Kinetics of Absorption of Carbon Dioxide into Aqueous Solutions of

Ind. Eng. Chem. Res. 2002, 41, 257-266

257

Kinetics of Absorption of Carbon Dioxide into Aqueous Solutions of Monoethanolamine + Triethanolamine Shyh-Yun Horng and Meng-Hui Li* Department of Chemical Engineering, Chung Yuan Christian University, Chung Li, Taiwan 32023, Republic of China

Kinetics of the absorption of CO2 into monoethanolamine (MEA) + triethanolamine (TEA) + water were investigated at 30, 35, and 40 °C using a laboratory wetted wall column. Ten systems of which 0.5 and 1.0 kmol m-3 TEA mixed with various MEA concentrations (0.1, 0.2, 0.3, 0.4, and 0.5 kmol m-3) were studied. Densities and viscosities of 10 blended-amine systems and solubilities and diffusivities of N2O in the solutions were measured. The N2O analogy was applied to estimate the solubilities and diffusivities of CO2 in amine systems. On the basis of the fast pseudo-first-order rate for CO2 absorption, the overall pseudo-first-order reaction rate constants were determined from kinetic measurements. The addition of small amounts of MEA to TEA results in a significant enhancement of CO2 absorption rates. A hybrid reaction rate model, a zwitterion mechanism for MEA and a first-order reaction for TEA, was used to model the kinetic data. The model is satisfactory to represent CO2 absorption rates in TEA + MEA aqueous systems. Introduction Solutions of alkanolamines are an industrially important class of compounds used in the natural gas, petroleum chemical plants, and ammonia industries for the removal of CO2 and H2S from gas streams. A wide variety of alkanolamines such as monoethanolamine (MEA), diethanolamine (DEA), di-2-propanolamine (DIPA), and N-methyldiethanolamine (MDEA) have been used industrially for a number of years.1 Owing to advantages of absorption capacity, absorption rate, and degradation resistance over conventional amines for CO2 removal from gases, sterically hindered amines have been recently suggested to be used in gas-treating processes.2 The absorption of acid gases in blended amines has specific advantages over the use of single amines. The addition of a small amount of primary amine to conventional tertiary amines can enhance the rate of absorption of CO2 to a large extent without appreciably affecting the stripping characteristics.3 Blends of primary and tertiary amines, such as MEA + MDEA + H2O, have been suggested for CO2 removal.3 Because of the increasing importance of blendedamine systems in acid-gas-treating processes, it is necessary to have an understanding of the kinetic phenomena in mixed-amine systems. Versteeg et al.4 conducted a study of CO2 absorption into aqueous mixed-amine solutions such as MMEA (methylmonoethanolamine) + MDEA, MEA + MDEA, DIPA + MDEA, and MEA + DEA + MDEA. Parallel reversible reactions according to both film and penetration models were used to model the CO2 absorption process. When the tertiary alkanolamine was assumed to completely deprotonate the primary alkanolamine, the predicted results were in good agreement with the experimental * Corresponding author. E-mail address: [email protected]. Phone: 886 03 4563171 ext 4110. Fax: 886 03 4563171 ext 4199.

results. For the reaction kinetics of CO2 absorption into MEA + TEA + H2O, only a few investigations are available in the literature. Rangwala et al.5 studied the absorption of CO2 into aqueous blends of MEA and TEA solutions in a stirred cell absorber at 20 °C. A modified pseudo-first-order model based on the film theory is used to predict the rate of absorption of CO2 into mixedamine solutions. The model accounts for the variation of amine concentration in the film and assumes a shuttle mechanism for rate enhancement. Bulk liquid concentrations of the various species present are obtained from a simplified thermodynamic model. The model predicts absorption rates that are in agreement with experimental measurements. The purpose of this work is to determine experimentally the reaction kinetics for the absorption of CO2 into MEA + TEA + H2O for temperatures over 30-40 °C. Experimental kinetic data were collected, and a hybrid kinetics model was used to interpret the results. For the rational design of the gas absorption units, physical properties such as solubility and the diffusivity of acid gases in amine solutions are required to model the rate of absorption. Because of the chemical reaction between CO2 and amines, neither the free-gas solubility nor the diffusivity of CO2 in amine solutions can be measured directly. The N2O analogy has been frequently used to estimate the solubility and the diffusivity of CO2 in amine solutions.6-10 The equations for solubility and diffusivity of N2O and CO2 in water proposed by Versteeg and van Swaaij7 will be applied directly in this study to calculate the corresponding solubility and diffusivity of CO2 in an amine solution. Theory Reactions of CO2 in Aqueous Solutions. For the reactions of CO2 in aqueous solutions, the first reaction to be considered is the hydration of CO2

CO2 + H2O T HCO3- + H+

10.1021/ie010671l CCC: $22.00 © 2002 American Chemical Society Published on Web 12/14/2001

(1)

258

Ind. Eng. Chem. Res., Vol. 41, No. 2, 2002

This reaction is very slow (k ) 0.026 s-1 at 25 °C)11 and usually is negligible.12 The second reaction is that of bicarbonate formation

