Ind. Eng. Chem. Res. 2007, 46, 233-241
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SEPARATIONS Competitive Absorption-Desorption of Acid Gas into Water-DEA Solutions Renaud Cadours* IFP-Lyon, BP3-69390 Vernaison, France
Damien Roquet TOTAL, Tour Coupole-La De´ fense 6, 92078 Paris La De´ fense, France
Gauthier Perdu PROSERNAT, Tour AREVA-La De´ fense 6, 92084 Paris La De´ fense, France
Measurements of the CO2 absorption rate into a diethanolamine (DEA) aqueous solution were made using a Lewis cell to determine the CO2-DEA kinetic parameters. The absorption rate was controlled by a gas-side mass-transfer phenomenon. This was achieved by monitoring the total pressure inside the Lewis cell. The gas-side mass-transfer coefficient was deduced from H2S absorption measurements in an aqueous solution of DEA. The kinetic rate constant of the reaction of formation of the zwitterion determined in this work is in good agreement with existing literature values. Measurements of the H2S absorption rate into a DEA aqueous solution previously loaded with CO2 showed a competition between H2S absorption and CO2 desorption. A mass-transfer rate model that takes the effects of chemical reactions into consideration enabled these opposite mass transfers to be modeled. Introduction Natural gas treatment involves several steps, resulting in increased costs and operating complexity. This treatment usually involves water and hydrocarbon dew-pointing and gas deacidification. Condensables are removed from natural gas, to prevent the formation of liquid phases and hydrates during gas transport. Acid gas removal is necessary to meet current pipeline specifications for hydrogen sulfide (H2S) and carbon dioxide (CO2) contents. Many processes are available to remove acid gases from sour gas mixtures. The most common solvent processes used for this operation are based on aqueous solutions of alkanolamines. Primary and secondary alkanolamines such as diethanolamine (DEA) are used for total deacidification. Tertiary amines, such as methyldiethanolamine (MDEA), can be used for selective H2S removal and to maximize CO2 slippage, to produce rich H2S gas for Claus treatment. Several parallel reversible reactions occur during the absorption of CO2 and H2S into an aqueous alkanolamine solution. The case of a reversible reaction is very complex, because of the fact that reaction rate expressions are not linear. Van Krevelen and Hoftijzer1 originally proposed approximate analytical solutions. Nevertheless, numerical solutions showed that these approximations generally were not valid.2 In the case of several reversible reactions, Onda et al.3-5 suggested some approximate solutions for specific cases. However, most of the recent work that was dedicated to the absorption of acid gases in aqueous alkanolamine solutions used a numerical technique to solve the equations; this numerical technique describes the * To whom correspondence should be addressed.
phenomenon of multicomponent mass transfer, coupled with parallel complex reversible chemical reactions. Cornelisse et al.6 studied the simultaneous absorption of CO2 and H2S in an aqueous solution of secondary alkanolamine, by means of the penetration theory. The resulting model was restricted to a few stoichiometric schemes. Bosch et al.7 described the simultaneous absorption of H2S and CO2 in aqueous solutions of alkanolamines, assuming reversibility, a generalized reaction rate, and a generalized stoichiometry for all reactions. The model was limited only by the type of reactions. Littel et al.8 used the same approach to describe the simultaneous absorption of H2S and CO2 in aqueous solutions of primary, secondary, or tertiary amines. They generalized the models of Bosch et al.8 by considering finite-rate reactions and instantaneous reactions, with respect to mass transfer. They modeled the latter as finite-rate reactions with very high reaction rate constants, as previously described by Glasscock and Rochelle.9 In this work, CO2 absorption rates into DEA aqueous solutions were measured in a thermoregulated reactor with a constant gas/liquid interfacial area. Mass-transfer resistance was controlled by monitoring the total pressure in the cell with nitrogen. Mass-transfer coefficients, for both gas and liquid phases, were deduced from specific chemical systems. The liquid-side mass-transfer resistance was obtained from N2O physical absorption into water-MDEA solvents. H2S absorption data into DEA aqueous solutions were used to determine the gas-side mass-transfer resistance. The absorption of CO2 or H2S into acid-gas-loaded solutions led to opposite mass transfers of H2S and CO2. A numerical model based on the film theory was used to explain the opposite mass transfers of H2S and CO2 observed in the Lewis cell.
