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Ind. Eng. Chem. Res. 2010, 49, 1150–1159
CO2 Absorption into Mixed Aqueous Solutions of 2-Amino-2-hydroxymethyl-1,3-propanediol and Piperazine Francis Bougie and Maria C. Iliuta* Chemical Engineering Department, LaVal UniVersity, Quebec, Canada, G1V 0A6
Solubility data of CO2 in aqueous mixtures of 2-amino-2-hydroxymethyl-1,3-propanediol (AHPD) and piperazine (Pz) were measured over a range of temperature from 288.15 to 333.15 K and for total amine concentrations up to 3.1 kmol · m-3. The CO2 partial pressure was kept within 0.21-2 637 kPa using a vapor-liquid equilibrium (VLE) apparatus based on a static-analytic method. The solubility of N2O in the Pz-AHPD aqueous solutions was also performed in order to determine, with the N2O analogy, the Henry’s law constant of CO2 in these solutions. The experimental data for the ternary system AHPD-CO2-H2O were correlated using a modified Pitzer’s thermodynamic model for the activity coefficients combined with the virial equation of state for representing the fugacity coefficients. The solubility of carbon dioxide in aqueous solutions of mixed amine (Pz + AHPD) was predicted by supposing that the parameters characterizing the single amine systems are essential for describing the quaternary system behavior. 1. Introduction In the past few decades, removal of CO2 has become one of the most important environmental issues facing the world community. This has motivated intensive research on CO2 capture where new and more energy-efficient absorbents are essential. Actual industrial CO2 absorption processes use aqueous solutions of alkanolamines. For technical, economical, and environmental concerns, this technique is widely applied for (i) acid gases (CO2, H2S) removal during natural gas sweetening and (ii) CO2 capture from fossil-fuel-fired power plants, as well as some other important industries such as chemical and petrochemical, steel, aluminum, and cement production. Industrially more often used alkanolamines are monoethanolamine (MEA), diethanolamine (DEA), diisopropanolamine (DIPA), N-methyldiethanolamine (MDEA), and 2-amino-2methyl-1-propanol (AMP).1 The choice of a certain amine is mainly based on the absorption capacity, reaction kinetics, and regenerative potential and facility. The key advantage of the primary and secondary alkanolamines such as MEA and DEA is their fast reactivity due to the formation of stable carbamates. Conversely, this will lead to very high solvent regeneration cost. On the absorption capacity side, they have the drawback of a relatively low CO2 loading (limited to 0.5 mol CO2/mol amine). Tertiary alkanolamines, like MDEA, have low reactivity with respect to CO2, due to the exclusive formation of bicarbonates by CO2 hydrolysis. However, this will lead to a very low solvent regeneration cost. Another advantage of these amines is the high CO2 theoretical loading capacity of 1 mol of CO2/mol of amine. The application of sterically hindered alkanolamines (SHA), e.g., AMP, in gas-treating technology offers absorption capacity, absorption rate, selectivity, and degradation resistance advantages over conventional amines for CO2 removal from gases.2 SHA form unstable carbamates due to the hindrance of the bulky group adjacent to the amino group. Hydrolysis of the voluminous carbamates leads to a preferential bicarbonate formation process, resulting in the theoretical loading capacity up to 1.0. Reaction kinetics significantly higher than those related to tertiary amines, coupled with a low solvent regeneration cost * To whom correspondence should be addressed. Tel.: 1 (418) 6562204. Fax: 1 (418) 656-5993. E-mail address: maria-cornelia.iliuta@ gch.ulaval.ca.
offer to SHA important industrial advantages. The use of blended alkanolamines solutions has also recently become very attractive because of the combination of each amine advantages: a fast reactivity from a primary or secondary alkanolamine coupled with the high absorption capacity and low solvent regeneration cost from a tertiary or sterically hindered alkanolamine. In our laboratory, extensive studies of CO2 capture in membrane contactors using activated (piperazine) aqueous SHA solutions are in progress. A set of four SHA was chosen.3 It concerns AMP, a simple hindrance form of MEA, and three SHA derived from AMP: 2-amino-2-methyl-1,3-propanediol (AMPD), 2-amino-2-ethyl-1,3-propanediol (AEPD), and 2-amino2-hydroxymethyl-1,3-propanediol (AHPD). The kinetics of these SHA has been discussed previously3 as well as the influence of the addition of an activator (Pz) in AHPD solutions.4 In order to have a more accurate insight of the properties of the studied solutions, data concerning the solubility of CO2 and N2O in aqueous amine solutions are needed respectively to (i) determine the equilibrium loading of CO2 in these solutions for a wide range of temperature, solutions concentration, CO2 partial pressure, and (ii) determine the Henry’s law constant of CO2 in these solutions by the application of the widely known N2O analogy. Henry’s law constants are particularly useful to calculate the CO2 diffusion coefficient in solution from values of the ratio DCO21/2/HCO2. This ratio is found by the use of the wetted wall column contactor as explain in our previous works.3,4 The number of studies about CO2 solubility and Henry’s law constant in aqueous solutions of AMPD, AEPD, or AHPD is quite low,5-11 and disagreements were found between the reported equilibrium solubility of CO2 in AHPD solution between the study of Park et al.8 and Le Tourneux et al.9 Furthermore, except for AMP, no study was found concerning the equilibrium solubility of CO2 and of N2O in Pz-activated aqueous solutions of these SHA. The main objective of this work is the experimental characterization and the thermodynamic modeling of the CO2 solubility in aqueous Pz-activated AHPD solutions. The solubility measurements were performed in a static vapor-liquid equilibrium apparatus for a large range of temperature, solution concentrations, and CO2 partial pressures. A thermodynamic model based
10.1021/ie900705y 2010 American Chemical Society Published on Web 11/24/2009
Ind. Eng. Chem. Res., Vol. 49, No. 3, 2010
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Figure 1. Schematic diagram of the solubility apparatus: (A) equilibrium cell; (B) magnetic rod; (C) platinum resistance thermometer; (Di) gears; (E) coil; (F) pressure transducer (F1, low pressure values; F2, high pressure values); (G) valve; (H) stirrer; (I) temperature controller; (J) computer; (K) circulating bath; (L) variable volume press for liquid introduction; (M) small gas cylinder; (N) gas cylinder; (Oi), needle valve; (Pi) valves; (Q) laboratory oven.