CO2 + OH- T HCO3-

(2)

This reaction is fast and can enhance mass transfer even when the concentration of hydroxyl ion is low. The forward reaction can be described as11

rCO2-OH- )

/ kOH -[CO2][OH ]

(3)

/ 3 -1 -1 s )) ) 13.635 - 2895/(T/K) log(kOH -/(m kmol (4)

Reaction Rate Dependence on MEA. The carbamate formation reaction occurs when CO2 reacts with primary and secondary alkanolamines

CO2 + 2R1R2NH T R1R2NCOO- + R1R2NH2+

(5)

where R1 is an alkyl group and R2 is Η for primary amines and an alkyl group for secondary amines. The zwitterion mechanism originally proposed by Caplow13 and reintroduced by Danckwerts14 is generally accepted as the reaction mechanism for the carbamate formation between CO2 with primary and secondary alkanolamines. The zwitterion mechanism has been used successfully in aqueous alkanolamine solutions12,15 as well as in some organic and viscous solutions.16-19 The reaction steps successively involve the formation of a zwitterion k2

8 R1R2NH+COOCO2 + R1R2NH 79 k -1

(6)

and the subsequent removal of the proton by a base B (base catalysis) kb

R1R2NH+COO- + B 98 R1R2NCOO- + BH+ (7) where B is a base that could be an amine, OH-, or H2O.12 For this mechanism, Danckwerts14 derived the forward reaction rate equation at quasi-steady state

Because the base could be MEA, TEA, OH-, or H2O, the forward reaction rate for CO2-MEA (i.e., eq 8) becomes rCO2-MEA ) 1+

k2,MEA[CO2][MEA] k-1 kH2O[H2O] + kOH-[OH-] + kMEA[MEA] + kTEA[TEA]

(9) Reaction Rate Dependence on TEA. For the reaction of CO2 with tertiary alkanolamines (R3N), Donaldson and Nguyen20 proposed the following reaction mechanism:

CO2 + R3N + H2O T R3NH+ + HCO3-

This reaction mechanism is essentially a base-catalyzed hydration of CO2, and the mechanism implies that tertiary amines cannot react directly with CO2. In most of the literature on CO2 kinetics with tertiary amines in aqueous solutions, it is assumed that the reaction of CO2 with TEA is a pseudo-first-order reaction5,21-23 as follows:

rCO2-TEA ) k2,TEA[CO2][TEA]

rov ) rCO2-MEA + rCO2-TEA + rCO2-OH-

1+

(8)

kH2O[H2O] + kOH-[OH-] + kR1R2NH[R1R2NH]

The zwitterion mechanism covers the shifting reaction orders (i.e., the overall reaction order changes between two and three) for the reactions of CO2 with different primary and secondary alkanolamines. Because the zwitterion mechanism is generally accepted as the reaction mechanism for the carbamate formation between CO2 with primary and secondary alkanolamines, the zwitterion mechanism will be applied in this study to the reaction between CO2 and MEA for CO2 absorption into MEA + TEA + H2O.

(12)

After the substitution of the zwitterion mechanism for MEA (i.e., eq 9) and the first-order reaction for CO2TEA (i.e., eq 11) into eq 12, one has rov ) kov[CO2] ) 1+

k2,MEA[CO2][MEA] k-1

+

kH2O[H2O] + kOH-[OH-] + kMEA[MEA] + kTEA[TEA] / k2,TEA[CO2][TEA] + kOH -[CO2][OH ] (13)

and the apparent reaction rate constant, kapp, is defined as / kapp ) kov - kOH -[OH ]

k-1

(11)

Reaction Rate for CO2 Absorption into MEA + TEA + H2O. For the absorption of CO2 into MEA + TEA + H2O, the CO2 overall reaction rate can be expressed as follows:

rCO2-AMINE ) k2,R1R2NH[CO2][R1R2NH]

(10)

(14)

The apparent reaction rate constant has the following expression:

[

((

k2,MEAkH2O [H2O] + k-1 k2,MEAkMEA k2,MEAkTEA k2,MEAkOH[OH-] + [MEA] + k-1 k-1 k-1

kapp ) [MEA]/ (1/k2,MEA) + 1/

[TEA]

))]

+ k2,TEA[TEA] (15)

Experimental Section Alkanolamine aqueous solutions were prepared from distilled water. The distilled water was degassed by

