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Dissociation Constant (pKa) and Thermodynamic Properties of Some Tertiary and Cyclic Amines from (298 to 333) K Ali Tagiuri, Mohanned Mohamedali, and Amr Henni* Clean Energy Technology Research Institute (CETRi), Faculty of Engineering and Applied Science, University of Regina, Regina, Saskatchewan S4S 0A2, Canada S Supporting Information *

ABSTRACT: The potentiometric titration method was used to measure pKa values of nine amines [2-(diisopropylamino)ethanol [2-DIPA], N,N,N′,N′-tetrakis(2-hydroxypropyl)ethylenediamine [THPEDA], 2-{[2-(dimethylamino) ethyl]methyl amino} ethanol [2-DMAEMAE], tris[2-(2-methoxyethoxy)ethyl]amine [TMEEA], N-(2-hydroxyethyl) aniline [2-HEAN], 1-(2hydroxyethyl)piperazine [HEP], piperazine [PZ], monoethanolamine [MEA], and N-methyldiethanolamine [MDEA] at (298.15, 303.15, 313.15, 323.15, and 333.15) K. pKa values of the last three amines were compared with published data to validate the procedure and assess the accuracy of the instrument. Thermodynamic quantities, such as standard enthalpy (ΔH°·kJ−1·mol−1) and entropy (ΔSo·kJ−1·mol−1) for the dissociation process, were determined at 298.15 K using the van’t Hoff equation. From the experimental results, the values of the standard state thermodynamic properties were derived and compared to the values of commercially available amines used as absorbents for CO2 capture. Among the studied amines, 2-DMAEMAE was identified as having a high pKa (9.18) and lower heat of dissociation than MEA (27.77 kJ/mol, as compared to 48.59 kJ/mol for MEA) and can therefore be considered a potential candidate for CO2 capture applications.

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INTRODUCTION Aqueous amine solutions are widely used in the removal of acid gases from various gas mixtures, including in CO2 capture from flue gases. Monoethanolamine (MEA) is considered the most widely used amine in postcombustion capture applications.1 CO2 is primarily absorbed by a chemical reaction that releases carbamates, bicarbonates, or protonated amines depending on the nature of the amine.2 Studying the protonation process of amines in aqueous solutions is of great importance to understand the overall reaction as well as the kinetic aspect of the CO2 capture process.3 Moreover, determination of the equilibrium constant of the CO2-amine reaction is essential to understand the influence of operating conditions (temperature and pressure) on CO2 removal. Therefore, an insight into the pKa values of different alkanolamine systems helps evaluate the CO2 absorption and regeneration processes and is useful in measuring the proton acceptor characterizing the amines. Several researchers have studied the pKa for different amino acids, alkanolamines, and cyclic amine absorbents.4,5 The dissociation constants for the most commonly used alkanolamines in CO2 capture applications were reported by Perrin.6 Sharma and co-workers7 have studied the reaction kinetics of CO2-amine systems and concluded the existence of a strong relationship between the rate constant and the basicity of the amine, a statement supported by a similar study published by Versteeg et al.8 and others. © XXXX American Chemical Society

pKa is a core property of any electrolyte and defines its biological and chemical behavior as mentioned by Khalili et al.9 and Hamborg et al.10 Potentiometric titration is the most common and convenient method for pKa determination. This method is the fastest experimental technique to measure the pKa of a compound and can provide accurate results.11 In this study, the potentiometric titration method was used to determine the pKa values of amines in the range of 3.48 to 9.75. The pKa values of the last five tertiary amines, listed in Table 1, have not been reported in the literature and are determined at temperatures between (298.15 and 333.15) K.