on the Pitzer’s equations for the activity coefficients coupled with the truncated virial equation of state for representing the nonideality of the vapor phase was used to correlate the experimental data for the ternary AHPD-CO2-H2O system. The solubility of carbon dioxide in aqueous solutions of mixed amine (Pz + AHPD) was predicted by supposing that the parameters characterizing the single amines system are essential for describing the quaternary system behavior. The solubility of N2O in the Pz-AHPD aqueous solutions was also tested in order to determine, with the N2O analogy, the Henry’s law constant of CO2 in these solutions. To our knowledge, similar data are not available in the open literature. 2. Experimental Section 2.1. Reagents. Aqueous Pz-AHPD solutions were prepared with degassed distilled water, 2-amino-2-hydroxymethyl-1,3propanediol, and piperazine. The amines (from Laboratoire MAT, Quebec, Canada) had a minimum purity of 99.9% and were used without further purification. CO2 and N2O gases were of commercial grade with a minimum purity of 99.5% and were supplied by Praxair. 2.2. Apparatus and Procedures. The experimental setup for the solubility measurements (Armines, France) used in this work is shown in Figure 1. It consists of an equilibrium cell made of TA6 V titanium with an internal volume of about 1.15 × 10-4 m3. The equilibrium cell is equipped with a magnetic rod covered with titanium, and the cell is located in a modified XU027 laboratory oven from France Etuves. This oven came with a C3000 temperature controller (by France Etuves) which allows temperature control of (0.1 K. A special feature of this apparatus is the addition in the oven of a coil refrigerated with a thermostatted bath (K-12108-10 from Cole-Palmer). This coil allowed us to made solubility measurement under room temperature (273.15-303.15 K) with the same temperature precision. Pressure in the cell was measured by means of one or two of the two installed absolute pressure transducers (Druck
PTX-611, 0-100 kPa and 0-16 000 kPa) according to the pressure range. Two 100 Ω platinum resistance thermometer were used for temperature measurements of the equilibrium cell. Liquid introduction inside the equilibrium cell was made with a variable volume press (stainless steel 316, internal diameter of 3.002 × 10-2 m). This press has been equipped with a linear encoder (Heidenhain, LS487C) which allowed knowing the exact longitudinal position of the piston in the press with an accuracy of (2 × 10-6 m. Gas introduction in the cell was made by a thermostatted small gas cylinder with an internal volume of about 7 × 10-5 m3. This small gas cylinder was equipped with a Druck PTX-611 0-16 000 kPa absolute pressure transducer. A standard experimental run consisted of a sequence of successive step. First, the amine aqueous solution (total amine molalities from 0.91 to 4.36 mol · kg-1) was prepared to its specific concentration by gravimetric method using a Mettler Toledo AE204 balance with a precision of (0.0001 g. Then, the solution was degassed under vacuum and the amine concentration of the resulting solution was checked with HCl and a methyl red-bromocresol green pH indicator mix to verify the possible change in concentration due to solvent or solute lost. The degassed solution was then transferred under vacuum inside the variable volume press and subsequently, with the piston, in the equilibrium cell previously brought to vacuum. The equilibrium cell was heated to the desired temperature and the solution was agitated. At this stage, the vapor pressure of the solution was measured by the low pressure transducer. This was followed by the introduction of the gas to be absorbed (CO2 or N2O) in the equilibrium cell via the small gas cylinder. The introduced gas mole number was calculated by using the cylinder volume, its temperature as well as the observed pressure drop in the cylinder after the gas introduction. The system equilibrium was reached when the pressure inside the equilibrium cell was varying less than 0.5% for at least 30 min. It took about 2 h after the gas introduction for chemical absorption
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Ind. Eng. Chem. Res., Vol. 49, No. 3, 2010
Table 1. Henry’s Constant for the Solubility of Carbon Dioxide in Pure Water (273 e T/K e 473)12,a
a
ln
ACO2,H2O
BCO2,H2O
CCO2,H2O
DCO2,H2O
192.876
-9624.4
0.01441
-28.749
m (T,Pwsat)/(MPa · kg · mol-1) HCO 2,H2O
) ACO2,H2O + BCO2,H2O/(T/K) + CCO2,H2O(T/K) + DCO2,H2O ln(T/K).
