Absorption of Carbon Dioxide into Aqueous Blends ofDiethanolamine and Methyldiethanolamine

Edward B. Rinker, Sami S. Ashour, and Orville C. Sandall*

Department of Chemical Engineering, University of California, Santa Barbara, California 93106-5080

In this work, a comprehensive model is developed for the absorption of carbon dioxide intoaqueous mixtures of primary or secondary alkanolamines with tertiary alkanolamines. Themodel, which is based on penetration theory, incorporates an extensive set of important reversiblereactions and takes into account the coupling between chemical equilibrium, mass transfer,and chemical kinetics. The reaction between CO2 and the primary or secondary amine is modeledaccording to the zwitterion mechanism. The key physicochemical properties that are needed forthe model are the CO2 physical solubility, the CO2 and amine diffusion coefficients, and thereaction rate coefficients and equilibrium constants. Data for carbon dioxide absorption intoaqueous mixtures of diethanolamine and methyldiethanolamine are compared to modelpredictions.

Introduction

Removal of acid gas impurities, such as CO2, isimportant in natural gas processing. Natural gas,depending on its source, can have varying concentra-tions of CO2. Some of the CO2 is often removed fromnatural gas because, when present at high levels, itreduces the heating value of the gas, and because it iscostly to pump this extra volume when it does not haveany heating value.

Aqueous solutions of alkanolamines are the mostwidely used solvents for removing CO2. The most com-monly used alkanolamines are the primary aminemonoethanolamine (MEA), the secondary amine dietha-nolamine (DEA), and the tertiary amine methyldietha-nolamine (MDEA). Primary and secondary amines reactrapidly with CO2 to form carbamates. By the additionof a primary or secondary amine to a purely physicalsolvent such as water, the CO2 absorption capacity andrate is enhanced manyfold. However, because there isa relatively high heat of absorption associated with theformation of carbamate ions, the cost of regeneratingprimary and secondary amines is high. Primary andsecondary amines also have the disadvantage of requir-ing 2 mol of amine to react with 1 mol of CO2; thus,their loadings are limited to 0.5 mol of CO2/mol ofamine. Tertiary amines lack the N-H bond required toform the carbamate ion and therefore do not reactdirectly with CO2. However, in aqueous solutions,tertiary amines promote the hydrolysis of CO2 to formbicarbonate and the protonated amine. Amine-promotedhydrolysis reactions are much slower than the directreactions of primary and secondary amines with CO2,and therefore, the kinetic selectivity of tertiary aminestoward CO2 is poor. However, the heat of reactionassociated with the formation of bicarbonate ions ismuch lower than that associated with carbamate forma-tion, and thus, the regeneration costs are lower fortertiary amines than for primary and secondary amines.Another advantage with tertiary amines is that thestoichiometry is 1:1, which allows for very high equi-librium CO2 loadings.

Chakravarty et al.1 suggested that, by mixing aprimary or secondary amine with a tertiary amine, theCO2 selectivity in the presence of H2S could be improvedand regeneration costs minimized. These blended aminesolutions also offer the advantage of setting the selectiv-ity of the solvent toward CO2 by judiciously mixing theamines in varying proportions, which results in anadditional degree of freedom for achieving the desiredseparation for a given gas mixture. This approach coulddramatically reduce capital and operating costs whileproviding more flexibility in achieving specific purityrequirements. Because of the need to exploit poorerquality crude and natural gas coupled with increasinglystrict environmental regulations, highly economical andselective acid gas treatment is more important todaythan at any time in the past. As a result, there has beena resurgence of interest in improved alkanolamine sol-vents and particularly in aqueous blends of alkanol-amines.

Design methods for acid-gas-treating processes em-ploying aqueous blends of alkanolamines vary widelyin their effectiveness at predicting process performance.Many acid-gas-treating processes are still designed byexperience and heuristics, resulting in overdesign,excessive energy consumption, and often failure to meetpurity requirements entirely (Chakravarty et al.1).Another common method uses equilibrium stage modelsmodified by tray efficiencies. This method, however,requires the use of existing plant data and lumps alluncertainties about the finite reaction rates of the gasesin the solvent into one parameter, the tray efficiency.Such a model cannot be predictive and will not capturethe essential interplay of mass transfer, chemical kinet-ics, and chemical thermodynamics that occur in complexchemical solvents such as aqueous blends of alkanola-mines. The third method of design is to develop modelsbased on the chemistry and physics of the process, whichaccounts for rates of mass transfer coupled with chemi-cal kinetics and thermodynamics. These models, whilestill requiring some experimental hydrodynamic infor-mation specific to different types of contacting devices,are capable of predicting column performance, thusminimizing the costs of design, equipment, and energyconsumption.* Corresponding author: sandall@engineering.ucsb.edu.

4346 Ind. Eng. Chem. Res. 2000, 39, 4346-4356

10.1021/ie990850r CCC: $19.00 © 2000 American Chemical SocietyPublished on Web 09/27/2000

The objective of the work presented here is to developa comprehensive model for the absorption of CO2 intoaqueous blends of tertiary and primary or secondaryamines. Experiments to test the model using methyldi-ethanolamine (MDEA) as the tertiary amine and di-ethanolamine (DEA) as the secondary amine are de-scribed.

