Kinetic and Equilibrium Reactions of a New Heterocyclic Aqueous 4

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Kinetic and Equilibrium Reactions of a New Heterocyclic Aqueous 4-aminomethyltetrahydropyran (4-AMTHP) Absorbent for Post Combustion Carbon Dioxide (CO) Capture Processes 2

Lichun Li, Sarah Norman, Graeme Puxty, Marcel Maeder, Robert C. Burns, Hai Yu, and William Owen Conway ACS Sustainable Chem. Eng., Just Accepted Manuscript • DOI: 10.1021/ acssuschemeng.7b02149 • Publication Date (Web): 18 Aug 2017 Downloaded from http://pubs.acs.org on August 21, 2017

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ACS Sustainable Chemistry & Engineering

Kinetic and Equilibrium Reactions of a New Heterocyclic Aqueous 4aminomethyltetrahydropyran (4-AMTHP) Absorbent for Post Combustion Carbon Dioxide (CO2) Capture Processes Lichun Li1, Sarah Norman1, Graeme Puxty2, Marcel Maeder1, Robert Burns1, Hai Yu2, and William Conway2,* (1)

Discipline of Chemistry, University of Newcastle, 1 University Drive, NSW 2304, Australia.

(2)

CSIRO Energy, 10 Murray Dwyer Circuit, Mayfield West NSW 2304 Australia.

* Corresponding author – Dr William Conway, [email protected] Keywords Chemical absorption, ethanolamine, coal fired power station, thermodynamics, sustainable industrial chemistry, 13C 1H NMR spectroscopy, stopped flow spectrophotometry. Abstract Aqueous amine absorbent processes remain at the forefront of existing technologies for the removal of CO2 from industrial and large scale power generation flue gas streams. It is essential that improvements in amine-based absorbent technologies are made in order to reduce both capital and operational costs. Intimate understanding of the fundamental chemical behaviour of new amine absorbents systems is an intelligent pathway towards higher efficiency amine based CO2 capture processes. Herein we investigate and report for the first time the complete temperature-dependent kinetic and equilibrium behaviour of a new heterocyclic amine 4-aminomethyltetrahydropyran (4-AMTHP), with CO2, in aqueous solutions. Stopped-flow spectrophotometry, 1H NMR spectroscopy, and potentiometric titration measurements have been performed over the temperature range 25.0 – 45.0oC and the corresponding rate constants for the reversible formation of the carbamic acid, together with equilibrium constants describing the stability of the carbamate, and the protonation of the amine, reported here. Thermodynamic analysis of the resulting constants using the Eyring, Arrhenius, and van’t Hoff relationships has revealed the activation energies, enthalpies, and entropies

for

the

reactions

allowing

a

comparison

to

the

industrial

standard

monoethanolamine (MEA). From the kinetic data the performance of 4-AMTHP was found to be superior to MEA and in line with the established Brønsted relationship between the second order rate constant and the protonation constant or basicity of the amine. The largely

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negative protonation enthalpy (-47 kJ/mol), among the key chemical drivers for CO2 regeneration, is again superior to MEA (-41 kJ/mol). Together the combination of kinetic and equilibrium properties of 4-AMTHP strongly position 4-AMTHP as a promising candidate for more intensive evaluations as a CO2 capture absorbent. Introduction Aqueous alkanolamine absorbent solutions have been widely employed for the chemical removal of carbon dioxide (CO2) containing gases emitted from industrial sources such as fossil fuel-fired power generation [1-3] . The technology has recently been deployed at full scale at the Saskpower Boundary Dam CO2 capture facility in Canada [4] and as part of the Petranova Project at the NRG energy power station in Texas USA [5]. Capital costs related to initial equipment purchases (i.e. solvent, process components, steel, cement, etc.) together with the ongoing energy requirements for operation cost, including high-temperature CO2 and absorbent regeneration steps, are still the major costs for a conventional CO2 capture process. Despite the seeming technology readiness of PCC, widespread uptake of the technology still demands improvements to make the capture process even more affordable. Fundamental improvements to the chemical properties of the reactive CO2 capture absorbent is perceived as one pathway to achieve significant decreases in the initial and ongoing costs. Many efforts have been made to increase the CO2 absorption rate of aqueous amine absorbent solutions through intelligent and targeted solvent screening and design, and by selective blending of two or more individual absorbent components together. Previous research has suggested the properties of fast kinetics and high absorption capacity, particularly the behaviour observed in the cyclic family of amine absorbents i.e. piperazine (PZ), 3- and 4-piperidinealkanols (3PM, 4-PM) making them an attractive family of absorbents [6-10]. The reported second-order rate constants for the reaction of CO2 with PZ, 2-PM, and 3-PM at 25.0oC are 2.43 ×104, 8.2×103 and 2.1×104 M-1s-1, respectively [8, 9, 11]. The second-order rate constants of the cyclic amines is significantly higher than the corresponding value for MEA, 4.9×103 M-1s-1 reaffirming the group of cyclic amines as attractive for CO2 absorption processes both independently and as potential promoters for blending with aqueous ammonia or tertiary and sterically hindered amine solutions. However, the heterocyclic amines/alkanolamines are a much less investigated family of absorbents. Benzylamine (BZA) is one such example which has recently been investigated for its CO2 capture behaviour and it was found to exhibit similar or improved kinetics and cyclic capacity to the industrial standard monoethanolamine