of CO2 and 30 min for physical absorption of N2O. The difference between the introduced and the remaining gas mole number in the head space of the equilibrium cell was then calculated which lead to the concentration of absorbed gas in the solution. 3. Thermodynamic Modeling of the Vapor-Liquid Equilibrium 3.1. Chemical Equilibrium in the Liquid Phase. Due to chemical reactions in the liquid phase, carbon dioxide can be found in the liquid phase in both neutral and nonvolatile ionic form. The model applied to correlate/predict the solubility of carbon dioxide in aqueous solutions of AHPD and Pz + AHPD considers the following equilibriums for the chemical species in the liquid phase: the formation and dissociation of bicarbonate reactions 1 and 2, the autoprotolysis of water (reaction 3), the protonation of AHPD reactions 4, the formation of AHPD carbamate (reactions 5), the protonation and diprotonation of piperazine (reactions 6 and 7), and the formation of piperazine carbamate, piperazine dicarbamate, and protonated piperazine carbamate (reactions 8-10). In reactions 4 and 5, “R” denotes the (HO-CH2)2-C group in AHPD. K1
CO2 + H2O y\z HCO3- + H+ K2
HCO3- y\z CO32- + H+ K3
H2O y\z H+ + HOK4
RNH2 + H+ y\z RNH3+ K5
RNH2 + HCO3- y\z RNHCOO- + H2O
(1)
(2)
(3)
(4)
(5)
In the addition of the above equilibrium equations, overall species mole and charge balances must be satisfied. In the balance equations for carbon dioxide, AHPD and Pz in the liquid ˜ Pz denote the stoichiometric phase (eqs 12-14) m ˜ AHPD and m molalities of AHPD and Pz, respectively, and R denotes the CO2 loading in the solutions expressed as total moles of CO2 absorbed both chemically and physically per mole of amine. m ˜ AHPD ) mRNH2 + mRNHCOO- + mRNH3+ m ˜ Pz ) mPz + mPzH+ + mPzH22+ + mPzCOO- + mPz(COO-)2 + mPzH+COO-
Pz + H+ y\z PzH+
(6)
mPzCOO- + 2mPz(COO-)2 + mPzH+COO-
PzH+ + H+ y\z PzH22+ -
K8
-
Pz + HCO3 y\z PzCOO + H2O K9
PzCOO- + HCO3- y\z Pz(COO-)2 + H2O K10
PzCOO- + H+ y\z PzH+COO-
(7)
(8)
(9)
(10)
The condition for chemical equilibrium for a chemical reaction R is KR(T) )
∏a
νi,R i
(R ) 1, ..., 10)
i
where ai is the activity of species i.
(11)
(14)
mH+ + mRNH3+ + mPzH+ + 2mPzH22+ ) mOH- + mHCO3- + 2mCO32- + mRNHCOO- + mPzCOO- + 2mPz(COO-)2
(15)
Solving this set of fourteen independent equations (eqs 11-15) for a given temperature and solution overall molality results in the true (equilibrium) composition of the liquid phase, expressed as the molality of each species, needed for solving the vapor-liquid equilibrium equations. 3.2. Vapor-liquid Equilibrium. Only water is treated as a solvent species. Carbon dioxide, AHPD, Pz, and the ions are treated as solute species. The reference state for the chemical potential of water is the pure liquid at the system temperature and pressure. The chemical potential of a solute species is a 1 molal solution in pure water at the system temperature and pressure. The condition of vapor-liquid equilibrium (VLE) is applied in order to calculate the total pressure and the composition of the gas phase. The extended Raoult’s law is used to express the VLE for water (eq 16) and the extended Henry’s law is used to express the equilibrium for carbon dioxide (eq 17):
[
]
Vw(P - Pwsat) aw ) Pywφw RT
[
*,m m mCO2γCO H (T, Pwsat) exp 2 CO2,H2O K7
(13)
R(m ˜ AHPD + m ˜ Pz) ) mCO2 + mHCO3- + mCO32- + mRNHCOO- +
Pwsatφwsat exp K6
(12)
∞ VCO (P - Pwsat) 2,H2O
RT
]
(16)
)
PyCO2φCO2
(17)
Because the vapor pressures of both amines used in this work are very low in the temperature range considered here, the presence of AHPD and piperazine in the vapor phase was neglected. 3.3. Thermodynamic Properties. The VLE calculation requires the knowledge of the following properties: (i) Henry’s constants for the solubility of carbon dioxide in pure water on the molality scale, HmCO2,H2O(T,Psat w ), were taken from the work of Rumpf and Maurer12 (Table 1). (ii) The temperature dependent equilibrium constants for reactions 1-10 are given in Table 2. Except for the equilibrium constant K5 which was calculated based on the experimental data for the system AHPDwater-CO2, all other constants were taken from the works of Edwards et al.,13 Perrin,14 Hetzer et al.,15 and Ermatchkov et al.16
Ind. Eng. Chem. Res., Vol. 49, No. 3, 2010 sat
(iii) The vapor pressure Pw and the molar volume Vw of pure water were taken from the work of Saul and Wagner.17 (iv) The fugacity coefficients φi were calculated using a truncated virial equation of state. Pure component second virial coefficients BH2O,H2O and BCO2,CO2 for water and carbon dioxide, respectively, were calculated on the basis of the data given by Dymond and Smith.18 The mixed second virial coefficients BCO2,H2O were taken from Hayden and O’Connell19 and correlated as a function of temperature. ∞ of carbon dioxide (v) The partial molar volumes VCO 2,H2O dissolved at infinite dilution in water were calculated as recommended by Brelvi and O’Connell20 and correlated as a function of temperature. 3.4. Pitzer’s GE Model for Activity Coefficients and Interaction Parameters. In the literature, several models are used to characterize the VLE of CO2 in aqueous amines solutions. Among them, the Kent and Eisenberg21 and the Deshmukh and Mather22 models are frequently used. However the former one does not take into account the activity coefficients in solution and the latter is limited to low concentration because the activity coefficients are calculated with Guggenheim’s equation.23 In this research, a more rigorous model is then used to cover the wide range of amine concentrations. Activity coefficients of both neutral and ionic species were calculated using a modified Pitzer model for the excess Gibbs energy of aqueous electrolyte solutions:23,24 E
G ) f1(I) + RTnwMw
∑ ∑ m m λ (I) + ∑ ∑ ∑ m m m τ i
j ij
i
i*w j*w
j
k ijk
i*w j*w k*w
(18) f1(I) is a modified Debye-Hu¨ckel term depending on ionic strength, temperature, and solvent (water) properties: f1(I) ) -Aφ(4I/1.2) ln(1 + 1.2√I)
(19)
where I is the ionic strength and Aφ is the Debye-Hu¨ckel parameter for the osmotic coefficient I)
1 2
∑mz
2
(20)
i i
i
(
e2 1 Aφ ) (2πNAFw)1/2 3 4πε0DkT
)
3/2
(21)
The dielectric constant of pure water, D, was taken from the work of Bradley and Pitzer.25 λij(I) is the ionic strength dependent second virial coefficient: (1) 2 -x λij(I) ) β(0) ij + βij [(2/x )(1 - (1 + x)e )]
(22)
where x ) 2I.
