Article pubs.acs.org/jced
Dissociation Constants (pKa) of Tertiary and Cyclic Amines: Structural and Temperature Dependences Aravind V. Rayer, Kazi Z. Sumon, Laila Jaffari, and Amr Henni* Acid Gas Removal Laboratory (AGRL), Faculty of Engineering and Applied Science, University of Regina, Regina, Saskatchewan S4S 0A2, Canada
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S Supporting Information *
ABSTRACT: The dissociation constants of the conjugate acids of 14 amines (diethylethanolamine, monoethanolamine, n-butyldiethanolamine, t-butyldiethanolamine, n,n-dimethylpropanolamine, methyl-diethanolamine, ethyldiethanolamine, monoethylethanolamine, n,n-dimethylisopropanolamine, triethanolamine, 4-methylpiperazine-1-amine, 3-morpholino propylamine, 4,2hydroxylethylmorpholine, and triethylamine) were measured over a temperature range between 293.15 and 333.15 K using the potentiometric titration method. The change in standard state thermodynamic properties was derived from the van’t Hoff equation. The influence of the steric hindrance, number of −OH groups, and length of alkyl chain on the dissociation constants was identified. Of the studied amines, few sterically hindered derivatives of piperazine, a secondary amine monoethylethanolamine, and a tertiary amine n,n-dimethylpropanolamine have high pKa values but lower standard enthalpy than those of the benchmark amine, monoethanolamine (MEA), and thus were deemed promising for CO2 capture technology. Monoethylethanolamine (MEEA) was found to have the highest basicity (pKa) with the lowest standard state enthalpy (ΔH°/kJ·mol−1).
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INTRODUCTION Aqueous amine solutions are important solvents for CO2 capture. CO2 is absorbed chemically, and upon absorption, carbamates, bicarbonates, and protonated amines are produced (for 1° and 2° amines). In a separate process operation, amines are regenerated at high temperature from carbamates or the protonated forms for reuse in a cyclic process. Although the carbamate formation of primary and secondary amines enhances the reaction rate of CO2 with amines, it is very difficult to regenerate the amine after the carbamate formation. Tertiary amines absorb CO2 only through bicarbonate formation with a slower reaction rate, a relative ease of amine regeneration. Therefore, insight into the protonation and deprotonation processes of amines in aqueous solution helps us understand the absorption and kinetic processes. Moreover, the corresponding equilibrium constants are helpful in nonrigorous screening of amines for developing new solvents.1,2 Finally, the dissociation constants are often of interest to analytical scientists and separation engineers when trying to identify the choice of techniques used to isolate compounds of interest.3 The dissociation constants for conjugate acids of amines, alkanolamines, amino acids, and cyclic amines were measured by various researchers.4−18 Sharma et al.19 reviewed the rate constants of many solvents used in CO2 capture studies, and suggested the existence of a correlation with the basicity of the amines. Versteeg et al.,20 in their classical review, correlated the second order rate constant for the formation of zwitterions and the dissociation constants of amines used for © 2014 American Chemical Society
CO2 capture studies, and proposed a correlation in the form of a Brönsted relationship. Penny and Ritter21 gave an explanation for the Brönsted relationship between the second order rate constant and the pKa of the amines. The pKa of less commonly studied amines (mainly tertiary and cyclic amines) are investigated in this work within the temperature range of (293.15 to 333.15) K. Thermodynamic quantities such as the standard enthalpy (ΔH°/kJ·mol−1) and entropy (ΔS°/kJ·mol−1) were calculated using the van’t Hoff equation. Trends in the variation of pKa values due to the addition of different radicals were analyzed for tertiary and cyclic amines using n-methyl diethanolamine (MDEA) and piperazine (PZ) as the base molecules.