10.1021/ie060019w CCC: $37.00 © 2007 American Chemical Society Published on Web 11/30/2006
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Figure 1. Experimental equipment for kinetic measurements.
Experimental Section The reaction between CO2 and DEA has been studied several times.10-15 The zwitterion mechanism that has been proposed by Caplow10 is the most widely accepted mechanism for primary and secondary alkanolamines. It involves the formation of an intermediate zwitterion that is deprotonated by the basic component to produce a carbamate and a protonated base.
CO2 + R2NH T R2NH+COOR2NH+COO- + base T (base)H + + R2NCOOThe contribution of each base is dependent on its concentration and its basicity. Depending on the mechanism-limiting step, which could be zwitterion formation or deprotonation, the reaction order in amine concentrations is observed to be between 1 and 2. Studies based on short gas-liquid contact time (for example, wetted-wall column12 or wetted-wall sphere13 or laminar jet apparatus15) led to a reaction order of 1 in amine concentrations. Reaction orders between 1 and 2 in DEA concentrations are usually obtained with studies using stirredcell reactors. The influence of the zwitterion-deprotonation step can be significant if long gas-liquid contact time measurements are made, especially at low alkanolamine concentrations. For example, Blauwhoff et al.16 and Littel et al.14 used the same methodology: a batch reactor, with respect to both gas and liquid phases. The absorption rate was deduced from the pressure decrease resulting from the CO2 injection. The overall reaction rate constant was defined when pseudofirst-order conditions were met. This procedure means that the CO2 absorption rate into a partially loaded solution is measured because of the absorption of a significant amount of the injected acid gas, according to the example given by Blauwhoff et al.16 These authors observed a significant influence of hydroxide ions on the overall reaction rate. Another difficulty in measuring the CO2 absorption rate into DEA aqueous solutions is the possible influence of a gas-side mass-transfer resistance. This resistance may appear, especially at low CO2 partial pressure. In this work, we propose an experimental approach using a Lewis cell to measure the CO2 absorption rate into an aqueous DEA solution, for temperatures ranging from 313 K to 333 K, and DEA concentrations ranging from 10 wt % to 30 wt %. Monitoring of the total pressure in the cell ensured the stability of gas-side mass-transfer resistance, which was characterized
by H2S absorption rate measurements into aqueous DEA solutions. The injection of very small amounts of CO2 in the cell made it possible to consider a pseudo-first-order approximation for determination of the apparent rate constant. This injection prevented acid gas from accumulating in the solvent and solvent composition from being modified, especially for hydroxide ion concentration. The overall apparent constant was compared to literature data. Experimental Setup. The experimental apparatus was developed to measure acid-gas absorption rates into wateralkanolamine solutions. The main equipment is the reactor (see Figure 1), which is composed of a (6.00 ( 0.02) × 10-2 m internal diameter quartz cylinder that has been closed at both ends by two stainless steel (SS304L) metallic flanges. The total volume available for gas and liquids is (346.5 ( 0.1) × 10-6 m3. The reactor is provided with a six-bladed Rushton turbine, (3.10 ( 0.02) × 10-2 m in diameter, in the liquid phase and a four-bladed impeller, (2.00 ( 0.02) × 10-2 m in diameter, in the gas phase. They are both magnetically driven by adjustablespeed motors. This technique prevents leaking, friction, and heat generation, which appear when using stems that pass through the top and bottom of the cell. Stirring speeds are checked periodically: they remain constant within 1 rpm during the tests. Four vertical baffles that are in place prevent vortices in the liquid. A horizontal plate and a ring, placed halfway up the cell, set both the level and the area of the interface available for gas/liquid transfer and ensure its stability during stirring. The gas-liquid interface area, (13.0 ( 0.1) × 10-4 m2, was geometrically estimated. The temperature in the reactor is known within (0.2 K through a 100Ω platinum probe, calibrated between 273 K and 403 K, against a Herau¨s TLH600 100Ω platinum probe that was calibrated by the Laboratoire National d’Essais. The absorption cell temperature is regulated by immersion of the cell in a thermoregulated bath. The absorption rate is measured by recording the absolute pressure drop using a Kulite pressure transducer that was operating in the pressure range of 0-10 × 105 Pa. This pressure transducer is calibrated in the operating range within 400 Pa against a Dru¨ck DPI605 reference pressure transducer calibrated by the Laboratoire National d’Essais. The Kulite pressure transducer is kept at a temperature slightly higher than the experimental temperature to avoid liquid condensation in its measuring chamber. A computer that was fitted with an acquisition card is used to record temperatures and pressures as a function of time.