Ind. Eng. Chem. Res., Vol. 41, No. 2, 2002 259

boiling. The alkanolamines are Riedel-de Hae¨n reagent grade with purities of MEA at 99% (mol) and TEA at 99% (mol). The concentration of the alkanolamine aqueous solution was determined by titration of a liquid sample with HCl using methyl orange as the indicator. The alkanolamine aqueous solutions were found to be within (0.2 mass % of the stated concentration. Density Measurement. The densities of solutions were measured by using a 25-mL Gay-Lussac pycnometer. The measurements were performed in a constanttemperature water bath, in which the temperature could be held constant to (0.05 °C. The experimental errors were estimated to be equal to (0.05% on the basis of comparisons with literature data. The method for density measurements is the same as that used in our previous work.24 Viscosity Measurement. The kinematic viscosities of solutions were measured by means of Cannon-Fenske routine viscometers over the temperature range of 3080 °C. The measurements were performed in a constanttemperature water bath, in which the temperature could be held constant to (0.05 °C. An electronic stopwatch with an accuracy of 0.01 s was used to measure the efflux times of the liquid samples. The kinematic viscosity of the solution was calculated from the multiplication of the efflux time with the viscometer constant. The absolute viscosity of fluid can be obtained by multiplying the kinematic viscosity by the density of the fluid. The accuracy of the viscosities was estimated to be (1.0% on the basis of comparisons with literature data. The procedure for viscosity measurement is the same as that in our previous work.25 Solubility Measurement. The solubility of N2O in amine aqueous solutions was measured by using a solubility apparatus, similar to those presented by AlGhawas et al.8 and Haimour.26 The accuracy of the temperature of the system is estimated to be (0.5 °C. The reproducibility between the various experiments is always within 3%. The estimated experimental error in the measured solubility is about (3%. The apparatus, experimental procedure, and methods of analysis are essentially the same as those used in our previous work.27 Diffusivity Measurement. The diffusivity of a gas in amine solutions was measured in a short wetted wall column absorber. A short stainless steel cylinder of an outside diameter (o.d.) of 2.54 cm and a height of 10 cm was used as the wetted wall column. The liquid flowed through an annular distributor cap and distributed uniformly as a thin film on the outside of the cylinder. The liquid flow rate was varied in the range of 3-6 cm3 s-1. The gas flow rate was measured by using a soapfilm meter. The absorption rate was measured by the difference between the inlet and outlet gas flow rates. The system was first purged with the saturated gas. The solution was heated to the desired temperature and flowed through the wetted wall column. The wetted wall column was adjusted to eliminate any ripples that it may have on the liquid surface. When the solution distributed uniformly as a thin film on the outside of the cylinder, the liquid flow rate and the gas absorption rate were measured. In physical absorption of a sparingly solute gas at a short contact time and with an initial gas-free liquid, the Higbie penetration theory gives the specific

absorption rate as

( )()

NA ) 2

DA πtc

1/2

pA H

(16)

where the contact time tc can be derived from wetted wall column hydrodynamics as

tc )

2h πd 2/3 3η 1/3 3 L Fg

( ) ( )

(17)

After the combination of eqs 16 and 17, one has

DA ) (NAH)2πtc/(2pA)2

(18)

The diffusivity of the gas in solution can be calculated by Henry’s constant, the specific gas absorption rate, the contact time, density, and viscosity of the solution. To calibrate the wetted wall column absorber, the diffusivities of N2O in water and CO2 in water were measured. The estimated experimental error in the measured diffusivity is estimated as (2%. The apparatus and the experimental procedure are the same as those described by Li and Lee.27 Kinetic Data Measurement. A short wetted wall 10-cm-long column with a 2.54 cm o.d. and was used for the kinetic data measurement, with the same apparatus as that used for the diffusivity measurement. For the condition of low partial pressure of CO2, nitrogen was introduced into the gas flow, and the partial pressure of CO2 was determined by gas chromatography. The CO2 absorption rate was measured from the difference in CO2 between inlet and outlet gas flows. The CO2 loading of the amine solution was determined by the CO2 absorption rate, the liquid flow rate, and the concentration of the amine solution. The apparatus and the experimental procedure are the same as those described by Shen et al.28 The conditions for the absorption of CO2 in amine solutions were selected in such a way as to ensure that absorption occurred in the fast pseudo-first-order reaction regime which requires that

3 < Ha 3, tanh Ha is approaching 1; thus, the specific rate of mass transfer of CO2 becomes

NA ) [CO2]ixDCO2kov )

pCO2

D k HCO2x CO2 ov

(22)

The specific absorption of CO2 into MEA + TEA + H2O was measured in various systems for temperatures

260


Table 1. Densities and Viscosities of MEA + TEA + H2O t (°C)

kmol m-3 MEA + kmol m-3 TEA

density F (g cm-3)

viscosity η (mPa s)


density F (g cm-3)

viscosity η (mPa s)

30

0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5

1.0067 1.0069 1.0072 1.0073 1.0076 1.0052 1.0053 1.0055 1.0057 1.0061 1.0033 1.0035 1.0038 1.0041 1.0042

1.028 1.045 1.065 1.085 1.105 0.918 0.934 0.951 0.967 0.985 0.826 0.839 0.855 0.870 0.884

0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0

1.0179 1.0182 1.0184 1.0187 1.0188 1.0161 1.0164 1.0166 1.0168 1.0170 1.0140 1.0143 1.0145 1.0147 1.0150

1.313 1.339 1.372 1.403 1.421 1.161 1.187 1.208 1.239 1.256 1.037 1.057 1.085 1.103 1.119