Reaction Mechanism

When CO2 is absorbed into an aqueous solution of atertiary alkanolamine, R(1)R(2)R(3)N, and a primary or asecondary alkanolamine, R(4)R(5)NH, the following reac-tions may occur:

For MDEA, we have R(1) ) -CH3 and R(2) ) R(3) )-CH2CH2OH, and for DEA, we have R(4) ) R(5) -CH2-CH2OH. Ki, ki, and k-i are the equilibrium constant, theforward rate coefficient, and the reverse rate coefficientfor reaction i. Reactions 1-11 are considered to bereversible with finite reaction rates. Whereas reactions12-16 are considered to be reversible and instantaneouswith respect to mass transfer and at equilibrium, sincethey involve only proton transfers.

Note that not all of the reaction equilibrium constantsare independent. Only eight equilibrium constants (K2,K4, K5, K12, K13, K14, K15, and K16) are independent. Theremaining eight can be obtained by appropriate combi-nations of the independent equilibrium constants.

The interaction between the protonated and unpro-tonated amines according to the reaction

involves only a proton transfer and is considered to beinstantaneous and at equilibrium. Reaction 17 is im-plicitly included in the reaction scheme above, as it canbe obtained by properly combining the instantaneousequilibria reactions 12 and 13. Hence, we have K17 )K13/K12.

The proposed mechanism for the reaction betweenCO2 and tertiary alkanolamines, R(1)R(2)R(3)N, indicatesthat they do not react directly with CO2. Instead,tertiary alkanolamines act as bases that catalyze thehydration of CO2 (Donaldson and Nguyen,2 Haimour etal.,3 Versteeg and van Swaaij,4 Littel et al.,5 Rinker etal.6). In contrast, the proposed mechanism for thereaction between CO2 and a primary or secondaryalkanolamine, R(4)R(5)NH, involves the formation of azwitterion, R(4)R(5)NH + COO- (see reaction 4), followedby the deprotonation of the zwitterion by a base toproduce carbamate, R(4)R(5)NCOO-, and protonated base(see reactions 5-10) (Caplow,7 Danckwerts,8 Blauwhoffet al.,9 Versteeg and van Swaaij,10 Versteeg andOyevaar,11 Versteeg et al.,12 Glasscock et al.,13 Littel etal.,14 Rinker et al.15). Any base present in the solutionmight contribute to the deprotonation of the zwitterion.The contribution of each base would depend on itsconcentration as well as its strength. Hence, the maincontribution to the deprotonation of the zwitterion inan aqueous solution of a mixture of a primary orsecondary alkanolamine, R(4)R(5)NH, and a tertiaryalkanolamine, R(1)R(2)R(3)N, would come from R(4)R(5)NH,R(1)R(2)R(3)N, and to a lesser extent OH- and H2O. Thereare two limiting cases in the zwitterion mechanism.When the zwitterion formation reaction is rate-limiting,the reaction rate appears to be first-order in both theamine and CO2 concentrations. In the case of mono-ethanolamine (MEA), a primary alkanolamine, theformation of the zwitterion has been shown to be therate-determining step (Danckwerts,8 Sada et al.,16 Ver-steeg and van Swaaij,10 Littel et al.14). On the otherhand, when the zwitterion deprotonation reactions are

CO2 + H2O 798K1,k1

H2CO3 (1)

CO2 + OH- 798K2,k2

HCO- (2)

CO2 + R(1)R(2)R(3)N + H2O 798K3,k3

R(1)R(2)R(3)NH+ + HCO3- (3)

CO2 + R(4)R(5)NH 798K4,k4,k-4

R(4)R(5)NH+COO- (4)

R(4)R(5)NH+COO- + R(4)R(5)NH 798K5,k5,k-5

R(4)R(5)NH2+ + R(4)R(5)NCOO- (5)

R(4)R(5)NH+COO- + R(1)R(2)R(3)N 798K6,k6,k-6

R(1)R(2)R(3)NH+ + R(4)R(5)NCOO- (6)

R(4)R(5)NH+COO- + H2O 798K7,k7,k-7

H3O + R(4)R(5)NCOO- (7)

R(4)R(5)NH+COO- + OH- 798K8,k8,k-8

H2O + R(4)R(5)NCOO- (8)

R(4)R(5)NH+COO- + HCO3- 798

K9,k9,k-9

H2CO3 + R(4)R(5)NCOO- (9)

R(4)R(5)NH+COO- + CO32- 798

K10,k10,k-10

HCO3- + R(4)R(5)NCOO- (10)

R(4)R(5)NCOO- + H2O 798K11,k11

R(4)R(5)NH + HCO3- (11)

R(1)R(2)R(3)NH+ + OH- 798K12

R(1)R(2)R(3)N + H2O (12)

R(4)R(5)NHH2+ + OH- 798

K13R(4)R(5)NH + H2O (13)

HCO3- + OH- 798

K14CO3

2- + H2O (14)

HCO3- + H3O

+ 798K15

H2CO3 + H2O (15)

2H2O 798K16

OH- + H3O+ (16)

R(4)R(5)NH2+ + R(1)R(2)R(3)N 798

K17

R(4)R(5)NH + R(1)R(2)R(3)NH+ (17)

Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000 4347

rate-limiting, the overall reaction rate appears to be oforder 2 in the amine concentration. For the secondaryalkanolamines diethanolamine (DEA) and diisopropanol-amine (DIPA), the rate-determining step is the depro-tonation of the zwitterion (Littel et al.13). Severalauthors have reported some rate coefficients for thislimiting case of the zwitterion mechanism for DEA andDIPA for a few temperatures (Blauwhoff et al.,9 Ver-steeg and van Swaaij,16 Versteeg and Oyevaar,10 Glass-cock et al.,12 Rinker et al.14). Similarly, if neitherreaction in the zwitterion mechanism is rate-limiting,the reaction rate exhibits a fractional order between 1and 2 with respect to the amine concentration. However,the rate expression is more complicated than the limit-ing cases. Fractional orders are usually only observedfor reactions between CO2 and secondary amines (Sadaet al.,16 Versteeg and Oyevaar,11 Littel et al.14); howeverfractional orders have also been observed for stericallyhindered primary amines such as 2-amino-2-methyl-propanol (AMP) (Bosch et al.,17 Alper18).

Mathematical Model

For convenience, the concentrations of the chemicalspecies are renamed as follows:

Liquid Bulk Concentrations of All ChemicalSpecies. The liquid bulk concentrations of all chemicalspecies can be estimated from the initial concentrationsof R(1)R(2)R(3)N and R(4)R(5)NH; the initial CO2 loadingof the solution, L1; and the assumption that all reactionsare at equilibrium. Because the concentration of wateris much larger than the concentrations of all otherchemical species, changes to its concentration over veryshort contact times are negligible, and we assume thatits concentration remains constant. Hence, we only needto solve for the concentrations of the remaining 11chemical species. We have the following equations forthe liquid bulk concentrations u1

0, ... , u110 :

We have 11 unknowns (u10, ... , u11

0 ) and 11 nonlinearalgebraic equations that we can solve for the liquidbulk concentrations. We have found that Newton’smethod did not converge unless the initial guessesfor the liquid bulk concentrations were very close tothe (unknown) solution. We, therefore, used theNewton homotopy continuation method (Hanna andSandall19), which exhibited better convergence be-havior.

The Partial Differential and Nonlinear Alge-braic Equations That Describe the Diffusion/Reaction Processes. Higbie’s penetration model (Hig-bie,20 Danckwerts21) was used to set up the diffusion/reaction partial differential equations that describe theabsorption/desorption of CO2 into/from aqueous solu-tions of tertiary amines, R(1)R(2)R(3)N, and primary orsecondary amines, R(4)R(5)NH, in a laminar-liquid-jetabsorber or a stirred-cell absorber. All reactions weretreated as reversible reactions. The first 11 reactionshave finite reaction rates, which are given by thefollowing reaction rate expressions, where Ri is thereaction rate expression for reaction i:

where

u1 ) [CO2] u2 ) [R(1)R(2)R(3)N]

u3 ) [R(1)R(2)R(3)NH+] u4 ) [HCO3-]

u5 ) [OH-] u6 ) [CO32-]

u7 ) [H3O+] u8 ) [R(4)R(5)NH]

u9 ) [R(4)R(5)NH2+] u10 ) [R(4)R(5)NCOO-]

u11 ) [H2CO3] u12 ) [H2O]

Overall tertiary amine, R(1)R(2)R(3)N, balance

u20 + u3

0 )[R(1)R(2)R(3)N]initial (18)

Overall primary or secondary amine, R(4)R(5)NH,balance

u80 + u9

0 + u100 )[R(4)R(5)NH]initial (19)

Overall carbon (from CO2) balance

u10 + u4

0 + u60 + u10

0 + u110 + L1{[R(1)R(2)R(3)N]initial +

[R(4)R(5)NH]initial} (20)

Electroneutrality balance

u30 + u7

0 + u90 - u4

0 - u50 - 2u6

0 - u100 ) 0 (21)

All reactions at equilibrium(only independent equilibrium constants)

K2 )u4

0

u10 u5

0(22)

K4K5K13K16 )u7

0 u100

u10 u8

0(23)

K12 )u2

0

u30 u5

0(24)

K13 )u8

0

u90 u5

0(25)

K14 )u6

0

u40 u5

0(26)

K15 )u11

0

u40 u7

0(27)

K16 ) u50 u7

0 (28)

R1 ) -k1u1 +k1

K1u11 (29)

R2 ) -k2u1u5 +k2

K2u4 (30)

R3 ) -k3u1u5 +k3

K3u3u4 (31)

R4,...,10 )-k4[u1u8 - (AB)u10]

1 + 1B

(32)

4348 Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000

and

Note that eq 32 was derived with the assumption ofa pseudo-steady-state approximation for the zwitterionreaction intermediate, R(4)R(5)NH+COO-. The partialdifferential equations that describe the diffusion/reac-tion processes were combined so as to eliminate the verylarge reaction rates for the instantaneous reactions 12-16. Because these reactions are assumed to be atequilibrium, their equilibrium constant expressionswere used to complete the equations that are requiredto solve for the concentration profiles of all chemicalspecies. Furthermore, the diffusion coefficients of theionic species were assumed to be equal. With thisassumption, the electrostatic potential gradient termsin the diffusion/reaction partial differential equationsfor the ionic species can be neglected, while the elec-troneutrality of the solution is preserved. The morerigorous approach of taking into account the electro-static potential gradient terms with unequal diffusioncoefficients for the ionic species requires much greatercomputational effort with essentially intangible effectson the predicted rates of absorption. The followingequations describe the diffusion/reaction processes:

We have 11 partial differential/algebraic equations thatwe can solve for the concentrations of the 11 chemicalspecies, u1, ... , u11.