Ind. Eng. Chem. Res., Vol. 46, No. 1, 2007 235
Figure 2. N2O absorption into a 40 wt % MDEA aqueous solution.
The reactor is connected to a solvent volumetric pump and an acid gas reservoir (see Figure 1), in which the pressure and temperature are recorded. The volume of solvent introduced into the reactor is known within 0.1 × 10-6 m3. The amount of acid gas transferred into the reactor is calculated from the pressure drop inside the acid gas reservoir of a known volume ((302.7 ( 0.1) × 10-6 m3). Mode of Operation. The solvent was prepared under vacuum from water and alkanolamine that had been degassed independently. The solvent composition was known by means of differential weighing to within 0.01 g of each component. Degassed solvent was introduced into the reactor using the volumetric pump. Stirring and heating were started. After the experimental temperature was attained, the liquid-phase volume in the reactor was calculated from the volume introduced, using the density correlation of Amararene et al.,17 and taking into account the temperature difference between the temperature inside the volumetric pump and the reactor. The absorbed gas partial pressure was obtained from the measured pressures:
Pabsorbed gas ) P - PI
303-333 K temperature range, for MDEA compositions in the 20-40 wt % range, and for various liquid-phase stirring speeds. The N2O physical absorption rate was calculated from the relation
φN2O ) kL(CN2O,int - CN2O,liquid bulk)
We assumed that the gas-phase resistance is negligible during pure N2O absorption. The interfacial concentration of N2O was obtained from Henry’s law:
CN2O,int )
Liquid-Side Mass-Transfer Characterization The liquid-phase mass-transfer coefficient was determined from the physical absorption of pure N2O into alkanolamine aqueous solutions. Measurements were conducted in the
P N2 O
(3)
HN2O
The N2O concentration in the liquid bulk was determined from the accumulated N2O absorbed into the alkanolamine aqueous solution:
CN2O,liquid bulk ) -
(1)
where PI is the inert gas pressure in the reactor. PI was measured before the acid gas injection and took solvent vapor pressure and nitrogen partial pressure into consideration. A controlled amount of acid gas from the thermostated highpressure gas reservoir was then fed into the vapor phase of the reactor through thermostated tubing. The pressure decrease versus time, as a result of acid gas absorption through the horizontal gas/liquid interface, was recorded. This experimental operation mode was followed for the different absorbing systems involved in this study. An example of the rough experimental results obtained for a typical absorption experiment is displayed in Figure 2.