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CHEMICALS AND APPARATUS All nine amines were purchased from Sigma-Aldrich. Table 1 lists the molecular structures of the solvents used along with their purities. A pH meter, model 270 Denver Instrument, was used to determine the pH values of the aqueous solutions. The high accuracy pH meter electrode was calibrated at each temperature using buffer solutions. VWR International supplied the buffer solutions with an accuracy of (± 0.01) for pH of 4.01 and 7.00, and (± 0.02) for pH 10.01. Measured pH values Received: June 19, 2015 Accepted: October 20, 2015

A

DOI: 10.1021/acs.jced.5b00517 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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Table 1. Molecular Structures of the Studied Amines for pKa Determination

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Table 2. Measured pH Values of the Calibration Buffersa

DETERMINATION OF PKA VALUES The procedure described above was used to calculate pKa values of [MEA] and all other amines. [MEA] is an acidic base and in aqueous solution can ionize along the following process

pH T/K 298.15 303.15 313.15 323.15 333.15 a

buffer 1 4.00 4.01 4.02 4.04 4.06

± ± ± ± ±

0.00 0.01 0.02 0.03 0.03

buffer 2 7.00 6.99 6.98 6.97 6.98

± ± ± ± ±

0.01 0.01 0.02 0.03 0.03

buffer 3 10.00 9.99 9.96 9.89 9.80

± ± ± ± ±

0.01 0.03 0.02 0.03 0.03

Ka

MEAH+ + H 2O ↔ MEA + H3O+

(1)

where {MEAH+} represents the protonated MEA, and {MEA} is the monoethanolamine molecule. Considering the mole fraction of water as one, and expressing the concentration of hydronium ion [H3O+] as [H+], the pKa equation of amine is as shown

Standard uncertainties: u(pH) = 0.03 and u(T) = 0.01 K.

of the buffer solutions, at different temperatures, are given in Table 2. Hydrochloric acid solution (HCl, 0.100 M (± 0.002)) was purchased from VWR International. High purity nitrogen gas (≥ 99.99 %) was provided by PRAXAIR for blanketing the solutions during the titration. A jacket beaker was employed to keep the temperature constant during the titration and was connected to an external water bath. Aqueous solutions of amines at 0.010 M (± 0.005) were prepared using deionized double distilled water. The pH meter was calibrated using buffer solutions after setting the required temperature for titration. The temperature was raised and a slow stream of gas nitrogen was used as a blanket until the solution reached the required temperature. For titration of the amine solutions (50 mL), a 0.100 M aqueous solution of HCl was used. Twenty portions of the titrant were added to the aqueous solution with each part being 0.5 mL. pH values were recorded as soon as equilibrium was reached after the addition of the titrant, as reported by Albert and Serjeant et al.11

pK a = pH + log

{MEAH+} {MEA}

(2)

pKa values are calculated using eq 2. These approximate equations are valid for pH between 4 and 10 as reported by Albert and Serjeant.11 The calculation procedure for the thermodynamic correction factor is explained in the Supporting Information. The ionic strength (I) can be calculated as I = 0.5 ∑ Cizi 2

(3)

where Ci is the molecular concentration of an ion and zi is the valency At the end point, nine pKa values (addition of 5 mL of 0.100 M HCl to neutralize 5 mL of amine solution) using eq 2, and the average values of pKa for [MEA], are shown in Table 3. The first and second pKa values, determined for all studied amines, are listed in Tables 4 and 5. B



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Table 3. pKa Values of [MEA] at 298.15 K HCL/ml

pH

[MEAH+]/[MEA]

log([MEAH+]/[MEA])

thermodynamic correction

pKa

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0

10.78 10.30 9.97 9.80 9.55 9.42 9.30 9.19 9.09 8.87 8.76

0.11 0.25 0.43 0.67 1.00 1.50 2.33 4.00 9.00

−0.95 −0.60 −0.37 −0.18 0.00 0.18 0.37 0.60 0.95

0.02 0.02 0.03 0.03 0.03 0.03 0.04 0.04 0.04

9.42 9.49 9.49 9.48 9.48 9.47 9.42 9.34 9.00

average

pKa

9.47 ± 0.02

Table 4. Measured First pKa Values for the Studied Amines at Different Temperaturesa T/K compound