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The influence of temperature on the binary interaction (1) parameters β(0) ij and βij is approximated by the relation: f(T) ) q0 +
q1 T/K
(23)
The ternary interaction parameters τijk are considered independent of temperature. The equation for the activity coefficients of dissolved species follows from the appropriate derivative of GE, and water activity is calculated from the Gibbs-Duhem equation:
(
)
√I 2 ln(1 + 1.2√I) + + 1.2 1 + 1.2√I βjk(1) 2 mjλij(I) - zi2 mjmk 2 1 - 1 + x + Ix j*w j*w k*w
) -Aφzi2 ln γ*,m i
∑
∑∑
(
) ]
[ (
x2 -x e +3 2 j*w I1.5
∑ ∑mm τ j
k*w
∑ ∑ m m (β M (2 ∑ ∑ ∑ m m m τ + ∑ m )
ln aw) Mw 2Aφ
1 + 1.2√I
w
-
i
j
(0) ij
i
j
k ijk
(24)
)
-x + β(1) ij e ) -
i*w j*w
i*w j*w k*w
k ijk
i
i*w
(25)
All interaction parameters used in this work are given in Table 3. 3.4.1. System AHPD-CO2-H2O. Interaction parameters for the ternary system AHPD-CO2-H2O were determined on the basis of experimental data taken from the literature7,9 and from the present work. In this system, eight species are present in the liquid phase: CO2, HCO3-, CO32-, RNH2, RNH3+, RNHCOO-, H+, and OH-. Due to the very low concentration of H+ and OH- with respect to the other species, their interactions with all other species were ignored, and therefore, the corresponding interaction parameters were set to zero. Binary and ternary interaction parameters between neutral species, CO2 and RNH2 were considered negligible and were set to zero. Except for the binary interaction parameter between RNH2 and RNH3+, all binary and ternary interaction parameters between RNH2 and any other species were also set to zero. In addition, the ionic strength dependence of the second virial coefficient (eq 22) was neglected for the all interactions except for RNH3+-HCO3-. In order to reduce the number of parameters, all binary and ternary interaction parameters involving species with the same sign of charge were neglected. Only the parameters which were found to have a significant influence on the liquid phase species distribution were optimized based (0) (0) (0) on the experimental data: βCO2,HCO-3 , βCO2,RNH+3 , βHCO3-,RNH+3 , (1) (0) (0) (0) βHCO3-,RNH+3 , βCO2,RNHCOO-, βCO32-,RNH+3 , and βRNH2,RNH+3 . A sensitivity study revealed that all other possible interaction parameters that
Table 2. Equilibrium Constants for Chemical Reactions 1-10a R
AR
BR
CR
DR
ER
FR
ref
T/K
1 2 3 4 5 6 7 8 9 10
-12091.1 -12431.7 -13445.9 0 0 3814.4 2192.3 1570.4 574.2 1517
-36.7816 -35.4819 -22.4773 0 0 0 0 0 0 0
235.482 220.067 140.932 22.61853 213.8527 14.119 10.113 -3.75 -1.587 4.354
0 0 0 0.591854 -2.123369 -1.51 × 10-2 -1.74 × 10-2 0 0 0
0 0 0 -2.360429 ×10-3 7.033246 × 10-3 0 0 0 0 0
0 0 0 2.814271 ×10-6 -7.884854 × 10-6 0 0 0 0 0
Edwards et al.13 Edwards et al.13 Edwards et al.13 Perrin14 this work Hetzer et al.15 Hetzer et al.15 Ermatchkov et al.16 Ermatchkov et al.16 Ermatchkov et al.16
273-498 273-498 273-498 273-323 288-333 278-328 273-323 273-323 273-323 273-323
a
ln KR ) AR/(T/K) + BR ln(T/K) + CR + DR(T/K) + ER(T/K)2 + FR(T/K)3.