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(MEA) [12, 13]. Therefore, identifying new structurally similar heterocyclic amines possessing rapid CO2 reactivity and high CO2 absorption capacity for the purpose of developing new CO2 absorption solvents is the main focus of the work contained herein. In the current study, we focus our attention on the characterisation of a new heterocyclic absorbent 4-aminomethyltetrahydropyran (4-AMTHP). To our knowledge, there is no open literature reporting the kinetic and equilibrium properties of CO2 reacting with 4AMTHP. Fundamental chemical data including the rate and equilibrium constants for the reaction of CO2 with 4-AMTHP, together with thermodynamic parameters, have been investigated here using stopped-flow spectrophotometric kinetic, 1H NMR spectroscopic, and potentiometric titration measurements at 25.0 - 45.0oC. Experimental section Chemicals High-purity CO2 gas (BOC), N2 (Coregas), 4-AMTHP (99%, Sigma-Aldrich), thymol blue sodium salt, Alizarin red S, and methyl orange, were all used as obtained without further purification. Ultra-high-purity Milli-Q water was boiled to remove CO2 and was used to prepare all solutions for stopped-flow kinetic and 1H NMR measurements. Hydrochloric acid (HCl), Sodium hydroxide (NaOH), Potassium carbonate (KHCO3-), and Sodium Carbonate (Na2CO3) solutions were prepared using Ultra-high-purity Milli-Q water and volumetric glassware. The concentration of NaOH was initially determined by potentiometric titration against a potassium hydrogen phthalate (KHP) solution and the concentration of HCl subsequently determined by titration with the standardised NaOH solution. The concentrations of KHCO3- and Na2CO3 were determined via titration with the standardised HCl solution. CO2(aq) solutions for the stopped-flow measurements were prepared by bubbling pure CO2 gas into a thermostatted water solution in the stopped flow. The concentration of CO2(aq) was determined from the known saturation constants at the desired temperature. Potentiometric titrations Titration of dilute aqueous 4-AMTHP solutions in the absence of CO2 was performed using an automated titration apparatus described in our previous work [14]. Solutions containing 15.0 mM 4-AMTHP were initially acidified with HCl (20.0 mM total HCl) and titrated to high pH using a standardised NaOH solution. Solutions were initially sparged with N2 gas to remove any remaining dissolved gas (other than CO2) and the headspace



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continuously purged with N2 gas throughout the titrations. The mV signal of the solution was recorded throughout the titration and converted into pH during the analysis of the titration data using Reactlab pH titration software (www.jplusconsulting.com). Titrations were repeated in duplicate or triplicate at each temperature from 25.0 - 45.0 oC. No attempt to maintain constant ionic strength was made. Alternatively, activity co-efficient corrections for all charged species were applied during the fitting procedure resulting in thermodynamically correct protonation constants at zero ionic strength. Equilibrium measurements - 1H NMR titrations 1