298.15

[MEA] [MDEA] [PZ] [HEP] [2-DIPA] [THPEDA] [2-DMAEMAE] [TMEEA] [2-HEAN] a

9.47 8.57 9.75 9.10 9.42 7.72 9.18 6.91 8.66

± ± ± ± ± ± ± ± ±

0.02 0.03 0.02 0.04 0.03 0.02 0.03 0.01 0.02

303.15 9.26 8.50 9.67 8.95 9.26 7.57 9.05 6.82 8.54

± ± ± ± ± ± ± ± ±

313.15

0.01 0.01 0.02 0.03 0.02 0.03 0.01 0.03 0.03

9.09 8.31 9.38 8.74 9.11 7.38 8.92 6.69 8.39

± ± ± ± ± ± ± ± ±

0.03 0.02 0.04 0.05 0.03 0.01 0.03 0.04 0.01

323.15 8.74 8.24 9.16 8.65 8.77 7.22 8.78 6.55 8.17

± ± ± ± ± ± ± ± ±

333.15

0.02 0.02 0.02 0.03 0.02 0.03 0.02 0.03 0.02

8.57 8.07 9.07 8.53 8.56 7.10 8.65 6.49 8.03

± ± ± ± ± ± ± ± ±

0.01 0.01 0.03 0.02 0.01 0.02 0.04 0.02 0.03

Standard uncertainties: u(pKa) = 0.03 and u(T) = 0.01 K.

Table 5. Measured Second pKa Values for the Studied Amines at Different Temperaturesa T/K compound

298.15

[PZ] [HEP] [THPEDA] [2-DMAEMAE] a

5.36 3.91 4.34 4.21

± ± ± ±

0.03 0.06 0.02 0.02

303.15 5.26 3.88 4.19 4.01

± ± ± ±

313.15

0.04 0.04 0.04 0.03

5.01 3.74 3.99 3.82

± ± ± ±

0.02 0.03 0.03 0.01

323.15 4.94 3.59 3.85 3.73

± ± ± ±

333.15

0.03 0.05 0.02 0.04

4.83 3.48 3.73 3.51

± ± ± ±

0.02 0.03 0.03 0.03

Standard uncertainties: u(pKa) = 0.03 and u(T) = 0.01 K.

amount of titrant (HCl) added by buret to a 50 mL solution of 0.010 M [MEA]. The pH values corresponding to each titrant added are shown in the second column, which was read from the pH meter. The ([MEAH+]/[MEA]) ratio of the first 10 rows is shown in the third column; after the addition of titrant (0.5 mL of 0.100 M HCl), one tenth of the base [MEA] would be protonated, and all the [MEA] molecules would be protonated [MEAH+] after the addition of 5 mL (0.100 M HCl). Thus, the solution would be [MEAH+] as the only ionized species. This point is called the first end point. A quantity of hydroxyl ion (OH−) would be produced as well due to the dissociation of water. However, this quantity is negligible compared to the concentration of (H+). The thermodynamic correction is shown in the fifth column, which includes the effect of activity coefficients, and these terms are presented in the Supporting Information. At the first end point, using eq 2, nine pKa values can be calculated, and the average value was reported as the pKa of [MEA] as shown in Table 3. First, pKa values were then determined in a similar manner for all amines studied and are listed in Table 4. The second pKa values for the diamines used in this work are listed in Table 5.

The following relation is used to calculate the activities +

+

{BH } = [BH ]. γBH+

(4)

where γBH+ is the activity coefficient of the ionized species and is usually smaller than one. To calculate the activity coefficients, the Debye−Hückel12 equation was used −log(γi) =

Azi 2 I1/2 1 + BaiI1/2

(5)

where the A and B terms are the Debye−Hückel equation constants, which vary with the dielectric constant and temperature of the solvent. The term ai is the ionic size parameter, i.e., the mean distance of approach of the ions. (zi) is the ion valence, and I is the ionic strength, which is a function of the concentration of the solution. The ionic size parameters (ai) were taken from Kielland et al.13 and A and B from Manov et al.14

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RESULTS AND DISCUSSION As shown in Table 3, pKa values were determined from pH values for [MEA] at 298.15 K. The first column is the C



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from those published by Hamborg and Versteeg,15 Albert and Serjeant et al.,11 Kim et al.,16 and Bates and Pinching et al.17 by (1.11, 1.25, 0.40, and 1.30) %, respectively. ln Ka values for [MDEA] were also measured, and the comparisons with available data from the literature are presented in Figure 2.