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EXPERIMENTAL DETAILS Table 1 lists the molecular structures, abbreviations, and the purity of the solvents studied. All solvents used in this work were purchased from Sigma-Aldrich. The potentiometric titration method developed by Albert and Searjeant22 was used in this study. A pH meter, model 270 manufactured by Denver Instrument, was used to determine the pH values of the aqueous solutions. The pH meter electrode was calibrated with buffer solutions at each desired temperature. The buffer solutions with accuracies of (± 0.01) for pH 4.00 and 7.00, Received: July 20, 2014 Accepted: October 7, 2014 Published: October 20, 2014 3805
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Table 1. Molecular Structure of the Studied Amines for pKa Determination
and (± 0.02) for pH 10.00 were supplied by VWR International. The measured pH values of the buffer solutions at different temperatures are presented in Table 2. Hydrochloric acid solution (HCl) 0.100 M (± 0.002) was supplied by VWR International. Nitrogen gas with a high purity
(≥ 99.99 %) was purchased from Praxair for blanketing the solutions during the titration. A jacket beaker was used to keep the temperature constant during the titrations. Aqueous solutions of amines at 0.010 M (± 0.005) were prepared with dionized double distilled water. After the required temperature for titration was set, the pH meter was calibrated with buffer solutions. A slow stream of nitrogen was used to blanket the solution while the temperature of the solution was raised to the required set point. A 0.100 M aqueous solution of hydrochloric acid was used to titrate the amine solutions (50 mL). Twenty equal portions of the titrant were added to the solution, each portion being 0.5 mL. The pH value was recorded as soon as equilibrium was reached after the addition of the titrant. Details of the calculation of pKa values from the experimental pH values and calculation of the thermodynamic correction are given in appendix A (Supporting Information). The pKa values are calculated from the pH values measured using eq 1:
Table 2. Measured pH Values of the Calibration Buffersa pH T/K 293 298 303 308 313 323 333
buffer 1 4.00 4.00 4.01 4.02 4.03 4.06 4.09
± ± ± ± ± ± ±
0.00 0.00 0.02 0.03 0.03 0.03 0.03
buffer 2 7.02 7.00 6.99 6.98 6.97 6.97 6.98
± ± ± ± ± ± ±
0.03 0.02 0.01 0.04 0.03 0.03 0.03
buffer 3 10.06 10.00 9.94 9.90 9.85 9.78 9.70
± ± ± ± ± ± ±
0.02 0.01 0.03 0.02 0.03 0.03 0.03
a
Standard uncertainties u are u(pH) = 0.02; u(T) = 0.01 K (level of confidence = 0.95). 3806
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Table 3. Measured pKa Values for the Studied Amine at Different Temperaturesa
a
compound
T/K = 293
T/K = 298
T/K = 303
T/K = 308
T/K = 313
T/K = 323
T/K = 333
MDEA EDEA n-BDEA t-BDEA TEA TREA DMIPA DMPA DEEA MEEA MEA 4MPZ1A 3-MOPA 4,2-HEMO
8.74 8.90 8.99 9.11 7.87 10.74 9.53 9.56 9.82 9.90 9.56 7.45 10.03 6.99
8.54 8.80 8.90 9.03 7.73 10.64 9.47 9.51 9.73 9.74 9.45 7.40 9.95 6.90
8.49 8.69 8.74 8.89 7.68 10.58 9.41 9.38 9.64 9.68 9.27 7.32 9.70 6.87
8.41 8.59 8.67 8.83 7.55 10.50 9.19 9.35 9.56 9.55 9.14 7.27 9.55 6.79
8.32 8.54 8.60 8.75 7.42 10.40 9.15 9.25 9.45 9.53 9.06 7.24 9.22 6.66
8.25 8.42 8.36 8.52 7.28 10.15 9.04 9.05 9.29 9.34 8.72 7.17 9.11 6.59
8.08 8.26 8.09 8.43 7.05 9.85 8.72 8.92 9.08 9.13 8.58 7.08 8.92 6.47
Standard uncertainties are u(pKa) = 0.02; u(T) = 0.01 K (level of confidence = 0.95).
Figure 1. Dissociation constants of MEA at different temperatures.
pK a = pH + log
{BH+} {B}
I = 0.5 ∑ Cizi 2 i
(1)
where Ci is the molar concentration of an ion and z is its valency. At the end of the stepwise titration of the amine solution, pKa values were computed and the average value is reported in Table 3.