Initial Condition and Boundary Condition at x ) ∞.At t ) 0 (for all x ) 0) and at x ) 8 (for all t ) 0), theconcentrations of all chemical species are equal to theirliquid bulk concentrations.

Boundary Condition at Gas-Liquid Interface (x ) 0).At x ) 0 (gas-liquid interface), the fluxes of thenonvolatile chemical species are equal to zero, whichleads to the following equations:

for all i except i ) 1 (CO2). For the volatile component,CO2, the mass transfer rate in the gas near the interfaceis equal to the mass transfer rate in the liquid near theinterface.

H1 is the physical equilibrium constant (Henry’s lawconstant) of CO2, which is defined as the interfacialpartial pressure of CO2 in the gas phase, P1

/, divided bythe interfacial concentration of CO2 in the liquid, u1

/.For the case of pure CO2 in the gas phase, the interfacialpartial pressure of CO2, P1

/, is the same as the bulkpartial pressure of CO2, P1, and there is not any masstransfer resistance in the gas-phase (kg,1 f ∞). Hence,

A )( k5

k-4) u9

K4K5+ ( k6

k-4) u3

K4K6+ ( k7

k-4) u7

K4K7+

( k8

k-4) u12

K4K8+( k9

k-4) u11

K4K9+(k10

k-4) u4

K4K10(33)

B ) ( k5

k-4)u8 + ( k6

k-4)u2 + ( k7

k-4)u12 + ( k8

k-4)u5 +

( k9

k-4)u4 + (k10

k-4)u6 (34)

R11 ) -k11u10 +k11

K11u4u8 (35)

CO2 balance

∂u1

∂t) D1

∂2u1

∂x2+ R1 + R2 + R3 + R4,...10 (36)

Total carbon (from CO2) balance

∂u1

∂t∂u4

∂t+

∂u6

∂t+

∂u10

∂t+

∂u11

∂t) D1

∂2u1

∂x2+ D4

∂2u4

∂x2+

D6

∂2u6

∂x2+ D10

∂2u10

∂x2+ D11

∂2u11

∂x2(37)

Total tertiary amine, R(1)R(2)R(3)N, balance

∂u2

∂t+

∂u3

∂t) D2

∂2u2

∂x2+ D3

∂2u3

∂x2(38)

Total primary or secondary amine, R(4)R(5)N,balance

∂u8

∂t+

∂u9

∂t+

∂u10

∂t) D8

∂2u8

∂x2+ D9

∂2u9

∂x2+ D10

∂2u10

∂x2(39)

Electroneutrality balance

∂u3

∂t+

∂u7

∂t+

∂u9

∂t-

∂u4

∂t-

∂u5

∂t- 2

∂u6

∂t-

∂u10

∂t)

D3

∂2u3

∂x2+ D7

∂2u7

∂x2+ D9

∂2u9

∂x2- D4

∂2u4

∂x2- D5

∂2u5

∂x2-

2D6

∂2u6

∂x2- D10

∂2u10

∂x2(40)

Carbamate, R(4)R(5)NCOO-, balance

∂u10∂t

) D10

∂2u10

∂x2+ R11 - R4,...10 (41)

Instantaneous reactions assumed to be at equilibrium

K12 )u2

u3u5(42)

K13 )u8

u9u5(43)

K14 )u6

u4u5(44)

K15 )u11

u4u7(45)

K16 ) u5u7 (46)

ui ) ui0, i ) 1, ... , 11 (47)

∂ui

∂x) 0 at x ) 0, t > 0 (48)

-D1

∂u1

∂x) kg,1[P1 - H1u1(0,t)] at x ) 0, t > 0 (49)

Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000 4349

the boundary condition for CO2 at the gas-liquidinterface reduces to

The differential equations are integrated from t ) 0to t ) τ, the contact time. For the laminar-liquid-jetabsorber, the contact time is given by

where Q is the volumetric liquid flow rate, and d and lare the diameter and length of the laminar-liquid jet,respectively. For Higbie’s penetration model, the liquid-phase mass transfer coefficient for physical absorptionof CO2 is defined as

The average rate of absorption of CO2 per unit interfa-cial area is then computed from the following equation:

The enhancement factor of CO2 is determined accord-ing to the following equation:

where u1/ and u1

0 are the interfacial and bulk concen-trations of CO2 in the liquid, respectively.

Method of Solution for the Partial Differential/Algebraic Equations. The method of lines was usedto transform each partial differential equation into asystem of ordinary differential equations in t (Hannaand Sandall19). The systems of partial differential/algebraic equations were transformed into larger sys-tems of ordinary differential/algebraic equations, whichwere then solved by using the code DDASSL (Petzold22)in double-precision Fortran on an HP-735 computer.