(2)
VG (P - PN2O,ini) RTVL N2O
(4)
The N2O absorption rate was calculated by the mass balance in the Lewis cell gas phase:
φ N2 O ) -
VG dPN2O ART dt
(5)
A pressure-time relation was obtained from eqs 2-5:18
[
PN2O +
ln
]
VG HN2O (P - PN2O,ini) RTVL N2O ) PN2O,ini -AkL
(
)
1 RT + t (6) VGHN2O VL
The liquid-side mass-transfer coefficient was determined from the slope of the curve corresponding to the logarithmic transformation of the pressure-time absorption profile. The N2O
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Table 1. Gas-Side Mass-Transfer Coefficient Measured from H2S Absorption into DEA Aqueous Solvents temperature (K)
DEA (wt %)
total pressure (× 105 Pa)
PH2S (× 105 Pa)
kG (× 10-5 mol m-2 s-1 Pa-1)
323.15 333.15 313.15 313.15
30 20 10 10
4.510 4.656 4.687 2.417
0.036 ((0.004) 0.033 ((0.004) 0.038 ((0.004) 0.028 ((0.004)
2.81 × 10-2 (σ ) 0.24 × 10-2) 2.71 × 10-2 (σ ) 0.07 × 10-2) 2.82 × 10-2 (σ ) 0.05 × 10-2) 3.49 × 10-2 (σ ) 0.10 × 10-2)
Table 2. CO2 and H2S Diffusion Coefficients at Infinite Dilution in Nitrogen Diffusion Coefficient at 313 K (× 10-4 m2/s) gas
Wilke and Lee
Diffusion Coefficient at 333 K (× 10-4 m2/s)
Fuller and Giddings
CO2 H2S
0.092 0.095
Ptotal ) 2 × 105 Pa 0.090 0.094
CO2 H2S
0.037 0.038
Ptotal ) 5 × 105 Pa 0.036 0.037
Henry’s constant was calculated from a correlation, representing, to within 3%, the more-consistent data available in the literature. The model was similar to that of Wang et al.;19 the new correlation’s parameters were fitted on the data selected by Li and Mather.20 We used a well-known dimensionless equation to correlate the data:21 where
Sh ) 0.230Re
Sc
0.430
(7)
kLdcell Dgas
(8)
FNdRushton2 µ
(9)
µ FDgas
(10)
Sh ) Re )
0.661
Sc )
The physicochemical parameters used in these equations were obtained from the Amararene et al. correlations17 for solvent density; from the Hsu and Li correlations22 for solvent viscosity; and from the Versteeg and van Swaaij correlations23 for N2O diffusivity. Gas-Side Mass-Transfer Characterization The gas-side mass-transfer coefficient was determined from the absorption rates of H2S into DEA aqueous solutions, in which an instantaneous reaction occurred between the absorbed acid gas and the alkanolamine. The total pressure inside the absorption cell was monitored by nitrogen injection. Critchfield24 observed that the gas-side mass-transfer coefficient was drivingforce-dependent, assuming no contribution from the liquid-phase resistance. This means that the liquid-phase resistance cannot be neglected. Taking these observations into consideration, the measurements were made with low driving forces and DEA concentrations of >1 kmol/m3, to minimize the effect of the liquid-phase gradients. For each experiment, the following inequality was checked and fulfilled:
kG < where
kLEi,H2S
(11)
HH2S
( )
HH2S DDEA (CDEA,bulk) Ei,H2S ) 1 + DH2S PH2S
Wilke and Lee
Fuller and Giddings
0.103 0.106
0.100 0.104
0.041 0.042
0.040 0.042
where CDEA,bulk is expressed in terms of mol/m3, HH2S is expressed in terms of Pa m3/mol, and PH2S is given in Pascal. The gas-side mass-transfer coefficient (kG) was obtained directly from the partial pressure of H2S in the reactor and the corresponding measured absorption rate:
φH2S ) kGPH2S
The gas-side mass transfer coefficient was measured for different operating conditions. The values reported in Table 1 are the mean values of five measurements; σ is the standard deviation observed in the five experimental values. Versteeg et al.21 reported the dependence of the gas-side masstransfer coefficient on gas diffusivity. The H2S and CO2 diffusion coefficients reported in Table 2 were estimated using the predictive methods of Wilke and Lee or Fuller and Giddings,25 and they are shown to differ by 0 (30b)
At the bulk side, x˜ ) 1:
Boundary conditions at the bulk side assume chemical equilibrium for all species.
