Article pubs.acs.org/IECR
Facile Method for Determination of Amine Speciation in CO2 Capture Solutions Naser S. Matin, Joseph E. Remias, James K. Neathery, and Kunlei Liu* Center for Applied Energy Research, University of Kentucky, Lexington, Kentucky 40511, United States S Supporting Information *
ABSTRACT: A simple and quantitatively reliable method for determination of amine speciation is introduced. The method employs three experimental methods that should be readily accessible. The results for CO2 loaded aqueous solutions of 30 wt % monoethanolamine (MEA) are used for demonstration of the method and show promising agreement with the more complicated spectroscopic methods reported in the literature. The measurements were done at ambient temperature and atmospheric pressure, since theoretical calculations and experimental data from the open literature revealed no considerable difference in speciation at different temperatures. The procedure is based on acid and base titration of the CO2 loaded amine solution along with the determination of total CO2 loading. The quantitative results for the different species concentration in the example MEA solution is in agreement with other available spectroscopic methods, mainly NMR, particularly for the free amine, carbamate, and protonated amine concentrations. Aspen Plus was also used for further assessment of the experimental data.
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necessary.9 They used the combination of proton and carbon NMR to study speciation in monoethanolamine (MEA)− H2O−CO2 and diethanolamine (DEA)−H2O−CO2 solutions at a wide CO2 loading range and temperatures between 20 and 80 °C. Of course, as noted, due to the fast proton transfer between molecular and protonated amine, only the sum of their concentrations can be observed. Furthermore, the NMR method is time-consuming and requires expensive equipment. Jakobsen et al. applied the same technique in the experimental part of their work to study species distribution in 2-[(2aminoethyl)amino]-ethanol (AEEA)−H2O−CO2 system.10 Fourier transform infrared spectroscopy (FTIR) was also used for the speciation study of CO2 absorbed into aqueous alkanolamines by Jackson et al.11 Their work was not quantitative, but they assigned the major peaks corresponding to a few important species in the liquid phase for MEA−CO2 and 2-amino-2-methyl-1-propanol (AMP)−CO2 system. On the other hand Derks et al. recently applied the FTIR technique to quantitative measurements of speciation in CO2 loaded aqueous solutions of N-methyldiethanolamine (MDEA).12 For performing the analysis, they initially prepared calibration curves for the species under study. Souchon et al. employed in situ Raman spectroscopy for the quantitative determination of the species distribution in alkanolamines (MEA, DEA, and MDEA)−H2O−CO2 systems at equilibrium at 40 °C.13 A gas chromatography technique also has been proposed.14 A titration method for the concentration measurement of different species in AMP−H2O−CO2 and DEA−H2O−CO2 was used by Haji-Sulaiaman et al.15 They employed a strong base (i.e., NaOH) to titrate the solution, followed by solving the equilibrium equations simultaneously to determine solution
INTRODUCTION Removal of acidic gases, particularly CO2, is a very important operation from an industrial and environmental point of view. Primary, secondary, and tertiary amines and their blends have found widespread application in the absorption and removal of carbon dioxide from process gases.1,2 In the designing of absorption or desorption columns for a CO2 capture process by amine or other reactive solvents, the precise knowledge of the reaction chemistry is crucial. In this regard, having an accurate determination of liquid phase species concentration is critical in CO2 absorption processes modeling vapor−liquid equilibria data. The contribution of the protonated amine, carbamate, and bicarbonate formed in amine solvents to, for example, the CO2 heat of absorption makes their quantities critical for thermodynamic understanding.3 Furthermore, providing a good estimation of the free amine concentration is important for mass transfer study of the CO2 absorption process, where the free amine concentration explicitly appears in the liquid film mass transfer coefficient.4 Finally, with good understanding of the above concepts it is possible to better design new solvents and blends. Various thermodynamic models, mainly based on activity coefficients, have been proposed for the CO2−H2O−alkanolamine systems.5−7 However, many of the models employ computational models, which are not readily accessible or require parameters not necessarily available for all solvents, making experimental methods of merit.5 Furthermore, reliable experimental values for species concentration in the liquid phase can improve the thermodynamic models. The activity coefficient models and their parameters can be better evaluated by the available amine speciation data.3 The NMR approach for speciation study of alkanolamine solutions containing dissolved CO2 is one of the basic methods for this purpose.8−10 As it was discussed by Böttinger et al., the quantification study in NMR spectroscopy is easier than optical spectroscopy or gas chromatography because no calibration is © 2012 American Chemical Society
Received: Revised: Accepted: Published: 6613
February April 20, April 27, April 27,
3, 2012 2012 2012 2012
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Article
0.1 M sulfuric acid solution was utilized here with sample size selected to require a minimum of 10 mL of titrant to minimize volumetric delivery error. For the base titrations, the dynamic equivalence titration (DET) method incorporated in the Metrohm Titrando 836 software was used as previously described.15 The method was validated by adding known amounts of potassium bicarbonate to 30 wt % MEA solutions and titrating using the method. A nominally 0.1 M sodium hydroxide solution was used for the base titrations with sample size selected to require a minimum of 10 mL of titrant. The base titrations were performed at ambient temperature. For the base titration, eq 1 can be used with M, n, v as the base molarity, number of hydroxide ions per base molecule, and volume of base (in mL) used for titration, respectively. (see Figure S.2, Supporting Information for typical DET titration curve)
speciation. Their approach is simple and reliable but still requires solving the simultaneous system of equations and requires the necessary thermodynamic parameters, which may not be available for all amines. In this work, a new method based on the total alkalinity measurements (acid titration), strong base titration, and CO2 loading of the solution was developed to calculate the different species concentrations in the MEA (monoethanolamine)− H2O−CO2 system. The method has several advantages in that it is fast, repeatable, and employs methods that are typically available in any laboratory studying CO 2 absorption. Furthermore, the approach does not require the knowledge of any thermodynamic parameters to determine speciation. The developed system is compared to experimentally determined values obtained using NMR for the same solution and to an Aspen Plus equilibrium model for MEA. Good agreement was found between the three methods.