The activity was calculated using the following relation: {BH+} = [BH+]·γBH+
(2)
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where,γBH+ is the activity coefficient of the ionized species and was calculated from the Debye−Hückel equation: −log(γi) =
RESULTS AND DISCUSSION Comparison with Previous Literature Data. The dissociation constants for MEA, MDEA, TREA, and TEA were measured as a function of temperature and compared with values from the literature in Figures 1 to 4, respectively. The comparison of the dissociation constants (ln Ka) of MEA are plotted in Figure 1. It can be observed that values obtained in this work deviated slightly from those measured by Hamborg and Versteeg,13 Bates and Pinching,25 Kim et al.8 and Antelo et al.11 by 0.36 %, 0.72 %, 0.65 % and 1.47 % (percentage absolute average deviations, AAD %). The reason for the larger deviations with Antelo et al.11 was mentioned by Hamborg and Versteeg13 and relates to the consideration of activity
Azi2I1/2 1 + BkiI1/2
(4)
(3)
A and B are the Debye−Hückel equation constants, which are functions of the dielectric constants and temperature of the solvent. zi is the valency of the ion, and ki is the mean distance of approach of the ions (ionic size parameter). A and B constants were taken from Manov et al.,23 and the ionic size parameter (ki) values were taken from Kielland et al.24 I is the ionic strength which is a function of the concentration of the solution: 3807
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Figure 2. Dissociation constants of MDEA at different temperatures.
Figure 3. Dissociation constants of TREA at different temperatures.
coefficients as unity for all species. ln Ka values of MDEA measured in this work and the available literatures are presented in Figure 2. The values deviated from Hamborg and Versteeg,13 Kamps et al.,26 Oscarson et al.,16 Kim et al.,8 Littel et al.,9 Schwabe et al.3 by 1 %, 0.85 %, 2.2 %, 2.51 %, 0.49 %, and 1.46 %, respectively. The deviations may be due to the use of different instruments in the measurement of the protonation reaction. Hamborg and Versteeg13 and Kamps et al.26 used electromotive force (emf) measurements, Oscarson et al.16 used a flow calorimeter, and Kim et al.,8 Littel et al.,9 and Schwabe et al.6 used potentiometric titrations to measure the dissociation constants. The values obtained in this work are in better agreement with values reported by Littel et al.9 ln Ka values of TREA measured in this work deviated from Hamborg and Versteeg,13 Cox et al.,27 Fyfe et al.,28 Ablard et al.,29 and Bergstrom and Olofsson30 by 0.6 %, 0.61 %, 0.54 %, 0.39 %,
and 0.77 %, respectively. The deviations are within acceptable experimental deviations for TREA and are plotted in Figure 3. ln Ka values for TEA measured in this work deviated from those published by Hamborg and Versteeg,13 Bates and Allen,31 Bates and Schwarzenbach,32 Kim et al.,8 and Antelo et al.9 by 0.28 %, 0.28 %, 0.19 %, 0.9 %, and 1.41 %, respectively, and are shown in Figure 4. The measured pKa values of all the amines studied in this work are listed in Table 3. Substituent Effect and Comparison of Experiment Results with Prediction. Few trends appear in variation of pKa values due to the addition of various substituents. Figure 5 demonstrates the general trend of weakening pKa due to the addition of −OH groups at various locations on the molecule. The pKa of TREA, DEEA, EDEA, and TEA with zero, one, two, and three OH groups, respectively, are 10.67, 9.73, 8.8, 7.73, an approximate decrease of one pKa unit per OH group. 3808
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Figure 4. Dissociation constants of TEA at different temperatures.
Figure 5. Trend in pKa values of the amines with respect to the addition of −CH3 and −OH groups at 298.15 K.
The effect of −CH3 or −CH2−CH3 addition is less clear, but, when they are directly attached to the N atom rendering the amine from 1° to 2°, the pKa value decreased.33 Figure 6 is an
illustration of the change in pKa values of the cyclic amine family (piperazine, piperidine, and morpholine) due to the systematic modifications of the basic ring structure due to 3809
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Figure 6. Trend in pKa values of the cyclic amines with respect to the addition of −CH3 and −OH groups.