Model Parameters

The nitrous oxide analogy method was used to esti-mate the CO2 solubility (Rinker and Sandall23) and theCO2 diffusivity (Ashour et al.24) in the aqueous aminesolutions. The kinetics for CO2/MDEA were measuredby Rinker et al.,6 and the kinetics for the CO2/DEAreaction were determined by Rinker et al.15 The diffu-sion coefficients of MDEA and DEA were estimated fromthe diffusivity data of Hikita et al.25

We include here correlations that were used toestimate various other reaction rate coefficients andequilibrium constants that were obtained from theliterature. Values for the forward rate coefficient ofreactions 1 and 2, k1 and k2, respectively, were calcu-lated from the following correlations, which were re-ported by Pinsent et al.26 for the temperature ranges of273-311 K and 273-313 K, respectively:

Table 1 gives the values of the equilibrium constantsused in the model calculations, with the exception ofK15. The value of K15 at 298 K was taken to be 2 × 10-4

m3/kmol (Cotton and Wilkinson31). K15 was then cor-rected for temperature dependence according to thefollowing equation:

where the standard enthalpy change of reaction, ∆H0,is assumed to be independent of T and is approximatedby its value at 298.15 K, ∆H298.15

0 (Smith and VanNess32). ∆H298.15

0 values were calculated from valuesreported in the CRC Handbook of Chemistry andPhysics (Lide33).

Experimental Apparatus and Procedure

The rates of absorption of CO2 into aqueous solutionsof DEA and MDEA were measured in a laminar-liquid-jet absorber and a stirred-cell absorber. A schematicdrawing of the laminar-liquid-jet absorber is shown inFigure 1. The laminar-liquid-jet absorber and its opera-tion are described in detail by Rinker et al.15 The stirred-cell absorber is shown schematically in Figure 2. Theabsorption chamber is made of a 30.5 cm long, 10.1 cmi.d. Pyrex cylinder and is enclosed in a constant-temperature heating jacket constructed from a 31 cmlong, 24 cm i.d. Lucite cylinder. The ends of the Lucitecylinder are sealed with rubber O-rings between twoanodized aluminum flanges, and the glass cylinder issandwiched between two stainless steel flanges with theends sealed by Teflon gaskets. Cooling or heating wateris supplied to the jacket and recycled to a constant-temperature circulating bath. Two separate stainlesssteel coils are placed in the heating/cooling jacket andare used to control the temperatures of the liquid andgas feeds. The liquid supply is introduced into thechamber by a 0.635 cm o.d. stainless steel tube that canslide in the vertical direction through the bottom flangeand can be locked in position by a Swagelok nut withTeflon ferrules. The end of this tube is plugged, and theliquid is discharged into the absorption chamber fromperforations on the side of the tube near the pluggedend. This assembly makes it possible to discharge theliquid at different heights if so desired. The absorptionchamber is also equipped with four flat stainless steelbaffles, which help in reducing vortex formation andpromoting better mixing of the liquid phase. The bafflesare 12.7 cm long, 1.0 cm wide, and 0.1 cm thick and arepinned to the bottom flange, and their top ends areconnected by a wire ring. The distance between thebaffles and the glass wall of the chamber is about 0.35cm. When the liquid height in the chamber is 11.0 cm,the baffles extend about 1.7 cm above the liquid surface.

The absorber has two concentric shafts that protrudeinto the chamber through the top flange. The inner shaftis 0.6 cm in diameter and 36.5 cm long and is made ofstainless steel. This shaft extends to a Teflon bushingin the bottom flange and is supported at the top end bya pin bearing held in a cup on a crossbar. There are

u1(0,t) ) u1/ )

P1

H1at x ) 0, t > 0 (50)

τ ) πd2l4Q

(51)

kl,10 ) 2xD1

πτ(52)

RA1 ) -D1

τ ∫0

τ ∂u1

∂x(0,t) dt (53)

E1 )RA1

kl,10 (u1

/- u10)

(54)

log10(k1) ) 329.85 - 110.54 log10(T) - 17265.4T

(55)

log10(k2) ) 13.635 - 2895T

(56)

ln[ K15(T)

K15(298.15)] ) -∆H0

R (1T

- 1298.15) (57)

4350 Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000

two liquid-phase impellers mounted on the shaft, onewith a diameter of 7.5 cm mounted at a height of 8.0cm from the bottom flange, and another with a diameterof 6.0 cm mounted at a height of 3.0 cm from the bottomflange. The second shaft, which is a hollow stainlesssteel tube 20.0 cm long with a 1.1 cm i.d., is supportedby a concentric tubular extension welded onto the topflange and extends 10.0 cm into the chamber. Mountedon the second shaft are two impellers for stirring thegas phase, a 7.5 cm diameter impeller at 10.5 cm fromthe top flange, and a 9.5 cm diameter impeller at 17.5cm from the top flange. When the liquid height in thechamber is 11.0 cm, the bottom edge of the second gasimpeller is 2.0 cm away from the gas-liquid interface.Two sets of ball bearings, Teflon packings, and spring-loaded Teflon seals support the tubular shaft. Twosimilar assemblies support the liquid shaft. The twoshafts are driven independently by two variable-speedmotors.