H NMR spectra of the equilibrated 4-AMTHP/Na2CO3/HCl solutions were measured

on a Bruker Ascend 600MHz instrument at temperatures from 25.0 - 45.0 oC. A series of solutions were prepared in 10.0 mL volumetric flasks containing 4-AMTHP (15.0 mM), Na2CO3 (30.0 mM), together with various amounts of HCl from 0.0 to 40.0 mM. A small volume of each solution (1.5 mL) was then transferred into a series of 5.0 mm diameter NMR tubes together with a sealed glass capillary insert containing TSP (3- (trimethylsilyl)propionic acid-d4, sodium salt) as the reference in a mixed H2O and D2O (1:1 by volume) solvent. Once filled, the NMR tubes were capped and sealed with Parafilm to prevent any loss of CO2 gas before undergoing submersion and thermal equilibration in a water bath (Julabo F20) overnight to ensure that equilibrium was completely attained. Following equilibration each sample was sequentially removed and the 1H NMR spectra recorded immediately inside the temperature controlled NMR. Additional details of the NMR procedure can be found in our previous work [15]. Kinetic measurements - stopped-flow spectrophotometry Kinetic measurements were performed using an Applied Photophysics DX-17 spectrophotometer equipped with a J&M Tidas MCS 500-3 diode-array detector. The reactions were followed by observing the absorption change of coloured acid-base indicators (thus monitoring the pH change) over the wavelength range 400 - 700 nm. All reactions were repeated at temperatures from 25.0 - 45.0 °C and thermostatted to within ±0.1 °C using a circulating Julabo F20 water bath. The exact temperature of the solution was monitored by a thermocouple located within the cell apparatus. Full details of the stopped-flow procedure can be found in our previous work [16].


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Forward reaction – formation of 4-AMTHP carbamic acid/carbamate The formation reactions were performed by mixing aqueous amine solutions containing thymol blue indicator in a 1:1 ratio with an equilibrated CO2(aq) solution in the stopped-flow spectrophotometer. A series of 4-AMTHP solutions were prepared by varying the initial concentration of 4-AMTHP in the solutions. The initial concentrations of the solutions were as follows: [4-AMTHP]0 = 0.5 – 5.0 mM together with 12.5 µM thymol blue indicator, and [CO2]0 = 4.7 mM. Decomposition reaction of 4-AMTHP carbamate The decomposition of 4-AMTHP carbamate was initiated by mixing aqueous solutions of a pre-equilibrated 4-AMTHP carbamate solution with a range of HCl solutions in the stopped-flow spectrophotometer. Carbamate solutions were initially prepared by mixing 50.0 mM 4-AMTHP and 100.0 mM KHCO3- and equilibrating at the appropriate temperature for the stopped-flow measurements overnight. The composition of carbamate in the equilibrated solutions were determined directly using

1

H NMR spectroscopy. The initial [HCl]0

concentration was varied independently from 45.0 – 55.0 mM. Acid-base indicators were mixed with HCl solutions to follow the pH change during the reactions. The overall pH change is substantial and more than 2 pH units, thus a combination of two acid-based indicators was employed for the decomposition reactions (including 50.0 µM Alizarin Red S and 25.0 µM methyl orange). Kinetic and equilibrium model The reaction scheme describing all interactions of amine with CO2 and all carbonate species in aqueous solutions has been intensively investigated and is shown in Figure 1. RNH2 represents all primary/secondary amines. For the current study, RNH2 represents 4AMTHP.



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K4

H2CO3

k1

HCO3-

k2

k-1 k9

CO2(aq) + H2O

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K3

CO32-

k-2

k-9

+ H2O

CO2(aq) + OH-

+ H+

k-7

+ RNH2

k7 RNHCO2H

K8

RNHCO2-

RNH3+

K6

RNH2

Figure 1. Full reaction scheme of the RNH2-CO2-H2O system illustrating the reactions and species involved. Solid arrows represent reactions that have been investigated previously and the values employed in this work were sourced from literature; dashed arrows represent the values being investigated and determined in the current work.