Table 4. Base Weakening Effect According to the Original Perrin−Dempsey−Serjeant (PDS) Group-Additivity Method base-weakening effects
error in ACD/prediction
amine
type
base value
ring
N-Me
β-OH
β-OR
γ-OH
PDS
expt
error
classic method
GALAS method
MEA MEEA DEEA n-BDEA t-BDEA DMPA MDEA EDEA DMIPA TEA 4,2-HEMO TREA rms error
1° 2° 3° 3° 3° 3° 3° 3° 3° 3° 3° 3°
10.77 11.15 10.5 10.5 10.5 10.5 10.5 10.5 10.5 10.5 10.5 10.5
0 0 0 0 0 0 0 0 0 0 0.2 0
0 0 0 0 0 −0.4 −0.2 0 0 0 0 0
−1.1 −1.1 −1.1 −2.2 −2.2 0 −2.2 −2.2 −1.1 −1.1 −3.5 0
0 0 0 0 0 0 0 0 0 −2.4 0 0
0 0 0 0 0 −0.44 0 0 0 0 0 0
9.67 10.05 9.4 8.3 8.3 9.66 8.1 8.3 9.4 7 7.2 10.5
9.45 9.74 9.73 8.9 9.03 9.51 8.54 8.8 9.47 7.73 6.9 10.64
−0.22 −0.31 0.33 0.6 0.73 −0.15 0.44 0.5 0.07 0.73 −0.3 0.14 0.43
0.29 0.04 −0.06 0 −0.06 0.24 −0.27 −0.1 0.04 −0.04 0.24 0.02 0.16
−0.05 −0.16 −0.17 −0.4 −0.57 0.31 −0.46 −0.5 0.67 −0.07 0.3 −0.06 0.37
Table 5. Base Weakening Effect According to the Updated Perrin−Dempsey−Serjeant (PDS) Group-Additivity Method base-weakening effects amine
type
base value
ring
N-Me
β-OH
β-OR
γ-OH
updated PDS
expt
error
MEA MEEA DEEA n-BDEA t-BDEA DMPA MDEA EDEA DMIPA TEA 4,2-HEMO TREA rms error
1° 2° 3° 3° 3° 3° 3° 3° 3° 3° 3° 3°
10.60 11.10 10.60 10.60 10.60 10.60 10.60 10.60 10.60 10.60 10.60 10.60
0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.20 0.00
0.00 0.00 0.00 0.00 0.00 −0.40 −0.20 0.00 0.00 0.00 0.00 0.00
−1.00 −1.00 −1.00 −2.00 −2.00 0.00 −2.00 −2.00 −1.00 −3.00 −1.00 0.00
0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.00 −2.80 0.00
0.00 0.00 0.00 0.00 0.00 −0.40 0.00 0.00 0.00 0.00 0.00 0.00
9.60 10.10 9.60 8.60 8.60 9.80 8.40 8.60 9.60 7.60 7.00 10.60
9.45 9.74 9.73 8.90 9.03 9.51 8.54 8.80 9.47 7.73 6.90 10.64
−0.15 −0.36 0.13 0.30 0.43 −0.29 0.14 0.20 −0.13 0.13 −0.10 0.04 0.23
substituent effects of all the groups to a fixed base value that depends on whether the amino moiety was primary, secondary or tertiary. The base-weakening (or strengthening) effects for ring-structure, n-methylation, and substituents like −OH and −OR in the beta or gamma position were considered. Both the
functionalization with alkyl, amino, or alcohol group. We have also analyzed the substituent effect using the group additivity method proposed by Perrin−Dempsey−Serjeant.34 The method considers an amine as a structural variation of aliphatic amines, and the pKa is predicted by linearly adding the 3810
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of the reaction center, and calculation of charge influences of the ionized groups to the neighboring ionization centers.36 Effect of Temperature. The enthalpy of protonation contributes significantly to the heat of absorption (regeneration) in CO2 capture technology using aqueous amine solutions.36 Indeed, for tertiary amines, this represents the bulk of the total heat of absorption. The standard state enthalpy change (ΔH°/kJ·mol−1) and entropy change (ΔS°/kJ·mol−1· K−1) were calculated using the van’t Hoff equation (see also ref 37), and are listed in Table 6.