The liquid feed is pumped to a surge tank and thenthrough a calibrated rotameter and through a stainlesssteel coil in the constant-temperature jacket. It is then

introduced into the absorption chamber through thebottom flange and is discharged into the chamber byseven perforations about 8 cm from the bottom of thechamber. The liquid leaves the chamber through thebottom flange and goes to a liquid-leveling reservoir.The leveling device is similar to that used in thelaminar-liquid-jet absorber.

The gases pass through two sets of regulators (highand low pressure) and then through mass flow control-lers in order to maintain constant gas feed rates. Thegases subsequently pass through soap bubble meters sothat their volumetric flowrates can be measured. Thegases are next mixed in a T-fitting and fed to asaturator. After the saturator, the overall gas flow rateis measured with a bubble meter. The temperature ofthe gas mixture is measured and recorded at this pointwith a type-J thermocouple. The gas mixture thenpasses through the stainless steel heating/cooling coilin the constant-temperature jacket before being intro-duced into the absorption chamber through the tube inthe top flange. The gas mixture is sampled just beforeit enters the absorption chamber and is analyzed with

Figure 1. Schematic drawing of the laminar-liquid-jet absorber.

Table 1. Equilibrium Constant Correlations Used for Model Calculations

equilibriumconstant equation

temp range(K) reference

K16 log10(K16) ) 8909.483 - 142613.6/T 293-573 Olofsson-4229195 log10(T) + 9.7384 T and Hepler27

-0.012 963 8T2 + (1.150 68 × 10-5)T3

-(4.602 × 10-9)T4

K2K16 log10(K2K16) ) 179.648 + 0.019 244T 273-523 Read28

-67.341 log10(T) - 7 495.441/TK14K16 log10(K14K16) ) 6.498 - 0.023 8T - 2 902.4/T 273-323 Danckwerts

and Sharma29

K4K5K13K16 log10(K4K5K13K16) ) -10.549 2 + 1 526.27/T 298-333 Barth et al.30

K13K16 log10(K13K16) ) -4.030 2 - 1 830.15/T + 0.0043T 298-333 Barth et al.30

K12K16 log (K12K16) ) -14.01 + 0.018T 298-333 Barth et al.30

Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000 4351

a gas chromatograph. The gas is discharged just abovethe larger impeller through perforations near the endof the feed tube. The gas exits the chamber through afitting in the top flange and is sampled for compositionby a gas chromatograph. The volumetric flow rate of theexit gas is measured with a soap bubble meter and thetemperature of the exit gas is measured with a type-Jthermocouple.

The gas samples were analyzed using a model 5890Hewlett-Packard gas chromatograph equipped with athermal conductivity detector. The column used in theGC was a 6 ft. long, 0.085 in. i.d., stainless steel columnpacked with Poropak Q with a mesh size of 80/100. Therun conditions were 70 °C for 1.5 min, ramp to 90 °C at5 °C per minute, cool, and equilibrate for 2 min.

Experimental Results

(A) Absorption Measurements at Short ContactTimes Using a Laminar-Liquid-Jet Absorber. Ac-cording to the zwitterion mechanism, any base cancontribute to the deprotonation of the zwitterion. Forthe aqueous DEA case, the potential bases are R2NH,OH-, H2O, HCO, and CO. In an earlier study (Rinkeret al.15), the contributions of OH-, H2O, HCO, and COto the deprotonation of the zwitterion were found to beinsignificant for the reaction of CO2 with aqueous DEAin the laminar-liquid-jet absorber. However, if CO2 isabsorbed into aqueous blends of DEA and MDEA, it ispossible that MDEA could contribute significantly to the

deprotonation of the zwitterion. Littel et al.14 report thatMDEA makes a measurable contribution to the depro-tonation of the zwitterion for CO2 absorption intoaqueous blends of DEA and MDEA.

To test the hypothesis that MDEA contributes to thedeprotonation of the zwitterion, the rates of CO2 absorp-tion into aqueous blends of DEA and MDEA weremeasured in the laminar-liquid-jet absorber at 25 °C.The total weight percents of the aqueous DEA/MDEAblends were 10, 30, and 50 wt %, and the molar ratiosof DEA to MDEA were varied from zero to infinity. Theabsorption data are listed in Tables 2-4. The rates ofCO2 absorption were predicted using the numericalmodel developed in this work without taking intoaccount the contribution of MDEA to the deprotonationof the zwitterion. The predicted and measured rates ofCO2 absorption are compared in the parity plot shown

Figure 2. Schematic drawing of stirred-cell absorber.

Table 2. CO2 Absorption Data in 10 wt % Blends of DEAand MDEA Obtained in the Laminar-Liquid-Jet Absorberat 25 °C

[MDEA](kmol/m3)

[DEA](kmol/m3)

RCO2 × 105

measured(kmol/m2s)

RCO2 × 105

predicted(kmol/m2s)

0.844 0 1.22 1.220.808 0.040 1.29 1.350.692 0.173 1.45 1.440.586 0.293 1.65 1.650.532 0.355 1.71 1.760.232 0.696 2.40 2.690 0.959 3.28 3.32

4352 Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000

in Figure 3. From this plot, it is clear that there is goodagreement between the measured rates of CO2 absorp-tion and the rates of CO2 absorption predicted by themodel neglecting the contribution of MDEA to thedeprotonation of the zwitterion. The average deviationof the predicted rates from the measured rates is 6.8%.On the basis of these experiments, it appears thatMDEA does not significantly contribute to the depro-tonation of the zwitterion.