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Table 3. Interaction Parameters in Pitzer’s GE Equation for the System AHPD-PZ-CO2-H2Oa parameter
q1
(0) βCO 2,HCO3
2.256
-379.5
(0) βCO + 2,RNH3
-4.547
917.7
(0) βHCO + 3 ,RNH3
0.700
-400.2
(1) βHCO + 3 ,RNH3
1.017
(0) βCO 2,RNHCOO
-6.600
ref
AHPD + CO2 + H2O
this work
PZ+CO2+H2O
Pe´rez-Salado Kamps et al.26
995.2
3.400
(0) βRNH + 2,RNH3
0.300
τCO2,HCO3-,RNH3+
0.0707
0
τHCO3-,HCO3-,RNH3+
0.0480
0
(0) βCO + 2,PzH
0.14624
-187.24
(0) βHCO + 3 ,PzH
0.55489
(1) βHCO + 3 ,PzH
1.8949
(0) βCO + 2,PzH COO
0.55705
-196.84
(0) βPzH +COO-,PzH+COO-
0.096213
-72.2
-1006
-100.0
2.0459
776.48
(1) βPzH +COO-,PzH+COO-
-0.83929
324.79
(0) βPzH +,PzCOO-
-2.0678
776.43
(0) βPzH +,Pz(COO-) 2
-1.3044
440.98
0.34964
subsystem
-1050
(0) βCO 2+ 3 ,RNH3
(0) βPz,PzCOO -
a
q0
-83.169
Ermatchkov et al.27
f(T) ) q0 + q1/(T/K).
appear in the expressions for the activity coefficients (eqs 24 and 25) can be neglected without reducing the accuracy of the VLE representation of this system. Parameters q0, q1, and the ternary ones τijk were fitted simultaneously to the selected experimental data chosen as it will be described in section 4.2. 3.4.2. System AHPD-Pz-CO2-H2O. On the basis of the thermodynamic description of the solubility of carbon dioxide
in a single amine system, AHPD-CO2-H2O and Pz-CO2-H2O, the carbon dioxide solubility in the mixed AHPD + Pz aqueous system was predicted using the available interaction parameters (Table 3). Interaction parameters for the ternary system AHPD-CO2-H2O were determined from the experimental data of the present work and from literature,7,9 as described in the previous section. No other parameters were
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Figure 2. CO2 solubility in water: comparison with literature values. Table 4. Henry’s Law Constants for N2O in Pz (1)-AHPD (2) Solutions T (K)
m1 (mol · kg-1)
m2 (mol · kg-1)
HN2O (kPa · m3 · kmol-1)
288.15 288.15 288.15 298.15 298.15 313.15 313.15 333.15 333.15 333.15
1.1141 2.5881 4.2016 1.1351 3.3634 1.1570 3.3629 1.1149 2.5912 4.2046
0.1114 0.6470 0.1401 0.3405 0.4036 0.5785 0.4035 0.1115 0.6478 0.1402
3111.0 3687.2 4673.5 3865.4 5131.0 6518.2 6960.0 9265.3 12817.5 13158.4
Figure 3. CO2 solubility in Pz aqueous solution: comparison with literature values (m ˜ Pz ) 2.0 mol · kg-1).
found in the literature concerning this ternary system. Interaction parameters for the ternary system Pz-CO2-H2O were taken from the works of Pe´rez-Salado Kamps et al.26 and Ermatchkov et al.27 4. Results and Discussions 4.1. Experimental Setup Verification. To check the validity of the experimental setup and procedures, physical absorption of CO2 in water was made at 293.15 and 313.15 K and for several CO2 partial pressures. The experimental data were compared with literature values28-30 in Figure 2. It is possible to see that our results are in excellent agreement with literature values over the entire pressure range. As expected, the CO2 concentration in water decrease when temperature increase at constant CO2 partial pressure. 4.2. Solubility Measurements. N2O Solubility. The absorption of N2O in Pz-AHPD aqueous solutions ranging from 0.10 to 0.50 kmol · m-3 Pz and 1.0 to 3.0 kmol · m-3 AHPD was performed between 288.15 and 333.15 K. The experimental results, expressed in term of the Henry’s law constant, are indicated in Table 4. The uncertainties of indicated values are calculated to be within 2%. As expected, Henry’s law constant values increase with increasing either temperature for a given solution concentration or amine concentration in aqueous solutions at constant temperature. CO2 Solubility. CO2 solubility measurements were made in three different reactive systems: CO2-Pz-H2O, CO2-AHPD-H2O, and CO2-Pz-AHPD-H2O in order to respectively (i) validate the apparatus and the procedures for chemical absorption at high pressure, (ii) obtain more CO2 solubility data in AHPD solution, needed to check the validity of literature sources between those available (this is necessary to obtain good interaction parameters in the thermodynamic model) and (iii) obtain the CO2 solubility in the mixed
Figure 4. (a) CO2 solubility in AHPD aqueous solution at 298.15 K, comparison with the works of Park et al.8 and Le Tourneux et al.9 (m ˜ AHPD ) 0.9172 mol · kg-1). (b) CO2 solubility in AHPD aqueous solution at 323.15 K, in comparison with the work of Park et al.7 (m ˜ AHPD ) 0.9172 mol · kg-1).
Pz-AHPD aqueous solutions to determine the effect of Pz on the equilibrium solubility of AHPD and to test the prediction capacity of the developed VLE model in representing the experimental data for the quaternary system based on the interaction parameters for the corresponding ternary systems.
