and Sandall, O. C. On the Effect of Basicity on the Kinetics of Reaction CO 2 Absorption in Tertiary Amines. Chem. Eng. Sci. 1991, 46, 2927.
ReceiVed for reView November 23, 2006 ReVised manuscript receiVed April 4, 2007 Accepted April 10, 2007 IE0614982