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RESULTS AND DISCUSSION Background. Liquid bulk concentrations of chemical species and the partial pressure of solutes in the gas phase are required for all calculations including kinetic analysis and system simulation. Therefore, the independent chemical reactions, which produce or consume stable species in the bulk phase, and phase equilibrium equations have to be known for the system under study. Having such information makes it possible to calculate the desired quantities mentioned above. The following reactions may occur when CO2 absorbs into and reacts with aqueous primary and secondary amines.17 Ionization of water:
EXPERIMENTAL SECTION The aqueous solutions of 30 wt % MEA with different CO2 loadings were prepared and used for all experiments by contacting the aqueous MEA solution with a CO2 containing gas stream for different time periods. From each sample, five pieces of data were collected: pH, density, CO2 loading, total alkalinity, and base titration. The pH was measured at ambient temperature using a temperature corrected probe and 2 point calibration. The density was measured by dispensing 1 mL of solution (calibrated pipet to water) in quadruplicate and determining the mass. The CO2 solution loading was measured by an adaptation of a total inorganic carbon method previously described.16 In this method, phosphoric acid liberates the dissolved CO2 in amine solutions. The CO2 gas is stripped out with a nitrogen carrier gas and is directed to a HORIBA CO2 analyzer (VIA-510). The area under the curve is integrated, and CO2 concentration is determined using a calibration curve from a known analytical standard of potassium carbonate. Approximately 1 mL of sample (mass determined) was injected into the phosphoric acid reactor yielding CO2 concentration in mol CO2/kg sln. The HORIBA CO2 analyzer was calibrated with a certified CO2/N2 gas mixture (PurityPlus, 14.00% CO2) each day. The uncertainty was checked with a known analytical standard prior and after each set of unknowns with the allowable discrepancy in CO2 measurement (|(expected − measured)/expected| × 100) set at less than ±2% absolute. The standard deviation for repeated measurements was ±2.7%. Total alkalinity measurements were conducted using a Metrohm Titrando 836 and a standard equivalence point determination method, (with accuracy of ±0.003 in pH, and standard deviation at ±3.6% in end point measurements). Having the consumed acid volume at the final equivalence point (typically two were observed, see Figure S.1, Supporting Information) and its molarity, the following equation is used for alkalinity determination (and also base titration) of sample solution. A typical total alkalinity titration curve is shown in Figure S.1, Supporting Information. Total alkalinity (moleCO2 /kg sln) =
nMv wsample
2H 2O ↔ OH− + H3O+
(2)
Dissociation of dissolved CO2 through carbonic acid: CO2 + 2H 2O ↔ HCO3− + H3O+
(3)
Dissociation of bicarbonate: HCO3− + H 2O ↔ CO32 − + H3O+
(4)
Carbamate reversion to bicarbonate (hydrolysis reaction): RNHCOO− + H 2O ↔ RNH 2 + HCO−3
(5)
Dissociation of protonated amine: RNH3+ + H 2O ↔ RNH 2 + H3O+
(6)
Considering the above reactions, the potential species concentrations in question are OH−, H3O+, HCO3−, CO2, CO 3 − , RNH 3 + , RNH 2 (unreacted free amine), and RNHCOO−. One approach to finding the concentrations of the above-mentioned species is solving eight of the equations below simultaneously. The chemical reaction equilibriums, the mass and charge balances in the aqueous phase, and the phase equilibrium between CO2 in the gas and aqueous phase give enough (eqs 7−15) to perform the calculations.17 Amine balance: [RNH 2]f + [RNH3+] + [RNHCOO−] = [RNH 2]t
(7)
Carbon balance: (1)
[CO2 ] + [HCO3−] + [CO32 −] + [RNHCOO−]
Where M, n, v, and wsample are molarity of acid, number of proton per acid molecule, volume of acid was used during titration (mL), and sample mass (g), respectively. A nominal
= α[RNH 2]t
(8)
Charge balance: 6614
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Table 1. Equilibrium Constants18,19,a K1, K2, K3, K4, K5, a
(mol/dm3)2 (mol/dm3) (mol/dm3) (mol/dm3) (mol/dm3)
A
B
C
D
132.89888 231.465439 216.050446 −0.52135 −3.038325
−13445.9 −12092.1 −12431.7 −2545.53 −7008.357
−22.4773 −36.7816 −35.4819 0 0
0 0 0 0 −0.0031348
ln (K) = A + B/T + C ln(T) + DT, where T is in Kelvin.