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Table 5. CO2 Solubility in AHPD Aqueous Solutions T (K)
mAHPD (mol · kg-1)
PCO2 (kPa)
CO2 loading (mol CO2 · mol-1 AHPD)
298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 298.15 323.15 323.15 323.15 323.15 323.15 323.15 323.15 333.15 284.84 284.82 284.76 284.53 284.70 284.65 284.71 284.78 284.82 303.16 303.18 303.20 303.18 303.16 303.14 303.11 303.12 303.11 333.14 333.13 333.12 333.11 333.11 333.11 333.11 333.13 284.29 284.43 303.26 303.29 303.31 303.29 303.31 303.32 303.27 303.28 333.20 333.11 333.12 333.13 333.12 333.12 333.14 333.22 293.26 293.20 303.16 303.17 303.17 303.17 303.15 303.13 333.17 333.20 333.20 333.08 333.09 333.11 333.22
0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 0.917 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 2.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 3.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000 4.000
0.31448 1.1729 4.4127 27.331 232.00 533.60 1237.6 1938.4 2637.6 5.8978 26.232 68.496 130.17 236.23 526.82 1303.9 2106.0 0.90620 2.3588 6.7321 23.350 100.88 241.76 450.35 927.36 1249.1 0.92003 2.8518 6.8983 15.729 35.435 83.534 259.50 619.37 1034.1 7.6416 25.197 55.973 94.976 152.18 216.84 377.32 683.61 0.51300 1.5500 0.90389 3.3283 8.9937 24.050 59.078 184.42 456.96 775.93 22.759 35.534 59.999 98.446 198.23 416.19 676.89 914.81 0.60554 2.2092 1.1836 4.6753 11.897 40.680 77.589 258.72 24.854 95.209 176.18 373.09 601.44 849.99 1079.1
0.0745 0.1817 0.3652 0.7010 1.0208 1.1497 1.3953 1.6171 1.8545 0.1967 0.4010 0.5830 0.7272 0.8540 1.0287 1.3921 1.6782 0.1602 0.3283 0.5307 0.7551 0.9452 1.0257 1.0825 1.1881 1.2345 0.0940 0.1914 0.3063 0.4396 0.5824 0.7289 0.8948 1.0011 1.0715 0.1209 0.2284 0.3357 0.4310 0.5138 0.5815 0.6940 0.8172 0.1422 0.2953 0.1057 0.2161 0.3399 0.4897 0.6338 0.8054 0.9163 0.9741 0.0796 0.1623 0.2524 0.3399 0.4743 0.6144 0.7063 0.7642 0.1303 0.2569 0.1211 0.2462 0.3651 0.5466 0.6436 0.8125 0.1933 0.3628 0.4585 0.5864 0.6704 0.7331 0.7765
Figure 5. CO2 solubility in Pz-AHPD aqueous solutions at 288.15 and 333.15 K.
Figure 6. Predicted species distribution in the AHPD + CO2 + H2O system at 298.15 K (m ˜ AHPD ) 0.9172 mol · kg-1).
Figure 7. Predicted species distribution in the Pz-AHPD + CO2 + H2O system at 298.15 K (AHPD ) 1.0 kmol · m-3 and Pz ) 0.3 kmol · m-3).
For the system CO2-Pz-H2O, CO2 chemical absorption was made at 313.15 K in a solution containing 2 kmol · kg-1 Pz and up to a total pressure of 2 900 kPa. This pressure is above the highest pressure reached for the CO2 absorption in the mixed solvent, and it is possible to see in Figure 3 that even at this high pressure, the correlation between our data and those of Pe´rez-Salado Kamps et al.26 is particularly good.
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Figure 8. (a) CO2 solubility in aqueous solution of AHPD; experimental results of this work, m ˜ AHPD ) 4.0 mol · kg-1. (b) CO2 solubility in aqueous solution of AHPD; experimental results by Le Tourneux et al.,9 different AHPD molalities.
For the system CO2-AHPD-H2O, some disagreements were found between the reported equilibrium solubility of CO2 by the research group of Park et al.8 when compared to the data of Le Tourneux et al.,9 as it can be seen in Figure 4a. Our results agree very well with those of Le Tourneux et al.,9 which were performed using a different experimental setup. We therefore consider the data from this source reliable to be used in the interaction parameter determination. A verification of some of the data reported in another article by Park et al.7 was also made and shown in Figure 4b. It is possible to notice that Park’s data disagree from our results at pressures larger than around 500 kPa. We therefore decided not to include these data (P > 500 kPa) in the database used for the parameters estimation. Consequently, the number of reliable data for the system CO2-AHPD-H2O is 177: 84 from the work of Le Tourneux et al.,9 17 from that of Park et al.,7 and 76 from this work (Table 5). For the system CO2-Pz-AHPD-H2O, all experimental results are listed in Table S1 (Supporting Information). The CO2 partial pressure and CO2 loading reported in Table 5 and Table S1 have respectively a maximum calculated uncertainty of 0.6% and 1.5%. From Figure 5, it is possible to observe that at constant amine concentrations and CO2 partial pressure an increase in temperature leads to a decrease of the CO2 loading capacity. Furthermore, as expected, at constant temperature, an
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increase in the total amine concentration leads to a decrease of the gas absorption capacity. The absorption of CO2 in the AHPD aqueous solution was shown to be similar to that in the AMP aqueous solution.31 The system pressure increases very slowly during the chemical absorption when the gas is mostly dissolved in nonvolatile ionic form then the pressure shows a sharp increase above the stoichiometric gas-amine ratio in the physical absorption section. An example of the speciation in an aqueous AHPD by carbon dioxide addition, based on the equilibrium model, is shown in Figure 6. The concentration profile for several species is represented as a function of the CO2 loading for an AHPD aqueous solution with a molality of 0.9172 mol · kg-1 at 298 