For nonvolatile species: C ˜ i(1,t˜) ) 1
At the gas/liquid interface, fluxes of nonvolatile species are equal to zero. For the volatile species, we assume continuity of the mass-transfer rate in the gas and liquid near the interface:
(27b)
)
For volatile species: C ˜ i(1,t˜) )
(24)
)
∂C ˜ i(x,t˜) ∂φ˜ - ziD ˜ iC ˜ i(x,t˜) + R ˜ i (28) ∂x˜ ∂x˜
For nonvolatile species: Ci(x,0) ) Ci,bulk (for 0 ex eδ) (23b)
Ci(x,δ) ) Ci,bulk
δ2 DiCi,bulk
(27a)
NC
(22)
The set of coupled differential and algebraic equations was solved with the appropriate initial and boundary conditions. The initial concentration profiles of the transferred species are linear, taking into account the boundary conditions at the gas/liquid interface and the bulk concentrations. For nonvolatile species, the initial concentrations in the mass-transfer zone are equal to the bulk concentrations.
(26)
Equations 19 and 22 become
NC
ziCi(x,t) ) 0 ∑ i)1
(25b)
where D1 is the diffusion coefficient of the first volatile component.
∂Cq(x,t)
∑zqDq RT q)1
∀t>0
Ci,bulkHi Pi
∀ ˜t > 0 ∀ ˜t > 0
(31a) (31b)
The set of partial differential equations was solved by the finite difference iterative method. The numerical treatment was identical to that used by Cadours and Bouallou.27 This model was used to represent the opposite mass transfer when H2S is absorbed into a CO2-loaded DEA aqueous solution
Ind. Eng. Chem. Res., Vol. 46, No. 1, 2007 239
Figure 4. Absorption of CO2 into H2S-loaded DEA aqueous solution. Conditions: 10 wt % DEA, 313.15 K, RH2S ) 0.015 molH2S/molDEA, and PCO2 ) 3400 Pa.
Figure 5. Absorption of H2S into CO2-loaded DEA aqueous solution. Conditions: 10 wt % DEA, 313.15 K, RCO2 ) 0.015 molCO2/molDEA, and PH2S ) 2800 Pa.
or when CO2 is absorbed into a H2S-loaded DEA aqueous solution. A simplified mechanism reduced to main reactions was selected:
CO2 + 2R2NH T R2NH2+ + R2NCOOH2S + R2NH T R2NH2+ + HSWe first considered the overall finite rate reaction between CO2 and DEA, where only the alkanolamine was involved in the deprotonation step of the zwitterion mechanism that was proposed by Caplow.10 Because absorption rate measurements have been achieved with very low acid-gas loading, and under specific operating conditions to set the zwitterion formation as the kinetic rate-limiting step, the other reactions proposed by Rinker et al.15 were not taken into account. We considered a reversible reaction, with reaction order of 1 in CO2 and DEA concentrations. The reaction order value of the reverse reaction was assumed to be 2, with reaction order 1, with respect to each ionic compound. The experimentally measured reaction rate was directly introduced into the numerical model for the
forward reaction between CO2 and DEA. The backward rate constant was then deduced from the reaction equilibrium constant. Reaction between H2S and DEA was assumed to be instantaneous, with respect to mass transfer, because it involves only a proton transfer. In the numerical treatment, this instantaneous equilibrium was modeled as a finite-rate reaction with a very high reaction rate constant, as proposed by Glasscock and Rochelle.9 A reaction order of 1, with respect to each component, was assumed. The reaction rate between H2S and DEA was assumed to be 100 times greater than the reaction rate between CO2 and DEA. As an initial approximation, the equilibrium constants were calculated from the Kent and Eisenberg correlations.28 The Henry’s law constant of CO2 in DEA aqueous solutions was obtained from Blanc and Demarais.12 The Henry’s law constant of H2S in aqueous DEA solutions was directly obtained from H2S solubility in pure water. The CO2 diffusion coefficient in the liquid phase at 298 K was determined from the Blanc and Demarais correlations.12 The H2S diffusion coefficient in the solvent was estimated from the diffusivity in pure water. The
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DEA diffusion coefficient was also obtained from the correlations reported by Hikita et al.29 The influence of the temperature was taken into consideration, according to the relation proposed by Hikita et al.29 for aqueous alkanolamine solutions. The diffusion coefficients of the different ionic species were assumed to be equal to the DEA diffusion coefficient. The model was used to highlight the opposite mass transfer in the case of the absorption of an acid gas into a solvent previously loaded with another acid gas. First, we considered CO2 absorption into a solvent with a loading of 0.015 molH2S/ molDEA. The measured reaction rate was not significantly influenced by H2S in the solvent, taking into account the experimental uncertainties reported in this work. The concentration profiles in the liquid phase, which have been calculated