In the absence of amine (RNH2), but in the presence of water, there are several equilibrium and protonation reactions involved in the CO2-H2O system. CO2 (aq) participates in two reversible reactions, by reacting with both H2O and OH-, as described in equations (1) and (2). The protonation of CO32-, HCO3- and OH- are also included in the mechanism, described here in equations (3)-(5). The rate and equilibrium constants for equations (1)-(5) have been reported elsewhere and the values were incorporated into the chemical model [1719]. k1 ,K1

→ H 2CO3 CO2(aq) +H2O ←

(1)

k -1

k 2 ,K 2

→ HCO-3 CO2(aq) +OH- ←

(2)

k -2

K3

 → HCO3CO 32- +H + ←

(3)

→ H 2 CO 3 HCO 3- +H + ←

(4)

K5

(5)

K4

→ H 2 O H + +OH - ←


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When RNH2 is added to the CO2-H2O system, a series of additional reactions occur, which have been shown as dashed lines in Figure 1. These include reversible reactions between RNH2 and CO2 to form the carbamic acid, RNHCO2H, in equation (7); protonation reactions of the RNH2 and amine carbamate, RNHCO2-, equations (6) and (8), respectively; and the decomposition reaction of carbamate, RNHCO2- in equation (9). Values of K6, k7, k-7, K7, K8 and K9 have been determined for NH3 and a series of other amines [11, 16, 20], but to our knowledge remain unreported for 4-AMTHP prior to the present study. Determination of these parameters and their temperature dependence, which allows for an overall thermodynamic assessment of 4-AMTHP for the CO2 capture processes is one of the primary tasks of the current study. (6)

K6

→ RNH 3+ RNH 2 +H + ←

(7)

k ,K

7 7  → RNHCO 2H CO2(aq) +RNH 2 ← 

k -7

(8)

K8

 → RNHCO 2 H RNHCO -2 +H + ←

(9)

K9

→ RNH 2 +HCO 3RNHCO -2 +H 2 O ←

Despite working at low concentrations here, the influence of species charges and activity coefficients can be significant. In favour of concentrations, the Debye-Hückel equation was used for the estimation of the activity coefficients of all ionic species, as shown in equation (10).

log γ =

(10)

−Az 2 µ 1+ µ

In this equation, γ is the activity coefficient of the solution, z is the charge of the species, and μ is the ionic strength of the solution. The parameter A is taken from Manov et al., [21]. Data analysis Potentiometric titrations Titration data in the form of mV vs volume of NaOH added to an acidified 4-AMTHP solution

were

analysed

directly

using

Reactlab

pH

titration

software

(www.jplusconsulting.com). A chemical model comprising reaction (5) for the dissociation of water and reaction (6) describing the protonation of the amine was used during the



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regression. Global analysis of duplicate and triplicate measurements was performed here and the resulting protonation constants at zero ionic strength represent the average fit over the three measurements. 1

H NMR spectroscopic analysis

Relative integrals of the 1H NMR peaks are a direct measure of the hydrogencontaining species in solution. 1H NMR spectral data were used to quantitatively determine the concentrations of the amine and carbamate species using their integrals. It should be noted that the integrals represent the sum of the protonated and deprotonated forms of the amine/protonated amine and carbamate/carbamic acid species due to the fast proton exchange of labile protons on the relative timescale of the 1H NMR measurements [15]. A series of concentration data representing the amine and carbamate species as a function of solution pH (established by addition of a range of HCl concentrations) can then be generated. The spectral measurements were repeated over a range of temperatures from 25.0 - 45.0 oC after reequilibration of the samples overnight at the new temperature. The resulting concentration data were analysed using ReactLab Equilibria software (www.jplusconsulting.com), with a chemical model that included the equilibrium reactions of equations (1) to (6) and equations (8) to (9). The values of K1 to K5 which describe all of the equilibrium reactions in the CO2H2O system, have been well-investigated and their values were obtained directly from the available literature [17-19]. The Reactlab software employs standard Newton-Gauss algorithms which are used for nonlinear least-squares fitting of the equilibrium constants for the decomposition of RNHCO2- to form HCO3-, K9. Due to the fact that all of the 1H NMR measurements were conducted at pH conditions > 8, and that under such conditions the carbamate species is primarily in its deprotonated form, RNHCO2-, determination of the protonation constants of the carbamic acid RNHCO2H, K8 values was not possible. Thus, only K9 values are reported from regression of the 1H NMR data here. Stopped-flow spectrophotometric kinetic analysis, 4-AMTHP Experimental absorbance data obtained via stopped-flow spectrophotometry was analysed

using

an

(www.jplusconsulting.com).