K. Because H+ and OH- concentrations are much lower than the concentrations of all other species, the corresponding curves were not represented. The amine concentration decreases rapidly at CO2 loadings less than about 1 mol/mol of amine, while the bicarbonate and the protonated amine sharply increase. It can be noted that at a CO2 loading of about 1 mol/mol of amine, practically all amine is converted preferentially into the protonated amine, RNH3+ and bicarbonate. Moreover, the carbamate (RNHCOO-) concentration is very low, which is consistent with the behavior of the sterically hindered amines, especially when the hindered character is very important, as in the case of AHPD.3,8 The same type of concentration profiles are represented in Figure 7 for the quaternary system AHPD-Pz-CO2-H2O for an aqueous solution containing 1 kmol · m-3 AHPD and 0.3 kmol · m-3 Pz at 298.15 K. For the same reasons mentioned for the ternary system AHPD-CO2-H2O, the curves corresponding to H+ and OH- concentrations were not represented. It can be seen that the AHPD behavior in the presence of Pz is similar to that observed in the single aqueous amine (AHPD) system. The AHPD concentration decreases rapidly at CO2 loadings less than about 1 mol/mol of amine, while the bicarbonate and the protonated amine (RNH3+) increase sharply. At a CO2 loading of about 1 mol/mol of amine, practically all amine is converted preferentially into the protonated amine and bicarbonate and the AHPD carbamate (RNHCOO-) concentration is very low. On the contrary, Pz reacts very rapidly at very low CO2 loadings (up to about 0.2 mol/mol of total amine), and it is preferentially converted into Pz carbamate, PzCOO- .4,16 Pz dicarbamate, Pz(COO-)2, and diprotonated Pz (PzH22+) concentrations are very low. Moreover, with the increase of CO2 loading, Pz carbamate, PzCOO-, and protonated Pz, PzH+ are converting into protonated Pz carbamate, PzH+COO-. All 177 selected experimental data points for the system AHPD-CO2-H2O, covering a large range of amine concentrations (between 0.0125 and 4 mol · kg-1), temperature (between 283.15 and 333.15 K), and total pressure (between 1.85 and 2640.8 kPa) were correlated together with an average relative deviation of 22.7%. The interaction parameters for this system are valid for the entire range of temperature, pressure, and amine concentration (Table 3). Generally, higher deviations were obtained at very large amine concentration and very high pressures. Figures 8a and b show some comparisons between experimental and calculated total pressure at low and large amine concentrations. The solubility of carbon dioxide in aqueous solutions of mixed amine (AHPD + Pz) was predicted by supposing that the parameters characterizing the single amines systems are essential for describing the quaternary system behavior. Predictions of the CO2 partial pressure correspond to an average relative
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deviation of 37%. This is believed to mainly due to the fact that 49% of our quaternary experimental data are obtained at temperatures lower than 313 K, the lowest valid temperature of the interaction parameters available in the literature for the system Pz-CO2-H2O. For the same reason, no attempt was made to correlate the experimental data of the quaternary system. New experimental work is presently in progress in our laboratory in order to enlarge the experimental database for the Pz-CO2-H2O system at lower temperatures, which will allow us to revise and extend the available interaction parameters using the new data. 5. Conclusions In the present work, new data concerning the solubility of CO2 and N2O in aqueous mixtures of 2-amino-2-hydroxymethyl1,3-propanediol (AHPD) and piperazine (Pz) were obtained over a large range of temperature (283.15-333.15 K) and amine concentrations (0.91-4.36 mol · kg-1). On the basis of the experimental data, Henry’s law constants for CO2 in these solutions were calculated using the N2O analogy. The experimental data for the ternary system AHPD-CO2-H2O were satisfactorily correlated using a modified Pitzer’s thermodynamic model for the activity coefficients combined with the virial equation of state for representing the fugacity coefficients. The solubility of carbon dioxide in aqueous solutions of mixed amine (AHPD + Pz) was predicted by supposing that the available parameters characterizing the single amines systems are appropriate for describing the new data for the quaternary system behavior. However, the quite large deviations obtained between experimental and calculated equilibrium pressure led to the conclusion that more experimental data for the Pz-CO2-H2O system at lower temperatures are necessary in order to allow the revision of the interaction parameters for this system. Acknowledgment Financial support from the Natural Sciences and Engineering Research Council of Canada (NSERC) is gratefully acknowledged. List of Symbols ai ) activity of species i D ) dielectric constant or relative permittivity of pure water e ) electronic charge f ) function, as defined by eq 19 k ) Boltzmann’s constant KR ) equilibrium constant for the chemical reaction R, expressed on the molality scale m HCO (T,Pwsat) ) Henry’s constant for the solubility of carbon 2,H2O dioxide in pure water on the molality scale M ) molarity (kmol · m-3) Mw ) molar mass of water (kg · mol-1) mi ) true molality of species i in solution m ˜ i ) stoichiometric molality of component i NA ) Avogadro’s number P ) pressure q0, q1 ) coefficients, as defined in eq 23 pwsat ) saturated vapor pressure of water R ) universal gas constant T ) absolute temperature V(partial) ) molar volume Y ) vapor phase mole fraction zi ) charge of ion i Greek Letters R ) CO2 loading, mol CO2/mol amine