Table 6. Thermodynamic Quantities for the First Dissociation of Amines in Aqueous Solutiona solvent [MEA] [MDEA] [PZ] [HEP] [2-DIPA] [THPEDA] [2-DMAEMAE] [TMEEA] [2-HEAN]

ΔrGm°

ΔrHm°

ΔS°

pKa

kJ·mol‑1

kJ·mol‑1

kJ·mol‑1·K‑1

± ± ± ± ± ± ± ± ±

53.89 48.88 55.55 51.65 53.78 43.85 52.25 39.33 49.38

48.59 26.49 38.68 29.94 46.50 33.19 27.77 23.25 34.24

−0.02 −0.07 −0.06 −0.07 −0.02 −0.04 −0.08 −0.05 −0.05

9.47 8.57 9.75 9.10 9.42 7.72 9.18 6.91 8.66

0.02 0.03 0.02 0.04 0.03 0.02 0.03 0.01 0.02

Standard uncertainties: u(pKa) = 0.03; u(T) = 0.01 K.; u(ΔrG°m) = 0.05; u(ΔrHm° ) = 0.05; (ΔS°) = 0.05.

a

Table 7. Thermodynamic Quantities for the Second Dissociation of Amines in Aqueous Solutiona solvent [PZ] [HEP] [THPEDA] [2-DMAEMAE]

Figure 2. Dissociation constants of MDEA at different temperatures.

ΔrGm°

ΔrHm°

ΔS°

pKa

kJ·mol‑1

kJ·mol‑1

kJ·mol‑1·K‑1

± ± ± ±

30.41 22.43 24.54 23.71

28.86 24.39 32.54 34.87

0.003 0.007 0.027 0.037

5.36 3.91 4.34 4.21

0.03 0.06 0.02 0.02

The values deviated from Hamborg and Versteeg,15 Kim et al.,16 Littel et al.,18 and Schwabe et al.19 by (1.03, 0.59, 0.54 and 1.90) %, respectively. The deviations are generally due to the use of different instruments and the procedures in the calculation of the protonation reaction. Hamborg and Versteeg15 used electromotive force measurements; Kim et al.16 and Littel et al.18 used potentiometric titration to obtain the dissociation constants. Finally, measured ln Ka values for [PZ] and data available in the literature are presented in Figure 3.

Standard uncertainties: u(pKa) = 0.03; u(T) = 0.01 K.; u(ΔrGm° ) = 0.05; u(ΔrH°m) = 0.05; (ΔS°) = 0.05.

a

Figure 1. Dissociation constants of MEA at different temperatures. Figure 3. Dissociation constants of PZ at different temperatures.

The dissociation constants were calculated and correlated with temperature using the following relationship using the measured pKa values. ln K a = −ln(10 pKa) = C1 +

C2 T /K

Measured values deviated from those published by Khalili et al.,9 Hetzer et al.,20,21 Hamborg and Versteeg,15 Pagano et al.,22 and Enea et al.23 by (0.77, 0.64, 0.43, 1.85, and 2.50) %, respectively. All deviations are within the acceptable experimental deviations for the three amines (MEA, MDEA, and PZ).

(6)

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Using the van’t Hoff equation, the standard state enthalpy change (ΔH°/kJ·mol−1) and entropy change (ΔS°/kJ·mol−1·K−1) were calculated and are the values are listed in the Tables 6 and 7. ln K a =

( −ΔH °) ΔS° + R RT

EFFECT OF SUBSTITUENT GROUPS ON THE PKA VALUES The variation of the pKa values in response to the effect of the substituent groups added are shown in Figure 4. The base weakening effect of the (−OH) group was confirmed, as it was observed that pKa values decrease significantly upon the insertion of (−OH) groups to the molecular structure of the amines. For instance, the pKa falls by approximately 1.7 units due to the addition of two (−OH) groups if 2-DIPA and

(7)

( −ΔH °) ΔS° ; C2 = (8) R R A comparison of the dissociation constants (ln Ka) of [MEA] is shown in Figure 1. The values of (ln Ka) deviated C1 =

D



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Figure 4. Effect of (−CH3) and (−OH) groups on the experimental pKa values at 298.15 K.