A possible explanation for the difference in the resultsof this study and the work of Littel et al.14 is that weused a laminar-liquid-jet absorber whereas they useda batch stirred-tank reactor. The laminar-jet absorberoperates at steady state, with a gas-liquid contact timeof about 0.005 s, whereas the reactor used by Littel etal. operates under transient conditions and over muchlonger gas-liquid contact times. For very short contacttimes, the deprotonation reaction with MDEA may nothave time to contribute significantly to the overall rateof absorption.

(B) Absorption Measurements at Long ContactTimes Using a Stirred-Cell Absorber. To predictrates of CO2 absorption obtained using the stirred-cellapparatus, accurate values of the liquid-phase masstransfer coefficient for physical absorption must bedetermined from experimental absorption data obtainedfor a model system. The liquid-phase physical masstransfer coefficient was determined by measuring theconcentration of dissolved CO2 in the liquid effluentfrom the stirred-cell absorber using wet-chemical analy-sis. The liquids used for these experiments were puredeionized water and pure poly(ethylene glycol) (PEG400).The gas phase was pure CO2 saturated with the vaporpressure of the liquid. The liquid feedstocks weredegassed at elevated temperatures. The concentrationof CO2 dissolved in the liquid was determined bytitration of samples of effluent. The water samples weredrawn slowly into a syringe and injected into an equalvolume of 0.05 M NaOH to convert all of the dissolved

CO2 into carbonate. The carbonate was precipitatedwith barium chloride, and the remaining NaOH wastitrated with 0.05 M HCl. This method was reproducibleto within 1%, and it was accurate enough to measurethe small rates of CO2 absorption over liquid stirringspeeds of 30-140 rpm. For the pure PEG400 experi-ments, the samples were injected into twice theirvolume of 0.05 M NaOH in order to convert the CO2 tocarbonate and dilute the PEG400 so that the sampleswere mostly water (on a molar basis) for accurate pHdetermination. Pure PEG400 was used to check theviscosity dependence of the mass transfer coefficient asit has a viscosity of 0.9527 P at 25 °C and water has aviscosity of 0.00895 P at 25 °C.

The physical mass transfer coefficients were calcu-lated from the measured data using the followingequations:

where CA/ is the interfacial concentration of CO2, CA

0 isthe bulk or effluent concentration of CO2, L is thevolumetric liquid flowrate, As is the area of the gas-liquid interface (79.64 cm2), and k is the physical liquid-phase mass transfer coefficient. The mass transfercoefficients were correlated in dimensionless form asfollows:

where

The dimensionless mass transfer coefficients used toobtain the correlation given above are plotted inFigure 4.

(i) CO2 Absorption into Aqueous MDEA. Rates of CO2absorption into aqueous solutions of MDEA were mea-sured in the stirred-cell absorber at 25 °C. The CO2partial pressure was varied from 0.15 to 0.5 atm, andthe diluent was N2. The measured absorption rates are

Table 3. CO2 Absorption Data in 30 wt % Blends of DEAand MDEA Obtained in the Laminar-Liquid-Jet Absorberat 25 °C

[MDEA](kmol/m3)

[DEA](kmol/m3)

RCO2 × 105

measured(kmol/m2s)

RCO2 × 105

predicted(kmol/m2s)

2.581 0 1.12 0.8702.118 0.529 1.97 1.491.376 1.376 3.08 2.830.712 2.135 3.77 3.800 2.947 4.86 4.752.581 0 1.06 0.8502.118 0.529 1.73 1.511.376 1.376 3.06 2.820.712 2.135 3.72 3.800 2.947 4.65 4.71

Table 4. CO2 Absorption Data in 50 wt % Blends of DEAand MDEA Obtained in the Laminar-Liquid-Jet Absorberat 25 °C

[MDEA](kmol/m3)

[DEA](kmol/m3)

RCO2 × 105

measured(kmol/m2s)

RCO2 × 105

predicted(kmol/m2s)

4.217 0.202 0.753 0.6924.217 0.202 0.868 0.7833.722 0.772 1.10 1.113.719 0.771 1.15 1.123.722 0.772 1.09 1.093.034 1.517 1.67 1.723.034 1.517 1.50 1.74

Figure 3. Comparison of predicted and measured rates of CO2absorption into aqueous blends of DEA and MDEA at 25 °C inthe laminar-liquid-jet absorber. The total amine concentrationswere 10, 30, and 50 mass %.

RA1 ) k°l,1(CA/- CA

0) (58)

RA1 )CA

0 LAs

(59)

Sh ) 0.0193Re0.845Sc0.5 (60)

Sh )k°l,1da

DARe )

FNLdi2

µSc ) µ

FDA(61)

Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000 4353

compared to the model predictions in the parity plotshown in Figure 5. The average deviation of the modelpredictions from the measured rates is 24%.