with the numerical model previously presented, show that the CO2 absorption in the solvent had a direct influence on the H2SDEA equilibrium (see Figure 4). The model calculated a significant molecular H2S formation in the mass-transfer film. CO2 absorption in the solvent was coupled to an H2S desorption phenomenon. The relative quantities of acid gas transferred were not observable with the equipment used in this work, and the reaction rate does not seem to be influenced by H2S, under the uncertainties associated with the experimental procedure. On the other hand, we observed experimentally that H2S absorption was significantly influenced by dissolved CO2 in solvent. The phenomenon modeling predicted CO2 desorption simultaneously with the H2S absorption (see Figure 5). The high reactivity of H2S with DEA led to a high H2S absorption rate and, as a result, to an important modification of the interface conditions. The equilibrium between CO2 and DEA was significantly displaced by the high H2S absorption rate and led to CO2 desorption at the interface. These experimental results and their representation highlight the importance of the competitive mass transfer, when the absorption of two gases is considered. The importance of the reverse reactions must also be taken into consideration to represent the phenomenon. Integration of this numerical tool into absorption simulation models will lead to improvement in the performance prediction of natural gas treatment units. Conclusion In this work, the kinetics of the zwitterion formation reaction between CO2 and diethanolamine (DEA) were determined from the CO2 absorption rate obtained in a Lewis cell. The total pressure control of the measuring cell with nitrogen made it possible to study the reaction under pseudo-first-order conditions, with unloaded solvent. Experiments performed with very low acid-gas loadingslower than 0.015 mole of acid gas per mole of DEAsshowed the influence of the first acid gas dissolved in the solution when the absorption rates of a second acid gas are measured. The experimental observations were represented with a general model coupling mass transfer and chemical reactions. This procedure involved a numerical approach to solve the set of coupled differential and algebraic equations. This model was used to represent the opposite mass transfer when CO2 or H2S is absorbed in a water-DEA solution partially loaded with the other acid gas. Real-time measurements of the quantity of CO2 and H2S present in the gas phase would be useful to arrive at a better understanding of this phenomenon. Nomenclature A ) gas/liquid interfacial area (m2) Ci ) concentration of component i (mol/m3)
Di ) diffusivity of component i (m2/s) dcell ) internal diameter of the reactor (m) dRushton ) diameter of the Rushton turbine in the liquid phase of the reactor (m) E ) enhancement factor Ei ) enhancement factor in instantaneous reaction region F ) Faraday’s constant; F ) 96 489 C/mol H ) Henry’s law constant in the concentration scale (Pa m-3 mol-1) Ha ) Hatta’s number k ) kinetic rate coefficient kG ) gas-side resistance mass transfer (mol m-2 s-1 Pa-1) kL ) liquid-side resistance mass transfer (m/s) NC ) total number of components NR ) total number of reactions P ) pressure (Pa) R ) ideal gas constant; R ) 8.314 J K-1 mol-1 Ri ) production term for component i (mol m-3 s-1) Re ) dimensionless Reynolds number Sc ) dimensionless Schmidt number Sh ) dimensionless Sherwood number t ) time (s) T ) temperature (K) V ) volume (m3) x ) distance from interface (m) zi ) ion charge for component i Greek Symbols R ) acid gas loading (molacid gas/molalkanolamine) β ) reaction order δ ) film thickness (m) λ ) stoichiometric coefficient µ ) viscosity (Pa s) F ) density (kg/m3) φ ) absorption rate (mol m-2 s-1) φ ) electrostatic potential (V/m) Subscripts and Superscripts bulk ) bulk G ) gas I ) Inert (solvent vapor pressure + inert gas partial pressure) ini ) initial time, corresponding to the absorbed gas injection int ) interface L ) liquid ∼ ) dimensionless notation Component AbbreViations CO2 ) carbon dioxide DEA ) diethanolamine H2S ) hydrogen sulfide HS- ) hydrosulfide ion N2O ) nitrous oxide OH- ) hydroxide ion R2NH ) primary or secondary alkanolamine R2NH2+ ) protonated primary or secondary alkanolamine R2NCOO- ) primary or secondary alkanolamine carbamate Literature Cited (1) Van Krevelen, D. W.; Hoftijzer P. J. Kinetics of gas-liquid reactions. Part 1: general theory. Recl. TraV. Chim PayssBas 1948, 67, 563. (2) Versteeg, G. F.; Kuipers, J. A. M.; van Beckum, F. P. H.; van Swaaij, W. P. M. Mass transfer with complex reversible chemical reactionssI. Single reversible chemical reactions. Chem. Eng. Sci. 1989, 44, 2295.