Table 8. Base-Weakening Effect According to Group Additivity Method (PDS) at 298.15 K base-weakening effect

a

amine

type

base value

ring

N-CH3

β-OH

β-OR

β-NHR

α-C6H5

β-NR2

PDS

exp.

error

[MEA] [MDEA] [PZ]a [HEP]a [2-DIPA] [THPEDA]a [2-DMAEMAE]a [TMEEA] [2-HEAN]

1° 3° 2° 2° 3° 3° 3° 3° 2°

10.77 10.50 11.15 11.15 10.50 10.50 10.50 10.50 11.15

0 0 0.2 0.2 0 0 0 0 0

0 −0.2 0 0 0 0 −0.4 0 0

−1.1 −2.2 0 −1.1 −1.1 −2.2 0 0 −1.1

0 0 0 0 0 0 0 −3.6 0

0 0 −1.8 0 0 0 0 0 0

0 0 0 0 0 0 0 0 −1.4

0 0 0 −1.8 0 −0.9 −0.9 0 0

9.67 8.10 9.85 8.75 9.40 7.70 9.50 6.90 8.65

9.47 8.57 9.75 9.10 9.42 7.72 9.18 6.91 8.66

−0.20 0.47 −0.10 0.35 0.02 0.02 −0.32 0.01 0.01

Statistical factor (log 2 = 0.3) was added to account for the number of sites available for proton addition.

estimate the pKa of a compound using a base value for the parent amine and adding an incremental pKa value corresponding to each substituent group incorporated into the amine structure. The base values of the parent amines are fixed depending on the nature of the amino moiety (primary, secondary, or tertiary). The influence of the incorporated substituent groups could be either a base-weakening or -strengthening effect. In this work, we have considered the effect of OH, OR, NHR, and NR2 attached to the β position in addition to the effect of methyl groups attached to the nitrogen atom. Table 8 shows both the experimental pKa values along with the PDS predictions and the error in prediction. Maximum errors of 0.47 and −0.45 were recorded for MDEA and HEP systems, respectively. The updated PDS method, reported by our group (Sumon et al.25), was also used to predict the pKa values. Both methods were found to provide satisfactory results

THPEDA are considered, which is in agreement with the findings of a previous study conducted by Rayer et al.24 Conversely, the influence of the addition of a (−CH3) group was less obvious; however, the base weakening effect of the (−CH3) group was observed when they are directly bonded to the nitrogen atom, as can be seen in 2-DMAEMAE. In addition, it can be clearly observed that the replacement of a hydrogen atom in MEA with a benzyl group decreased the pKa by one unit, which could be attributed to the hindrance effects of the bulky benzyl group.

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PREDICTION OF PKA VALUES USING GROUP ADDITIVITY METHOD The Perrin−Dempsey−Serjeant empirical method (PDS)26 was used to predict pKa values for the different amines used in this study. The method uses a linear free energy relationship to E



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with an acceptable range of errors. The original method was slightly better than the updated database with average absolute errors of 0.17 and 0.20, respectively.

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EFFECT OF TEMPERATURE

In CO2 capture applications with liquid absorbents, a major challenge is the high energy requirements for the regeneration process. The heat of absorption/regeneration is a very important criterion when evaluating the performance of different aqueous amine systems. The enthalpy of protonation represents a considerable fraction of the total heat, especially for tertiary amines, as it constitutes the largest share in the total heat of absorption. In Figure 5, the dissociation constants are

Figure 5. Trend in dissociation constant (pKa) values at different temperatures for the studied amines.

Table 9. Base-Weakening Effect According to Updated Group Additivity Method (PDS) at 298.15 K base-weakening effect amine

type

base value

ring

N-CH3

β-OH

β-OR

β-NHR

α-C6H5

β-NR2

PDS

exp.

error

[MEA] [MDEA] [PZ]* [HEP]* [2-DIPA] [THPEDA]* [2-DMAEMAE]* [TMEEA] [2-HEAN]

1° 3° 2° 2° 3° 3° 3° 3° 2°

10.60 10.60 11.10 11.10 10.60 10.60 10.60 10.60 11.10

0 0 0 0 0 0 0 0 0

0 −0.2 0 0 0 0 −0.4 0 0

−1.0 −2.0 0 −1.0 −1.0 −2.0 0 0 −1.0

0 0 0 0 0 0 0 −4.2 0

0 0 −2.0 0 0 0 0 0 0

0 0 0 0 0 0 0 0 −1.4

0 0 0 −2.0 0 −1.0 −1.0 0 0

9.60 8.40 9.40 8.40 9.60 7.90 9.50 6.40 8.70

9.47 8.57 9.75 9.10 9.42 7.72 9.18 6.91 8.66

−0.13 0.17 0.30 0.70 −0.18 −0.18 −0.32 0.51 −0.04

*

Statistical factor (log2 = 0.3) was added to count for the number of sites available for proton addition.