CO2 + RNH 2 + H 2O ↔ H3O+ + RNHCOO−
[RNH3+] + [H3O+] = [HCO3−] + [OH−] + 2[CO32 −] + [RNHCOO−]
Therefore, it can be assumed that, during titration for total alkalinity, the free amine also includes all of the carbamate in the solution. Considering the pH for carbonated alkanolamine solutions, which are generally between 8 to 10, and the dissociation constant (pK2) of dissolved CO2 through carbonic acid at ambient temperature is ∼6.0 (see Table 1); it is a plausible approximation that the bicarbonate−carbon dioxide equilibrium, eq 3 is shifted toward the formation of bicarbonate. Thus, the concentration of free CO2 is generally very low.15 The alternative form of eq 8, considering carbon balance, gives
(9)
Independent equilibrium constants:
K1 = [OH−][H3O+]
(10)
K 2 = [HCO3−][H3O+]/[CO2 ]
(11)
K3 = [CO32 −][H3O + ]/[HCO3−]
(12)
K4 = [RNH 2][HCO3−]/[RNHCOO−]
(13)
K 5 = [RNH 2][H3O + ]/[RNH3+]
(14)
PCO2 = He × [CO2 ]
[CO2 ]t = [CO2 ]f + [HCO3−] + [CO32 −] + [RNHCOO−]
Alka t = {[RNH 2]f + [RNHCOO−]} + 2[CO2 ]t
Where [RNH2]f, [RNH2]t, α, and He are unreacted free amine and total amine concentrations, carbon loading, and Henry’s constant. The equilibrium constants are given in Table 1. Using eqs 2−15 is the approach applied to thermodynamic modeling of CO2 loaded solutions (alkanolamines) to calculate each species concentration in the liquid and gas phase, and it is necessary to solve and optimize the system of nonlinear equations over the experimental solubility data. In the next section, a simple and reliable titration method is introduced to obtain the speciation data in alkanolamine solutions based on the mass and charge balance without the need for equilibrium constants. Method Description. For total alkalinity measurement, the reactions 3, 4, and 6 can take place with approximate pK values of 7, 11, and 10, respectively (see Table 1). Therefore, considering the pK of the mentioned reactions, with addition of proton (acid) the bicarbonate ion and protonated amine formation take place at approximately pH ≅ 6 (pH of 7.3 in experiments), that is, the first equivalence point. Further addition of acid results in the bicarbonate ion converting to CO2 through eq 3 at pH ≅ 4 (pH of 4.4 in experiments), that is, the second equivalence point. Depending on the amine protonation equilibrium constant of various alkanolamines, the pH versus acid volume graph might display three equivalence points, that is, a large difference in equilibrium constants of eqs 4 and 6, but the pattern and interpretation remain the same. The total alkalinity of the solution can be calculated using an acid titration, as follows:
− [HCO3−] − [RNHCOO−]
(16-a)
which is simplified as follows, Alka t = [RNH 2]f + 2[CO2 ]t − [HCO3−]
(16-b)
Considering eq 4, at basic conditions, the following equilibrium reaction can be considered between bicarbonate and carbonate ions CO32 − + H 2O ↔ HCO3− + OH−
(4-a)
Using the dissociation constant for bicarbonate to carbonate ion and proton at room temperature (pK3 ≅ 10.3), and the water dissociation constant (pKw ≅14), the equilibrium constant for eq 4-a can be expressed as K3‐a = K w /K3 = [HCO3−][OH−]/[CO32 −] ≅ 10−3 (18)
Considering eq 18, the basicity determines the bicarbonate/ carbonate ratio in the solution. Therefore, it can be inferred at the mentioned solution’s pH (8 ≤ pH ≤ 10) that 10 ≤ [HCO3−]/[CO32 −] ≤ 1000
So, for most CO2 loading ranges considered and particularly at the loadings above 0.25 molCO2/molamine, the concentration of CO32− compared to HCO3− is very low, and it can be discarded from the eq 8 or 8-a. Therefore, the simplified form of the eq 8a is
Alka t = [RNH 2]f + [HCO3−] + 2[CO32 −] + [RNHCOO ]
(8-a)
Therefore, from eq 16, 8-a, and discussion followed by eq 17, the following equation can be reached,
(15)
−
(17)
[CO2 ]t = [HCO3−] + [RNHCOO−] (16)
(8-b)
where [CO2]t was determined experimentally using the abovedescribed method. Moving from the alkalinity measurements, as the acid titration gives the free amine concentration, the dynamic
The addition of the acid (proton) in alkalinity measurement causes protonation of the carbamate to produce free amine and CO2 through the reverse reaction 6615