35

40

of 30, 35, and 40 °C. kov can then be calculated from the diffusivity of CO2 in solution, the partial pressure of CO2, and the Henry constant of CO2 using eq 22. k2,MEA, k2,TEA, k2,MEAkH2O/k-1, k2,MEAkOH-/k-1, k2,MEAkMEA/ k-1, and k2,MEAkTEA/k-1 in the expression of kapp (i.e., eq 15) can be determined from the measured value of kov.

viscosity of MEA + TEA + H2O increases as the concentration of the MEA increases, and it decreases as the temperature increases. The viscosities of MEA + TEA + H2O were correlated using a Redlich-Kistertype equation for the viscosity deviation.25 The viscosity deviation of aqueous blends of alkanolamine solutions is assumed to be

Results and Discussion The aqueous blended-amine systems chosen for the CO2 kinetics study are 0.5 and 1.0 kmol m-3 TEA with various MEA concentrations: 0.1, 0.2, 0.3, 0.4, and 0.5 kmol m-3. The temperature range is 30, 35, and 40 °C. The density of MEA + TEA + H2O was measured in various temperatures of 30, 35, and 40 °C, and the results are presented in Table 1. The density of MEA + TEA + H2O increases as the concentration of MEA increases, and it decreases as the temperature increases, as shown in Table 1. The density of MEA + TEA + H2O was correlated using the density correlation24 for the excess molar volume as follows:

VE ) VE12 + VE13 + VE23

(23)

where the superscript E denotes the excess molar volume. A Redlich-Kister-type equation is used for the excess molar volume as follows:

VE12/(cm3 mol-1) ) x1x2

n

Ai(x1 - x2)i ∑ i)0

(24)

where the temperature dependence of A is assumed as

Ai ) ai,1 + ai,2(T/K) + ai,3(T/K)2

(25)

The deviation of the density calculation for MEA + TEA + H2O is 0.1%. In the calculation, the parameters for the pair MEA + TEA are determined as a0,1 ) 48 348.3, a0,2 ) -296.11, a0,3 ) 0.452 46, a1,2 ) 596.193, a1,3 ) -1.8586, and a2,3 ) 3.8052; the other parameters are adopted from the literature values (pure fluids are H2O,24 MEA,24 and TEA;30 binary pairs are MEA + H2O24 and TEA + H2O31). The viscosities of MEA + TEA + H2O for the temperatures ranging from 30-40 °C were measured and are also presented in Table 1. At a constant temperature, the viscosity of MEA + TEA (1.0 M) + H2O is larger than that of MEA + TEA (0.5 M) + H2O, as in Table 1. Also, at the same concentration of TEA, the

δν ) δνE12 + δνE13 + δνE23

(26)

δνEij is the excess kinematic viscosity deviation32 and is assumed to have the Redlich-Kister-type expression m

Ai(x1 - x2)i ∑ i)0

δν12 ) x1x2

(27)

where Ai values are pair parameters and are assumed to have the following temperature dependence:

Ai ) ai,1 +

ai,2 (T/K) + ai,3

(28)

The deviation of the viscosity calculation is 0.3%. In the calculation, the parameters for the MEA + TEA pair are determined as a0,1 ) -102.41, a0,2 ) 18 810, a0,3 ) -165, and a1,1 ) 5718; the other parameters are adopted from the literature values (pure fluids are H2O,25 MEA,25 and TEA;31 binary pairs are MEA + H2O25 and TEA + H2O31). To confirm the apparatus and the experimental procedure of the solubility measurement, the solubilities of N2O in water were measured. The measured solubilities of N2O in water at 30, 35, and 40 °C are 4579, 5118, and 5919 kPa m3 kmol-1, respectively. In Figure 1, a comparison between the literature values7-9,26,33 and the values obtained in this study for the solubility of N2O in water is shown. The solid line in Figure 1 is for the calculated values using the correlation of Versteeg and van Swaaij.7 As can be seen from Figure 1, the solubility of N2O in water measured in this study is in good agreement with the literature values. The results for the solubility of N2O in MEA + TEA + H2O at 30, 35, and 40 °C are presented in Table 2. At a constant temperature, the solubility of N2O in MEA + TEA + H2O increases systematically as the concentration of MEA increases. Also, the solubility of N2O in MEA + TEA (1.0 M) + H2O is larger than that of corresponding MEA + TEA (0.5 M) + H2O at the same temperature,

Ind. Eng. Chem. Res., Vol. 41, No. 2, 2002 261 Table 2. Estimated Solubility of CO2 in MEA + TEA + H2O Using N2O Analogy t (°C) 30

35

40

kmol m-3 MEA + kmol m-3 TEA 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5

kPa m3 kmol-1 HN2O HCO2 4676 4716 4774 4853 4891 5312 5343 5371 5420 5441 5844 5900 5952 5968 6023