βij(0), βij(1) ) binary interaction parameters between species i and j in Pitzer’s equation ε0 ) permittivity of free space νi,R ) stoichiometric coefficient of component i in the reaction R φi ) fugacity coefficient of component i γi*,m ) activity coefficient of component i normalized to infinite dilution, on molality scale λij(I) ) second virial coefficient in Pitzer’s equation τijk ) ternary interaction parameter in Pitzer’s equation Fw ) water density (kg · m-3) Subscripts and Superscripts R ) reaction R ∞ ) infinite dilution in pure water m ) molality w ) water
Supporting Information Available: CO2 solubility in Pz-AHPD aqueous solutions (Table S1). This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited (1) Kohl, A. L.; Nielsen, R. B. Gas Purification, fifth ed.; Gulf Publishing Company: Houston, 1997. (2) Sartori, G.; Savage, D. W. Sterically Hindered Amines for CO2 Removal from Gases. Ind. Eng. Chem. Fundam. 1983, 22, 239. (3) Bougie, F.; Iliuta, M. C. Kinetics of absorption of carbon dioxide into aqueous solutions of 2-amino-2-hydroxymethyl-1,3-propanediol. Chem. Eng. Sci. 2009, 64, 153. (4) Bougie, F.; Lauzon-Gauthier, J.; Iliuta, M. C. Acceleration of the reaction of carbon dioxide into aqueous 2-amino-2-hydroxymethyl-1,3propanediol solutions by piperazine addition. Chem. Eng. Sci. 2009, 64, 2011. (5) Paul, S.; Ghoshal, A. K.; Mandal, B. Physicochemical Properties of Aqueous Solutions of 2-Amino-2-hydroxymethyl-1,3-propanediol. J. Chem. Eng. Data 2009, 54, 444. (6) Park, J.-Y.; Yoon, S. J.; Lee, H.; Yoon, J.-H.; Shim, J.-H.; Lee, J. K.; Min, B.-Y.; Eum, H.-M.; Kang, M. C. Solubility of carbon dioxide in aqueous solutions of 2-amino-2-ethyl-1,3-propanediol. Fluid Phase Equilib. 2002, 202, 359. (7) Park, J.-Y.; Yoon, S. J.; Lee, H.; Yoon, J.-H.; Shim, J.-G.; Lee, J. K.; Min, B.-Y.; Eum, H.-M. Density, Viscosity, and Solubility of CO2 in Aqueous Solutions of 2-Amino-2-hydroxymethyl-1,3-propanediol. J. Chem. Eng. Data 2002, 47, 970. (8) Park, J.-Y.; Yoon, S. J.; Lee, H. Effect of Steric Hindrance on Carbon Dioxide Absorption into New Amine Solutions: Thermodynamic and Spectroscopic Verification through Solubility and NMR Analysis. EnViron. Sci. Technol. 2003, 37, 1670. (9) Le Tourneux, D.; Iliuta, I.; Iliuta, M. C.; Fradette, S.; Larachi, F. Solubility of carbon dioxide in aqueous solutions of 2-amino-2-hydroxymethyl-1,3-propanediol. Fluid Phase Equilib. 2008, 268, 121. (10) Baek, J.-I.; Yoon, J.-H.; Eum, H.-M. Physical and Thermodynamical Properties of Aqueous 2-Amino-2-Methyl-1,3-Propanediol Solutions. Int. J. Thermophysics 2000, 21, 1175. (11) Baek, J.-I.; Yoon, J.-H. Solubility of Carbon Dioxide in Aqueous Solutions of 2-Amino-2-Methyl-1,3-propanediol. J. Chem. Eng. Data 1998, 43, 635. (12) Rumpf, B.; Maurer, G. An Experimental and Theoretical Investigation on the Solubility of carbon Dioxide in Aqueous Solutions of Strong Electrolytes. Ber. Bunsen-Ges. Phys. Chem. 1993, 97, 85. (13) Edwards, T. J; Maurer, G.; Newman, J.; Prausnitz, J. M. VaporLiquid Equilibria in Multicomponent Aqueous Solutions of Volatile Weak Electrolytes. AIChE J. 1978, 24, 966. (14) Perrin, D. D. Dissociation Constants of Organic Bases in Aqueous Solution; Butterworth and Co. Ltd.: London, 1966. (15) Hetzer, H. B.; Robinson, R. A.; Bates, R. G. Dissociation constants of piperazinium ion and related thermodynamic quantities from 0 to 50 deg. J. Phys. Chem. 1968, 72, 2081. (16) Ermatchkov, V.; Pe´rez-Salado Kamps, A.; Maurer, G. Chemical equilibrium constants for the formation of carbamates in (carbon dioxide + piperazine + water) from 1H-NMR-spectroscopy. J. Chem. Thermodyn. 2003, 35, 1277.
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(26) Pe´rez-Salado Kamps, A.; Xia, J.; Maurer, G. Solubility of CO2 in (H2O + Piperazine) and in (H2O + MDEA + Piperazine). AIChE J. 2003, 49, 2662. (27) Ermatchkov, V.; Pe´rez-Salado Kamps, A.; Speyer, D.; Maurer, G. Solubility of Carbon Dioxide in Aqueous Solutions of Piperazine in the Low Gas Loading Region. J. Chem. Eng. Data 2006, 51, 1788. (28) Munjal, P.; Stewart, P. B. Correlation Equation for Solubility of Carbon Dioxide in Water, Seawater, and Seawater Concentrates. J. Chem. Eng. Data 1971, 16, 170. (29) Lide, D. CRC Handbook of Chemistry and Physics, 89th ed.; CRC: Boca Raton, FL, 2008-2009. (30) Nova´k, J.; Fried, V.; Pick, J. Lo¨slichkeit des kohlendioxyds in wasser bei verschiedenen dru¨cken und temperature. Collect. Czech. Chem. Commun. 1961, 26, 2266. (31) Silkenba¨umer, D.; Rumpf, B.; Lichtenthaler, R. N. Solubility of Carbon Dioxide in Aqueous Solutions of 2-Amino-2-methyl-1-propanol and N-Methyldiethanolamine and Their Mixtures in The Temperature Range from 313 to 353 K and Pressures up to 2.7 MPa. Ind. Eng. Chem. Res. 1998, 37, 3133.
ReceiVed for reView May 1, 2009 ReVised manuscript receiVed October 29, 2009 Accepted November 2, 2009 IE900705Y