in-house The

extended

in-house

version

version

of

incorporates

Reactlab

Kinetics

activity

coefficient

corrections for all charged species. All of the kinetic and equilibrium reactions are coupled with each other, and with the pH of the solution, and as such they cannot be analysed independently. To overcome this limitation, global analysis of a series of kinetic


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measurements acquired at a single temperature, but under different concentration conditions, was used to obtain the kinetic rate and equilibrium constants of the 4-AMTHP-CO2-H2O system. Similar to the 1H NMR analysis, standard Newton-Guess algorithms were applied to conduct nonlinear least-squares fitting of the spectral data to determine the unknown kinetic and equilibrium constants. In the present study, at each temperature, a set of 10 measurements including 7 carbamate formation reactions, and 3 carbamate/carbamic acid decomposition reactions were analysed together in a single unit in order to determine the constant, k7. The decomposition rate constant, k-7, is then calculated during the regression from the relationship between the formation and decomposition rate constants, and the overall equilibrium constant: K7 = k7/k-7. Using the independently determined value of K9 from the 1H NMR analysis and K7 above, the protonation constant for the carbamate to carbamic acid, K8, was determined from values of the equilibrium constants in a loop as K8=K4K7/K3K2K9 according to the principle of microscopic reversibility. As a measure of the robustness of the fitting procedure the initial concentrations of [CO2]0 were also fitted as part of the analysis and were always found to be within the error limits (± 5.0%). Thermodynamic parameters Arrhenius, Eyring, and van’t Hoff relationships were applied to the kinetic constants, k7 and k-7, and equilibrium constants, K6, K7, K8, and K9, at each temperature, in order to determine the corresponding thermodynamic parameters. The standard molar enthalpy and entropies, ∆Hø and ∆Sø were calculated from a van’t Hoff plot of lnK against 1/T, as outlined in equation (11). The activation energy and Arrhenius parameter were calculated from the slope and intercept of a plot of lnk against 1/T, as shown in equation (12). The ∆H † and ∆S † were determined from the Eyring equation shown in equation (13). van’t Hoff ln K (T ) = -

Arrhenius

ln k (T ) =

Eyring ln

∆H φ ∆S φ + RT R

- Ea + ln A RT

k -∆H † 1 ∆S † = ⋅ + T R T R

(11)

(12)

(13)

In the above equations, K represents the equilibrium constants, ∆H φ represents the standard molar enthalpy change (kJ.mol-1), ∆S φ represents the standard molar entropy change



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(kJ.mol-1), A is Arrhenius constant, Ea is the activation energy for the reaction (J.mol−1), ∆H † represents the enthalpy of activation (kJ.mol-1), ∆S † represents the entropy of activation (kJ.mol-1) and R is the universal gas constant (8.314 J.mol-1. K-1). Results Potentiometric titrations The protonation constants of 4-AMTHP at 25.0, 35.0 and 45.0 oC, together with van’t Hoff Parameters ( ∆H φ and ∆S φ ) are listed in Table 1. The corresponding values from Fernandes et al for MEA are included in Table 1 for comparison. Table 1. Protonation constants at zero ionic strength and thermodynamic parameters for the protonation of 4AMTHP from 25.0, 35.0 and 45.0 oC. MEA values from Fernandes et al.[14] included for comparison. Note – values expressed in molarity (M). Temperature (oC)

K6

→ RNH 3+ RNH 2 +H + ←

van’t Hoff Parameters

25.0

35.0

45.0

∆H ( kJ.mol-1)

4-AMTHP

9.90(1)

9.62(1)

9.39(2)

-47

32

MEA

9.44

9.21

8.96

-41

42

φ

∆S (J.mol-1.K-1)

* MEA values taken from Fernandes et al.