(ii) CO2 Absorption into Aqueous DEA. Rates of CO2absorption into 10 wt % aqueous DEA were measuredin the stirred-cell absorber at 25 °C. The CO2 partialpressure was varied from 0.088 to 0.712 atm, and thediluent was N2. The measured and predicted rates ofCO2 absorption are compared in the parity plot shownin Figure 6. In this case, there is fairly good agreementbetween the model predictions and the measurements,

with an average deviation of 10.2% from the measuredrates.

(iii) Absorption of CO2 into Aqueous Blends of DEAand MDEA. Rates of absorption of CO2 into aqueousblends of DEA and MDEA were measured in the stirred-cell absorber at 25 °C. The total amine concentrationwas approximately 50 wt %, and the molar ratios ofDEA to MDEA were 0.050, 0.207, 0.250, and 0.500. Themodel predictions are compared to the measured ab-sorption rates in Figure 7. The average deviation of thepredictions from the measured rates of absorption is17.6%. The predicted rates of absorption were made byneglecting the contribution of MDEA to the deprotona-tion of the zwitterion. Some of the absorption rates areoverpredicted and some are underpredicted when theMDEA contribution is neglected. As a result, the CO2absorption data obtained in the stirred-cell absorbersupport the findings for CO2 absorption into blends inthe laminar-liquid-jet absorber that MDEA does notsignificantly contribute to the deprotonation of thezwitterion.

Conclusions

The model developed in this work for the rates ofabsorption of carbon dioxide into a aqueous mixed aminesolutions was found to agree reasonably well with theexperiments. The reaction between carbon dioxide andthe secondary amine, DEA, was described by the zwit-terion mechanism in this model. For our experimentswith gas-liquid contact times varying from approxi-mately 0.005 to 10 s, it appears that the tertiary amine,MDEA, does not contribute significantly to the depro-tation of the zwitterion. The only species that contrib-utes significantly to the deprotonation of the zwitterionis DEA, and thus, eqs 6-10 in the reaction model couldbe deleted according to the results obtained in thisstudy.

The key physicochemical property needed for themodel calculations is the physical solubility of CO2. Anyuncertainty in this property translates to an equivalenterror in the predicted absorption rate.

Acknowledgment

This work was sponsored by the Gas Research Insti-tute and the Gas Processors Association.

Figure 4. Liquid-phase physical absorption mass transfer coef-ficient correlation.

Figure 5. Comparison of predicted and measured rates of CO2absorption into aqueous MDEA at 25 °C in the stirred-cellabsorber.

Figure 6. Comparison of predicted and measured rates of CO2absorption into aqueous DEA at 25 °C in the stirred-cell absorber.

Figure 7. Comparison of predicted and measured rates of CO2absorption into aqueous 50 mass % blends of DEA and MDEA at25 °C in the stirred-cell absorber.

4354 Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000

Notation

As ) gas-liquid interface area, m2

CA ) diameter of the laminar-liquid jet, mDi ) coefficient of species i in the aqueous alkanolamine

solution, m2/sDEA ) diethanolamineDIPA ) diisopropanolamineE1 ) factor of CO2, defined by eq 54h0 ) spacing next to the gas-liquid interface for the

discretized spatial variable x, mH1 ) physical equilibrium constant (Henry’s law constant)

for CO2, H1 ) P1//u1

/, (kPa m3)/kmolkapp ) rate coefficient for the reaction between CO2 and a

secondary or a primary alkanolamine, defined by eq 70,m3/(kmol s)

kg,1 ) gas-phase mass transfer coefficient for CO2, kmol/(kPa m2 s)

kl,10 ) mass transfer coefficient for CO2 in the liquid phase,defined by equation (52), m/s

ki ) rate coefficient of reaction i, s-1 for first-orderreactions, m3/(kmol s) for second-order reactions

k-i ) rate coefficient of reaction iKi ) constant for reaction il ) length of the laminar-liquid jet, mL1 ) CO2 loading of the aqueous amine solution, (kmol

CO2)/(kmol amine)MDEA ) N-methyldiethanolamineMEA ) monoethanolamineNL ) impeller speed, rev/sP1 ) pressure of CO2 in the gas phase, kPaP1

/) partial pressure of CO2 in the gas phase, kPaQ ) volumetric flow rate of the liquid, m3/sRi ) reaction rate of reaction i, kmol/(m3 s)RA1 ) rate of absorption of CO2 per unit interfacial area,

defined by equation (53), kmol/(m2 s)Re ) Reynolds number defined by eq 80Sc ) Schmidt number defined by eq 80Sh ) Sherwood number defined by eq 80t ) independent time variable, sT ) absolute temperature, Kui ) concentration of species i in the liquid phase (which

is a function of x and t), kmol/m3

ui0 ) liquid bulk concentration of s pecies i (which is aconstant), kmol/m3

u1/ ) interfacial concentration of CO2 in the liquid, kmol/m3

x ) independent spatial variable, m

Greek letters

µ ) viscosity of the aqueous alkanolamine solution, kg/(ms)

F ) density of the aqueous alkanolamine solution, kg/m3

τ ) gas-liquid contact time, defined by eq 51 for a laminar-liquid-jet absorber, s

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Received for review November 22, 1999Revised manuscript received July 17, 2000

Accepted July 20, 2000

IE990850R

4356 Ind. Eng. Chem. Res., Vol. 39, No. 11, 2000