3411 3441 3483 3540 3568 3825 3848 3868 3903 3918 4157 4197 4234 4245 4284

kPa m3 kmol-1 H N 2O HCO2


4777 4835 4891 4941 5001 5398 5439 5490 5546 5598 5961 5989 6063 6099 6156

3484 3527 3568 3604 3648 3887 3917 3954 3994 4032 4240 4260 4313 4338 4379

Figure 1. Solubility of N2O in water as a function of temperature. The solid line is calculated using the correlation of Versteeg and van Swaaij.7

Figure 2. Diffusivity of N2O in water as a function of temperature. The solid line is calculated using the correlation of Versteeg and van Swaaij.7

as given in Table 2. Using the N2O analogy, the solubilities of CO2 in MEA + TEA + H2O can be calculated and are presented Table 2. The diffusivity of N2O in an amine solution was measured by the wetted wall column absorber. To confirm the apparatus and the experimental procedure, the diffusivities of N2O in water were measured at 30, 35, and 40 °C. The measured diffusivities of N2O in water at 30, 35, and 40 °C are 2.00 × 10-9, 2.37 × 10-9, and 2.58 × 10-9 m2 s-1, respectively. A comparison between the literature values6-8,33,34 and the values obtained in this study for the diffusivity of N2O in water is also shown in Figure 2. The solid line in Figure 2 is for the calculated values using the correlation of Versteeg and van Swaaij.7 As shown in Figure 2, the values obtained in this study for the diffusivity of N2O in water are in good agreement with the literature values. The diffusivities of N2O in MEA + TEA + H2O are presented in Table 3. Because of the higher viscosities of MEA + TEA (1.0 M) + H2O at a constant temperature, the diffusivity of N2O in MEA + TEA (1.0 M) + H2O is smaller that that of corresponding MEA + TEA (0.5 M) + H2O, as given in Table 3. Also, the diffusivity of N2O in MEA + TEA + H2O decreases as the concentration of MEA increases at the same temperature, and it

increases as the temperature increases at the same amine concentration. The diffusivities of CO2 in MEA + TEA + H2O are estimated using the N2O analogy and are presented in Table 3. The specific absorption rate of CO2 in an amine solution was measured by the wetted wall column absorber. The specific rates of absorption of CO2 into MEA + TEA + H2O are presented in Tables 4 and 5. From eq 22, the overall pseudo-first-order rate constant kov can be calculated from the partial pressure of CO2, the Henry constant, and the diffusivity of CO2 in MEA + TEA + H2O. The calculated kov and kapp are presented in Tables 6 and 7. The Hatta number and the enhancement factor are also presented in Tables 6 and 7. The reaction rate constant for the reaction of the bicarbonate formation was estimated by eq 4. The hydroxyl ion concentration was estimated from the relation given35 as

[OH-] ) )

Kw 1 - R , R g 10-3 Kp R

(

x

)

Kw [AM], R e 10-3 Kp

(29)

262


Table 3. Estimated Diffusivity of CO2 in MEA + TEA + H2O Using N2O Analogy t (°C)

D (109 m2 s-1) DN2O DCO2


30

35

40

1.78 1.69 1.63 1.59 1.54 2.07 1.97 1.83 1.75 1.70 2.25 2.15 2.11 2.03 2.00

D (109 m2 s-1) DN2O DCO2


1.89 1.80 1.74 1.69 1.64 2.17 2.07 1.92 1.84 1.79 2.33 2.22 2.19 2.10 2.07

1.55 1.45 1.34 1.23 1.09 1.69 1.61 1.50 1.40 1.24 2.01 1.86 1.76 1.52 1.49

1.65 1.54 1.43 1.31 1.16 1.78 1.69 1.58 1.47 1.30 2.08 1.93 1.83 1.58 1.54

Table 4. Kinetics Data Obtained for CO2 in a MEA + TEA (0.5 M) + H2O System t (°C)


pA (kPa)

L (107 m3 s-1)

φ (108 kmol s-1)

tc (s)

R (102 kmol of CO2/ kmol of amine)

NA (106 kmol m-2 s-1)

kL (105 m s-1)

30

0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5

19.9 18.8 15.0 15.4 13.3 19.0 16.0 14.3 14.5 13.1 17.6 15.0 15.2 13.0 11.3

6.21 8.83 8.70 6.33 10.2 6.80 8.94 6.62 6.86 7.77 11.7 14.1 7.73 10.8 7.77

3.36 5.29 5.33 6.07 6.09 3.15 4.67 4.98 5.99 6.48 3.19 4.55 5.44 5.64 5.93

1.37 1.09 1.11 1.38 1.01 1.24 1.04 1.28 1.26 1.16 0.84 0.74 1.12 0.90 1.12

9.02 8.56 7.66 10.7 5.99 7.72 7.46 9.40 9.70 8.34 4.53 4.62 8.80 5.81 7.63

3.50 5.50 5.55 6.32 6.33 3.28 4.86 5.19 6.24 6.75 3.32 4.73 5.67 5.87 6.18

4.19 4.58 4.47 3.95 4.54 4.72 5.02 4.37 4.31 4.42 5.97 6.17 5.00 5.46 4.84

NA (106 kmol m-2 s-1)

kL (105 m s-1)