The protonation constants of 4-AMTHP are larger than the values of MEA at similar temperatures indicating that 4-AMTHP is a stronger Lewis base than MEA. The increase in the Lewis basicity of 4-AMTHP is expected to increase its reactivity towards CO2. This will be discussed further in future sections of the manuscript. 1

H NMR equilibrium

The equilibrium concentrations of 4-AMTHP, and the corresponding carbamate, determined from the 1H NMR spectra together with the calculated species concentration profiles as a function of pH from the analysis, are shown in Figure 2(a). The markers in Figure 2(a) are the experimental values determined from the 1H NMR measurements while the solid lines are the calculated curves resulting from the fitting of the carbamate stability constant, K9, at each temperature. From Figure 2(a), the pH of the initial solution containing 15.0 mM 4-AMTHP and 30.0 mM Na2CO3 and in the absence of HCl is high (~ pH 11.0) and, at these conditions, the majority of the amine is in its deprotonated form (RNH2). From Figure 2(b) the 4-AMTHP carbamate (RNHCO2-) is stable at high pH and is present in small quantities (~ 2.0mM, or ~13.3% of the total amine concentration). As the HCl concentration


Page 11 of 18

is increased, and subsequently the pH of the solution is decreased, the concentration of carbamate increases, reaching a maximum concentration between pH ~9.5 - 10.0. The concentration of carbamate is strongly temperature dependent such that both the maximum concentrations and the location of the maximum peak height shift downwards in concentration and towards lower pH with increasing temperature. Under all three investigated temperatures, the maximum concentrations of the carbamate species were observed to reach ~ 4.5 – 6.0 mM or ~ 30.0 to 40.0 % of the total amine concentration. Further addition of HCl and hence a decrease in pH drives the chemical equilibrium away from carbamate towards HCO3-. Correspondingly, the amine is protonated during this process and RNH3+ becomes the dominant form of the amine in the solution. 0.016

0.030 (a)

(b)

0.014 0.025

0.010

Concentration (M)

RNH2/RNH3+

0.012

Concentration (M)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


25 oC 35 ooC 45 C

0.008 0.006

0.020

HCO3CO32-

0.015 RNH2

RNH3+

0.010

RNHCO2-

0.004 0.005

RNHCO2-/RNHCO2H

0.002 0.000

OH-

0.000 8

8.5

9

9.5

10

10.5

11

11.5

8

8.5

9

9.5

pH

10

10.5

11

11.5

pH

Figure 2. (a) Equilibrium concentrations of RNH2/RNH3+ and RNHCO2-/RNHCO2H as a function of pH at temperatures from 25.0 – 45.0 oC; (b) Calculated species concentration in the equilibrated solutions as a function of pH at 25.0 oC.

The calculated log K9 values are reported in Table 2 together with the corresponding theromdynamic values calculated from the van’t Hoff analysis. Previously determined values for MEA from Fernandes et al have been included in the table for comparison. Table 2. Equilibrium constants at 25.0, 35.0 and 45.0 oC and the relative thermodynamic parameters of the decomposition of RNHCO2-. MEA values from Fernandes et al.[14] are included for comparison. Note – values expressed in molarity (M). Temperature (oC) 25.0

35.0

45.0

van’t Hoff Parameters ∆ΗØ

∆SØ -1

K

9 → RNH 2 +HCO -3 ← RNHCO -2 +H 2 O

( kJ.mol )

( J.mol-1.K-1)

4-AMTHP

1.89

1.83

1.67

-20

-6

MEA

1.76

1.66

1.55

-18

-25

Log K9 values are similar but marginally larger for 4-AMTHP compared to MEA at similar temperatures. More pertinent than the values themselves is the temperature dependence of the stability constant K9 which is a key driver of the chemical change during the regeneration of the amine and CO2 at higher temperatures. Large changes in the



equilibrium constant with temperature, as observed by a more negative ∆Ηǂ, could result in lower energy requirements when compared to similar carbamate forming absorbents with lower enthalpies. The enthalpy for the carbamate formation reaction here is similar to MEA and falls within the range of values for similar carbamate forming amines [20]. Stopped-flow kinetic data To follow the formation of the carbamate/carbamate acid, the reaction between 4AMTHP solutions, and CO2(aq), was conducted under conditions where the formation of the carbamate is the dominant reaction. Thymol blue acid-base indicator was used to observe the pH changes in the solutions during the CO2 absorption. Example series of the experimental absorbance profiles together with the calculated concentration profiles of the reactions between 4-AMTHP and CO2(aq) (4.46 mM) at 25.0 oC in the presence of 12.5uM thymol blue indicator are shown in Figure 3. NOTE - markers represent the experimental results from the measurement, and the lines are the calculated values after the final regression of the rate and equilibrium constants. As can be seen from