Table 6. Standard Enthalpy and Entropy Changes of the Studied Aminesa −ΔrGm0 ‑1
solvent
pKa
kJ·mol
TREA 3-MOPA MEEA DEEA DMPA DMIPA MEA t-BDEA n-BDEA EDEA MDEA TEA 4MPZ1A 4,2-HEMO
10.67 9.95 9.74 9.73 9.51 9.47 9.45 9.03 8.90 8.80 8.54 7.73 7.40 6.90
61.75 56.49 56.69 56.43 55.30 54.86 54.03 52.18 51.14 51.04 49.88 45.17 43.19 40.04
ΔrHm0 kJ·mol
‑1
45.60 54.46 33.86 34.22 30.81 36.99 48.05 33.02 41.10 28.97 34.92 22.72 16.89 25.15
ΔSo kJ·mol‑1·K‑1
δ (%)b
−0.05 −0.01 −0.07 −0.07 −0.08 −0.06 −0.02 −0.06 −0.03 −0.07 −0.05 −0.07 −0.08 −0.05
0.52 0.31 0.27 0.13 0.32 0.52 0.33 0.24 0.33 0.22 0.58 0.23 0.17 0.30
ln K a =
C1 =
( −ΔH ◦) ΔS◦ + R RT
ΔS◦ R
and
C2 =
(5)
( −ΔH ◦) R
(6, 7)
In Figure 7, the dissociation constants (ln Ka) at different temperatures (1/T) for tertiary amines along with one secondary amine (2-ethylamino ethanol, MEEA) and one primary amine (MEA) were compared. The positive slope of the primary amine (MEA) was greater than that of the secondary amine (MEEA). The slopes of the tertiary amines were lower than that of both the primary and secondary amines, showing that the standard enthalpy changes for tertiary amines are small. Figure 8 shows the comparison of the dissociation constants for the cyclic amines studied in this work and those available in the literature (Khalili et al.3 and Xu et al.10). As can be seen in Table 6, amines with higher pKa tend to have a higher ΔH0, which is undesirable. However, from the kinetics point of view, there exists a linear Brønsted relationship between the second order rate constants (k2) and the dissociation constants (pKa). At a certain temperature, the rate increases with pKa.18 In Figure 9, some amines with relatively high pKa as well as low ΔH0 are identified. The linear amine MEEA, a tertiary amine such as DMEA, and sterically hindered cyclic amines such as 1-MPZ, 1-EPZ, and 4,2-HEPZ have the lowest standard enthalpies and the highest pKa values.
a
Standard uncertainties are u(pKa) = 0.02; u(T) = 0.01 K; u (−ΔrGm0) = 0.05; u(ΔrHm0) = 0.05; u(ΔS°) = 0.05 (level of confidence = 0.95). b δ % = ∑i n= 1|ln Ka,calc − ln Ka,exp/ln Ka,exp| × 100/n; n = number of experimental points.
original values and the recent updated values35 for the substituent effect were used, and the results for the amines for which all substituent effects were available are presented in Tables 4 and 5. The root-mean-square (rms) error with the original PDS and the updated PDS were 0.43 and 0.23, respectively. In Table 4, we included the prediction of pKa by the classic algorithm and also the global adjusted locally according to similarity (GALAS) algorithm commercially available in the ACD/Percepta software.36 It was found that the classic algorithm which is based on Hammet-type equations and electronic substituent constants to predict pKa values for ionizable groups gave good prediction compared to the GALAS method which is based on a multistep procedure involving the estimation of pKa microconstants for all possible ionization centers in the hypothetical state of an uncharged molecule with numerous corrections according to the chemical environment
Figure 7. Trend in dissociation constants (Ka) values at different temperatures for the studied amines. 3811
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Figure 8. Trend in dissociation constants (Ka) values at different temperatures for the cyclic amines.
Figure 9. Comparison of the studied amines at 298.15 K {amine, ΔH°/kJ·mol−1, pKa} for the feasibility2 in CO2 capture process.
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CONCLUSIONS
ASSOCIATED CONTENT
S Supporting Information *
The dissociation constants of 14 amines (primary, secondary, and tertiary) were measured at different temperatures (293.15 to 333.15) K. The standard enthalpies and entropies were calculated using the van’t Hoff equation. The effect of substituent groups on amine pKa was identified. The experimental results were compared with predictions by the group-additivity method and ACD software package. Some of the amines were identified to have high pKa, as well as low heat of absorption, two desirable criteria in solvent development.