2.47 4.13 5.00 5.90 5.20 3.61 4.76 5.04 5.92 6.26 3.55 4.57 5.05 5.68 6.18

3.91 4.04 3.85 3.60 3.37 4.69 4.51 4.41 4.14 3.74 4.82 5.15 4.39 4.28 4.45

35

40

Table 5. Kinetics Data Obtained for CO2 in a MEA + TEA (1.0 M) + H2O System t (°C)


pA (kPa)

(107 m3 s-1)

30

0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0

11.3 14.8 14.9 16.3 13.2 17.2 16.0 15.4 14.4 13.6 16.7 15.4 14.1 13.2 11.7

6.93 8.54 8.40 7.88 7.84 10.1 9.82 10.2 9.57 8.49 8.18 11.3 7.72 8.95 10.5

35

40

L

(108

φ kmol s-1)

tc (s)

R (102 kmol of CO2/ kmol of amine)

2.37 3.97 4.81 5.67 5.00 3.47 4.58 4.85 5.69 6.02 3.41 4.40 4.85 5.46 5.94

1.38 1.21 1.23 1.29 1.30 1.03 1.06 1.03 1.09 1.19 1.14 0.93 1.20 1.10 0.99

3.11 3.87 4.40 5.14 4.25 3.13 3.89 3.64 4.25 4.73 3.79 3.24 4.83 4.36 3.76

Table 6. Kinetic Data for the Absorption of CO2 in MEA + TEA (0.5 M) + H2O t (°C) 30

35

40


[H2O] (kmol m-3)

kov (s-1)

kOH-*[OH-] (s-1)

kapp (s-1)

Ha

E

0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5

51.0 50.7 50.4 50.0 49.7 51.0 50.7 50.4 50.0 49.7 51.0 50.7 50.4 50.0 49.7

190.4 562.7 948.4 1261 1757 202.1 665.2 1023 1531 2291 262.0 786.3 1137 1749 2656

0.11 0.11 0.13 0.09 0.17 0.21 0.22 0.17 0.16 0.19 0.59 0.58 0.29 0.46 0.34

190.3 562.6 948.3 1261 1757 201.9 665.0 1023 1531 2291 261.4 785.7 1137 1749 2656

14.3 22.0 28.7 36.9 37.4 14.0 23.3 32.1 38.9 45.8 13.1 21.4 31.6 35.1 48.4

14.3 22.0 28.7 36.9 37.4 14.0 23.3 32.1 38.9 45.8 13.1 21.4 31.6 35.1 48.4

Ind. Eng. Chem. Res., Vol. 41, No. 2, 2002 263 Table 7. Kinetic Data for the Absorption of CO2 in MEA + TEA (1.0 M) + H2O t (°C) 30

35

40


[H2O] (kmol.m-3)

kov (s-1)

k*OH-[OH-] (s-1)

kapp (s-1)

Ha

E

0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0

46.9 46.6 46.2 45.9 45.5 46.9 46.6 46.2 45.9 45.5 46.9 46.6 46.2 45.9 45.5

348.7 623.3 1003 1303 1780 372.4 801.8 1068 1823 2668 391.0 834.3 1301 2216 3498

0.33 0.27 0.23 0.20 0.24 0.54 0.43 0.46 0.39 0.35 0.71 0.84 0.55 0.62 0.72

348.4 623.1 1003 1303 1780 371.9 801.3 1068 1823 2668 390.3 833.4 1300 2215 3497

19.4 24.3 31.1 36.4 42.7 17.4 25.8 29.4 39.5 49.8 18.7 24.6 35.1 43.7 52.1

19.4 24.3 31.1 36.4 42.7 17.4 25.8 29.4 39.5 49.8 18.7 24.6 35.1 43.7 52.1

Figure 3. Apparent rate constant for the reaction of CO2 with MEA + TEA (0.5 M) + H2O as a function of MEA concentration. Solid lines are calculated using eq 15.

Figure 4. Apparent rate constant for the reaction of CO2 with MEA + TEA (1.0 M) + H2O as a function of MEA concentration. Solid lines are calculated using eq 15.

where R is the CO2 loading in an amine solution. The expression for Kw given by Barth et al.36 is applied

pseudo-first-order reaction rate for the CO2-MEA reaction will not be able to accurately represent the kinetic data for CO2 absorption in MEA + TEA + H2O. Consequently, the zwitterion mechanism for the reaction of CO2 with MEA in MEA + TEA + H2O will then be an appropriate choice. The addition of small amounts of MEA to TEA results in a significant increase in the CO2 absorption rates. kapp for CO2 absorption into TEA (1.0 M) + H2O is 4.78, 6.05, and 7.61 s-1 for 30, 35, and 40 °C,22 respectively. kapp for CO2 absorption into MEA (0.5 M) + TEA (1.0 M) + H2O is 1780, 2668, and 3498 s-1, as given in Table 7, for 30, 35, and 40 °C, respectively. The ratios of kapp for CO2 absorption into MEA (0.5 M) + TEA (1.0 M) + H2O to that for CO2 absorption into TEA (1.0 M) + H2O are 372, 440, and 460 for 30, 35, and 40 °C, respectively. Thus, a substantial increase in the CO2 absorption rate is observed by adding a small amount of MEA into aqueous TEA solutions. In eq 15, kapp is a function of k2,MEA, k2,TEA, k2,MEAkH2O/ k-1, k2,MEAkOH-/k-1, k2,MEAkMEA/k-1, and k2,MEAkTEA/k-1. In eq 15, the value of kapp is reduced to k2,TEA[TEA] when the concentration of MEA approaches zero. Thus, k2,TEA

log Kw ) -22.795 + 0.0294(T/K)