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(17) Amararene, F.; Balz, P.; Bouallou, C.; Cadours, R. Lecomte, F.; Mougin, P. Richon, D. Densities of water + diethanolamine + methanol and water + N-methyldiethanolamine + methanol at temperatures ranging from (283.15 to 353) K. J. Chem. Eng. Data 2003, 48, 1565. (18) Lemoine, B. Absorption of acid gases by aqueous MDEA solutions. Acquisition of data and kinetic and thermodynamic parameters, Ph.D. Thesis, Paris School of Mines, Paris, France, 1995. (19) Wang, Y. W.; Xu, S.; Otto, F. D.; Mather, A. E. Solubility of N2O in alkanolamines and in mixed solvents. Chem. Eng. J. 1992, 48, 31. (20) Li, Y.-G.; Mather, A. E. Correlation and prediction of the solubility of N2O in mixed solvents. Fluid Phase Equilib. 1994, 96, 119. (21) Versteeg, G. F.; Blauwhoff, P. M. M., van Swaaij, W. P. M. The effect of diffusivity on gas-liquid mass transfer in stirred vessels, experiments at atmospheric and elevated pressures. Chem. Eng. Sci. 1987, 42, 1103. (22) Hsu, C.-H.; Li, M.-H. Viscosities of aqueous blended amines. J. Chem. Eng. Data 1997, 42, 714. (23) Versteeg, G. F.; van Swaaij, W. P. M. Solubility and diffusivity of acid gases (CO2, N2O) in aqueous alkanolamine solutions. J. Chem. Eng. Data 1988, 33, 29. (24) Critchfield, J. E. CO2 absorption/desorption in methyldiethanolamine solutions promoted with monoethanolamine and diethanolamine: mass transfer and reaction kinetics, Ph.D. Thesis, The University of Texas, Austin, TX, 1988. (25) Reid, R.; Prausnitz, J.; Poling, B. The Properties of Gases and Liquids, 4th Edition; McGraw-Hill: New York, 1987. (26) Danckwerts, P. V.; Sharma, M. M. The absorption of carbon dioxide into solutions of alkalis and amines. Chem. Eng. 1966, 10, 244. (27) Cadours, R.; Bouallou, C. Rigorous simulation of gas absorption into aqueous solutions. Ind. Eng. Chem. Res. 1998, 37, 1063. (28) Kent, R. L.; Eisenberg, B. Better data for amine treating. Hydrocarbon Process. 1976, 55, 87. (29) Hikita, H.; Ishikawa, H.; Uku, K.; Murakami, T. Diffusivities of mono-, di-, and triethanolamines in aqueous solutions. J. Chem. Eng. Data 1980, 25, 324.
ReceiVed for reView January 5, 2006 ReVised manuscript receiVed September 27, 2006 Accepted September 29, 2006 IE060019W