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agreement with the Aspen Plus predicted values. The pH is defined as the activity of the proton in the solution. On the other hand, the single ion activity coefficients are related to the activity coefficient of the neutral electrolyte (i.e., γ+γ− = γN2).21 Therefore, as it can be deduced from the fundamental equations for ionic species activity in the electrolyte systems, and as it has also been mentioned in the literature that the single ion activity cannot be measured experimentally using pH and compared to predicted activity coefficient models for high ionic strength solutions.22 However, Chan et al. displayed that the pH measurement using a glass electrode even at high ionic strength comparing to the Pitzer model for single ion activity coefficient of hydrogen ion is in the acceptable range of error (generally less than 1.4%).23 Speciation Comparisons. Figure 2 displays the speciation measured by the method developed in this work compared to
titration of the amine solution with a strong base yields the concentration of bicarbonate, protonated amine and free CO2,15 so that the following balanced equation can be written [Base] = [RNH3+] + [HCO3−] + [CO2 ]f
(19)
with the same reasoning for eqs 16 to 16-a and 8-a to 8-b, [Base] = [RNH3+] + [HCO3−]
(19-a)
Considering the charge balance equation (eq 9) and the assumption that the concentrations of H3O+, OH−, and CO32− species are negligible compared to other species (i.e., pH ≅ 9), then the charge balance can be written as follows, [RNH3+] = [HCO3−] + [RNHCOO−]
(9-a)
Then, from eq 8-b and the simplified charge balance (eq 9-a), the following useful equation can be derived: [Base] = [CO2 ]t + [HCO3−]
(19-b)
Therefore, titration of a sample solution with strong base (e.g., NaOH) and using eq 19-b gives bicarbonate concentration. Then, applying eqs 19-a and 8-b the protonated amine and carbamate concentrations can be calculated, respectively. Using the above, along with the alkalinity measurements (acid titration), total inorganic carbon content of the solution and using eq 16-b the free amine concentration can be determined. Table S.1 (Supporting Information) shows the experimental values for the concentration of the different species in MEA solution with different CO2 loading. As it can be seen from Table S.1 (Supporting Information), in terms of the carbamate, protonated amine, and free amine, the agreement is good. In the case of bicarbonate concentration, the values determined in this work generally give higher values compared to the work done by Jakobsen et al., which is discussed later in this section.8 Table S.2 (Supporting Information) displays the uncertainty in this work compared to literature values. Aspen Model Verification. For further verification of the experimental method, Aspen Plus with the ENRTL-RK activity coefficient model was chosen to compare the results. Figure 1
Figure 2. MEA speciation as a function of carbon loading: this work at 21 °C, blank symbols; Jakobsen et al.,8 filled symbols; and Aspen Plus, dotted and filled lines at 20 °C.
the Aspen Plus prediction and literature.8 The agreement between data produced in this work using total carbon and HCO3− analytical measurement with the data from Jakobsen et al.8 and the Aspen Plus prediction are fairly good. As can be seen at the intermediate CO2 loading (i.e., 0.3 to 0.45) the predicted values for the total bicarbonate and carbonate concentration reveals slightly different values. At increased CO2 loading, 0.45 molCO2/molMEA, the predicted values for the three different sources come closer. The bicarbonate concentration determined in this work is higher than the data from Jakobsen et al.,8 particularly in the range of 0.3 to 0.45 CO2 loading. However, as it is discussed by Jakobsen et al.,8 their data shows the average uncertainties of 10.1% and 11.7% and the maximum uncertainties were 29.2% and 34.4% for data at 20 and 40 °C, respectively, in electroneutrality. They also pointed out that because of the sensitivity of their results to the calibration method used, for the species with fractions