(30)

On the basis of the literature values,22 Kp,TEA was expressed as a function of temperature as follows:

log Kp,TEA ) -13.223 + 0.017 92(T/K)

(31)

Because the MEA concentration is small in MEA + TEA + H2O, eq 29 was assumed to be valid for the calculation of the hydroxyl ion concentration in MEA + TEA + H2O. The differences in kov and kapp are as small as 1% of kov, as shown in Tables 6 and 7. As pointed out by some investigators,12,37-39 the contribution from the reaction of the bicarbonate formation is almost negligible. The plots of kapp versus the concentration of MEA are shown in Figures 3 and 4. kapp values for CO2 absorption in TEA (0.5 M) + H2O and TEA (1.0 M) + H2O obtained from the values of k2,TEA and amine concentration of Littel et al.22 are also shown in Figures 3 and 4, respectively. As can be seen in Figures 3 and 4, kapp versus the concentration of MEA is not linear; thus, the

264


Table 8. Diffusivity of Amines in Solutions Used To Calculate the Instantaneous Enhancemnet Factor for CO2 in MEA + TEA (0.5 M) + H2O t (°C) 30

35

40


[CO2]i (103 kmol m-3)

DMEA (109 m2 s-1)

DTEA (109 m2 s-1)

E∞,MEA

E∞,TEA

E∞,total

0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5 0.1 + 0.5 0.2 + 0.5 0.3 + 0.5 0.4 + 0.5 0.5 + 0.5

5.84 5.47 4.32 4.34 3.74 4.95 4.15 3.70 3.72 3.34 4.24 3.58 3.59 3.06 2.64

1.21 1.20 1.19 1.18 1.17 1.36 1.35 1.34 1.33 1.32 1.52 1.51 1.50 1.49 1.48

0.654 0.647 0.638 0.631 0.623 0.659 0.651 0.644 0.636 0.629 0.663 0.656 0.648 0.641 0.634

14.9 31.1 58.7 78.3 114.4 17.2 40.2 68.8 92.5 129.9 20.3 47.3 70.3 111.1 161.4

52.1 56.5 71.8 72.1 84.2 57.4 69.5 79.9 80.8 90.6 64.7 77.8 77.5 92.0 106.8

67.0 87.5 130.5 150.5 198.7 74.6 109.7 148.7 173.3 220.5 84.9 125.1 147.8 203.1 268.3

Table 9. Diffusivity of Amines in Solutions Used To Calculate the Instantaneous Enhancemnet Factor for CO2 in MEA + TEA (1.0 M) + H2O t (°C) 30

35

40


[CO2]i (103 kmol m-3)

DMEA (109 m2 s-1)

DTEA (109 m2 s-1)

E∞,MEA

E∞,TEA

E∞,total

0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0 0.1 + 1.0 0.2 + 1.0 0.3 + 1.0 0.4 + 1.0 0.5 + 1.0

3.25 4.21 4.18 4.51 3.62 4.43 4.09 3.88 3.61 3.36 3.94 3.61 3.28 3.04 2.66

1.21 1.20 1.19 1.18 1.17 1.36 1.35 1.34 1.33 1.32 1.52 1.51 1.50 1.49 1.48

0.555 0.548 0.539 0.531 0.527 0.563 0.555 0.549 0.539 0.534 0.570 0.563 0.553 0.547 0.542

27.5 43.0 66.6 85.2 139.8 20.9 44.8 72.2 106.2 150.7 22.9 50.2 84.1 128.9 185.0

180.2 143.3 148.7 142.7 187.7 128.7 141.8 153.6 169.3 192.3 134.9 151.7 169.8 195.5 224.6

207.6 186.3 215.3 227.9 327.6 149.6 186.6 225.8 275.5 342.9 157.8 201.9 253.9 324.4 409.6

is set equal to the expression of Littel et al.22 as follows:

[

k2,TEA/(m3 kmol-1 s-1) ) 1.01 × 107 exp -

]

4415 (T/K) (32)

When the term k2,MEAkOH-/k-1 is neglected, the values of k2,MEA, k2,MEAkH2O/k-1, k2,MEAkMEA/k-1, and k2,MEA kTEA/ k-1 are expressed as Arrhenius equations, and the parameters are determined from the values of kapp in Tables 6 and 7. The overall absolute percentage deviation for the calculation of kapp is about 6.2% for 30 data points. The solid lines in Figures 3 and 4 are the calculated kapp values from eq 15. The determined rate constants are as follows:

k2,MEA/(m3 kmol-1 s-1) )