Experimental pKa derivation; calculation of thermodynamic correction. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Tel: 306 585 4960; Fax: 306 585 4855; Email: amr.henni@ uregina.ca. 3812
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Funding
(19) Sharma, M. M. Kinetics of Reactions of Carbonyl Sulphide and Carbon Dioxide with Amines and Catalysis by Bronsted Bases of the Hydrolysis of COS. Trans. Farady Soc. 1965, 61, 681−687. (20) Versteeg, G. F.; van Dijck, L. A. J.; van Swaaij, W. P. M. On the Kinetics between CO2 and Alkanolamines both in Aqueous and Nonaqueous Solution, An Overview. Chem. Eng. Commun. 1996, 144, 113−158. (21) Penny, D. E.; Ritter, T. J. Kinetic Study of the Reaction between Carbon Dioxide and Primary Amines. J. Chem. Soc. Faraday Trans. I 1983, 79, 2103−2109. (22) Albert, A.; Serjeant, E. P. The Determination of Ionization Constants; A Laboratory Manual. 3rd ed. Chapman and Hall: London and New York, 1984. (23) Manov, G. G.; Bates, R. G.; Hamer, W. J.; Acree, S. F. Values of the Constants in the Debye-Hückel Equation for Activity Coefficients. J. Am. Chem. Soc. 1943, 65, 1765−1767. (24) Kielland, J. Individual Activity Coefficients of Ions in Aqueous Solutions. J. Am. Chem. Soc. 1937, 59, 1675−1678. (25) Bates, R. G.; Pinching, G. D. Acidic Dissociation Constant and Related Thermodynamic Quantities for Monoethanolammonium Ion in water form 0° to 50°C. J. Res. Natl. Bur. Stand. 1951, 46, 349−352. (26) Kamps, A. P.-S.; Maurer, G. Dissociation constant of NMethyldiethanolamine in Aqueous Solution at Temperatures from 278 to 368 K. J. Chem. Eng. Data 1996, 41, 1505−1513. (27) Cox, M. C.; Everett, D. H.; Landsman, D. A.; Munn, R. J. The Thermodynamics of the Acid Dissociation of Some Alkylammonium Ions in Water. J. Chem. Soc. (B) 1968, 1373−1379. (28) Fyfe, W. S. Complex-ion Formation. Part III. The Entropies of Reaction of the Silver and Hydrogen Ions in Some Aliphatic Amines. J. Chem. Soc. (B) 1955, 1347−1350. (29) Ablard, J. E.; McKinney, D. S.; Warner, J. C. The Conductance, Dissociation Constant and Heat of Dissociation of Triethylamine in Water. J. Am. Chem. Soc. 1940, 62, 2181−2183. (30) Bergström, S.; Olofsson, G. Thermodynamic Quantities for the Solution and Protonation of Four C6-Amines in Water over a Wide Temperature Range. J. Sol. Chem. 1975, 4, 535−555. (31) Bates, R. G.; Allen, G. F. Acidic Dissociation Constant and Related Thermodynamic Quantities for Triethanolammonium Ion in Water From 0°C to 50°C. J. Res. Natl. Bur. Stand. 1960, 64A, 343− 346. (32) Bates, R. G.; Schwarzenbach, G. Triät hanolamin als Puffersubstanz. Helv. Chim. Acta 1954, 37, 1437−1439. (33) Sartori, G.; Savage, D. W. Sterically Hindered Amines for CO2 Removal from Gases. Ind. Eng. Chem. Fundam. 1983, 22, 239−249. (34) Perrin, D. D.; Dempsey, B.; Serjeant, E. P. pKa Prediction for Organic Acids and Bases; Chapman and Hall: New York, 1981. (35) Sumon, K. Z.; Henni, A.; East, A. L. L. Predicting pKa of Amines for CO2 Capture: Computer versus Pencil-and-Paper. Ind. Eng. Chem. Res. 2012, 51 (37), 11924−11930. (36) ACD/Structure Elucidator, version 12.01; Advanced Chemistry Development, Inc.: Toronto, ON, Canada, 2014; www.acdlabs.com. (37) Kim, I.; Grimstvedt, A.; da Silva, E. F. Thermodynamics of Protonation of Alkanolamines in Aqueous Solutions. Energy Proc. 2011, 4, 576−582.
The financial support of the Natural Sciences and Engineering Research Council of Canada (NSERC), Natural Resources Canada (NRCan), Canadian Foundation for Innovation (CFI), Petroleum Technology Research Centre (PTRC) and the International Test Center for Carbon Dioxide Capture (ITC, University of Regina) are gratefully acknowledged. Notes
The authors declare no competing financial interest.
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REFERENCES
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