Figure 6. Standard enthalpy and pKa values of different amines studied at 298.15 K for applications in the CO2 capture process; the red rectangle represents the favored region for efficient CO2 capture applications. F



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presented at different temperatures for all amines studied. It is evident that MEA had the highest absolute slope value as compared to the remaining secondary and tertiary amines. A small slope, as observed for tertiary amines, means that the change in standard enthalpy with temperature is small, suggesting a lower heat requirement in the regeneration process. As shown by the chemical structures of the amines in Table 1, the presence of an OH group at various locations on the molecule resulted in lower pKa values. As presented in Table 9, and among the amines studied in this work, [2-DIPA] and [PZ] had the highest pKa values, and [TMEEA] had the lowest. Generally, primary and secondary amines are stronger bases than tertiary amines. Cyclic amine , [2-HEAN], with two tertiary amine, [THPEDA] and [TMEEA] are the weakest bases. Most likely because of a steric effect, [2-HEAN] is a weaker base than [2-DIPA]. Conversely, [HEP], an alkanolamine with an intermolecular hydrogen bonding tendency, has as lower pKa value than those of [PZ], [2-DIPA], and [2-DMAEMAE]. Normally, amines with higher pKa values react in an acid− base reaction with a stronger interaction with the absorbed molecules, and thus, higher heat is required during the regeneration step. This behavior is attributed to the existence of a linear correlation between the pKa value and the second order rate constant as described by Hetzer et al.21 combining both the pKa values along with the values of the standard enthalpy change. Figure 6 was constructed to categorize different amines based on their pKa and enthalpy changes. The new data were added to data we have reported earlier in a similar manner.24 It was found that 2-DMAEMAE and HEP both fell in the preferred region with (pKa, enthalpy change) of (9.18, 27.77 kJ/mol) and (9.10, 29.94 kJ/mol), respectively, followed by 2-HEAN, which had a relatively lower pKa value (8.66) but acceptable standard enthalpy change (34.24 kJ/mol).

CONCLUSIONS The dissociation constants of nine primary, tertiary, and cyclic amines were measured at five temperatures (298.15, 303.15, 313.15, 323.15 and 333.15) K. The thermodynamic quantities for the dissociation of amines in water were calculated using the van’t Hoff equation. The influence of the addition of (OH) and (CH3) groups to the basic amine structure was investigated. [2-DMAEMAE] was identified as having a high pKa as well as a low enthalpy of absorption and can therefore be considered, in this regard, as a promising amine in CO2 removal applications. TMEEA had the lowest standard state enthalpy change. pKa prediction using PDS was found to give results in very good agreement with experimental data. ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.5b00517. Additional calculations and equations (PDF)

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*E-mail: [email protected]. Notes

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(23) Enea, O.; Houngbossa, K.; Berthon, G. Chaleurs de Protonation de la Piperazine et de Quelques-uns de ses Derives. Electrochim. Acta 1972, 17, 1585−1594. (24) Rayer, A. V.; Sumon, K. K.; Henni, A. Dissociation Constants (pKa) of Tertiary and Cyclic Amines: Structural and Temperature Dependences. J. Chem. Eng. Data 2014, 59, 3805−3813. (25) Sumon, K.; Henni, H.; East, A. Predicting P K A of Amines for CO2 Capture: Computer Versus Pencil-and-Paper. Ind. Eng. Chem. Res. 2012, 51, 11924−11930. (26) Perrin, D.; Dempsey, B.; Serjeant, E. pKa Prediction for Organic Acids and Bases; Springer, 1981.

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Dissociation Constant - ACS Publications - American Chemical Society

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