[

5376.2 (33) (T/K)

[

2135.1 (34) (T/K)

3.014 × 1011 exp k2,MEAkH2O /(m6 kmol-2 s-1) ) k-1

-1.7 × 104 exp k2,MEAkMEA 6 /(m kmol-2 s-1) ) k-1

k2,MEAkTEA 6 /(m kmol-2 s-1) ) k-1

] ]

[

1.643 × 104 exp -

]

366.4 (36) (T/K)

As shown in Figures 3 and 4, kapp for the reaction of CO2 with MEA + TEA + H2O can be satisfactorily represented by eq 15. The larger deviations for kapp occur at MEA (0.3 M) + TEA (1.0 M) + H2O and MEA (0.5 M) + TEA (1.0 M) + H2O at 40 °C. For the absorption of CO2 in MEA + TEA + H2O to fall into the fast pseudo-first-order region, eq 19 has to be satisfied. The calculations of the enhancement factor in the instantaneous reaction region are presented in Tables 8 and 9. The diffusivities of MEA and TEA in MEA + TEA + H2O are estimated on the basis of the correlations of Snijder et al.40 and Hikita et al.,41 respectively. For all conditions, the Hatta numbers are all greater than 3, as in Tables 6 and 7. Also, the enhancement factors for instantaneous reaction are an order higher than the values of the Hatta number, as in Tables 8 and 9. Thus, the condition for eq 19 is satisfied, and eq 22 is valid for interpretation of the specific rate of mass transfer of CO2 in MEA + TEA + H2O. Conclusion

[

]

3569.8 1.566 × 10 exp (35) (T/K) 9

Kinetics of the absorption of CO2 into MEA + TEA + H2O were investigated at 30, 35, and 40 °C using a laboratory wetted wall column. Ten systems of which

Ind. Eng. Chem. Res., Vol. 41, No. 2, 2002 265

0.5 and 1.0 kmol m-3 TEA mixed with various MEA concentrations (0.1, 0.2, 0.3, 0.4, and 0.5 kmol m-3) were studied. Densities, viscosities, solubilities, and diffusivities of the blended-amine systems were also measured. The N2O analogy was applied to estimate the solubilities and diffusivities of CO2 in amine systems. The addition of small amounts of MEA to TEA results in a significant enhancement of CO2 absorption rates. On the basis of the fast pseudo-first-order rate for the CO2 absorption, the overall pseudo-first-order reaction rate constants were determined from kinetic measurements. A hybrid reaction rate model consisting of a zwitterion mechanism for MEA and a first-order reaction mechanism for TEA was used to represent the kinetics data. The overall average absolute percentage deviation for the calculation of the apparent reaction rate constant by the model is about 6.2%. This reaction rate model is found to be satisfactory for representing the CO2 absorption rate into MEA + TEA + H2O systems. Acknowledgment This research was supported by Grant NSC 89-2214E-033-011 of the National Science Council of the Republic of China. Notation d ) outside diameter of the wetted wall column, m D ) diffusivity of a gas in liquid, m2 s-1 E ) enhancement factor E∞ ) value of enhancement factor in instantaneous reaction region g ) gravitational constant h ) height of the wetted wall column, m H ) Henry law constant, kPa m3 kmol-1 Ha ) Hatta number, defined in eq 20 k ) reaction rate constant k-1 ) rate constant for the reverse reaction in eq 6 k2 ) second-order reaction rate constant in eq 6 k2,MEA ) rate constant in eq 15 kMEA ) rate constant in eq 9 k2,TEA ) rate constant in eq 15 kTEA ) rate constant in eq 9 kapp ) apparent reaction rate constant, s-1 kb ) second-order reaction rate constant for base B, in eq 7 kH2O ) reaction rate constant in eq 8 kL ) liquid-phase mass-transfer coefficient, in eq 20 k*OH- ) reaction rate constant for CO2 hydration, eq 3 kOH- ) reaction rate constant in eq 8 kov ) overall pseudo-first-order reaction rate constant, s-1 Kp,TEA ) TEA protonation constant ()[RR′N][H+]/[RR′NH+]), kmol m-6 Kw ) dissociation constant for water ()[H+][OH-]), kmol2 m-6 L ) liquid flow rate, m3 s-1 NA ) specific absorption rate, kmol m-2 s-1 pA ) partial pressure of CO2, kPa r ) reaction rate t ) temperature, °C tc ) contact time, s T ) temperature, K Greek Symbols R ) loading of CO2 in amine, kmol of CO2/kmol of amine η ) viscosity, mPa s

F ) density, g cm-3 φ ) absorption rate, kmol s-1

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Received for review August 7, 2001 Revised manuscript received October 18, 2001 Accepted October 18, 2001 IE010671L

Kinetics of Absorption of Carbon Dioxide into Aqueous Solutions of

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