Figure 3(a),

the reaction appears to be

slower with the increase of amine concentration. This phenomenon is due to the increased buffer action of the amine resulting in increasingly smaller pH changes as its concentration is increased. The concentration profiles for the reactions, Figure 3(b) provide a more representative insight into the chemical behaviour during the carbamate formation reactions where CO2 predominantly reacts with RNH2 to form RNHCO2- and HCO3-, thus decreasing the pH of the solution and correspondingly RNH2 is protonated forming RNH3+. Notably the reactions are rapid and reach completion within ~0.5 secs. (b)

0.2 5mM 4-AMTHP

0.16

4mM 4-AMTHP

0.14

3mM 4-AMTHP

0.12

2mM 4-AMTHP

10.00 pH

1mM 4-AMTHP

0.08

12.00

0.004

1.5mM 4-AMTHP

0.1

0.005

0.5mM 4-AMTHP 0.06

8.00 0.003 6.00

CO2(aq) RNH3+

0.002

4.00

RNHCO2-

0.04

0.001

0.02

pH

0.18

Concentration (M)

(a)

Absorbance at 590 nm

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 12 of 18

RNH2

2.00 HCO3-

0 0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4

0.45

0.5

Time (s)

0

0.00 0.00

0.10

0.20

Time (s)

0.30

0.40

0.50

Figure 3. (a) Absorbance traces at 590 nm of 4.46 mM CO2(aq) reacting with a series of 4-AMTHP solutions with concentrations ranging from 0.5 – 5.0 mM; (b) Concentration profile together with pH vs time describing the reaction of 4.0 mM 4-AMTHP with 4.46 mM CO2(aq).

The decomposition of 4-AMTHP carbamate/carbamate acid, RNHCO2-/RNHCO2H, was initiated here by reacting a pre-equilibrated 4-AMTHP/KHCO3- solution with a range of


Page 13 of 18

HCl solutions spanning a range of concentrations. An example series of measurements demonstrating the effect of HCl concentration, and thus pH, on the kinetics of decomposition of 4-AMTHP carbamate/carbamate acid at 25.0 oC is shown in Figure 4(a). Example calculated concentration profiles, and pH profile, as a function of time during the reaction of 55.0 mM HCl with carbamate solutions is shown in Figure 4(b). (a) 0.45

(b) 0.045

0.4

0.040

6.00 CO2(aq) pH

55mM HCl

5.00 0.035

50mM HCl 45mM HCl

0.3 0.25 0.2 0.15

0.030

4.00 RNH3+

0.025

3.00

pH

0.35

Concentration (M)

Absorbance at 590 nm

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.020 0.015

0.1

0.010

0.05

0.005

2.00

RNHCO2H

HCO31.00

0

H2CO3

0.000 0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4

0.45

0.5

0.00 0.00

0.10

0.20

Time (s)

0.30

0.40

0.50

Time (s)

Figure 4 (a) Absorbance traces at 550 nm of 4-AMTHP/KHCO3- solution reacting with 45.0, 50.0 and 55.0 mM HCl solutions; (b) Calculated concentration profile and pH vs time in the reaction of 55.0 mM HCl with pre-equilibrated 0.05 M 4-AMTHP/0.1 M KHCO3- solution.

Immediately after mixing with HCl, and before reaction, the pH of the solution instantaneously drops below 6.0, resulting in protonation of 4-AMTHP, HCO3- and 4AMTHP carbamate to form RNH3+, H2CO3 and RNHCO2H respectively. Following decomposition of the carbamic acid species, evolved free amine immediately protonates to RNH3+ and similarly carbonic acid re-equilibrates to form HCO3- and CO2(aq). The pH of the solution increases from ~3.0 to ~6.0 during the reaction. Global analysis of the 7 forward and 3 backward sets of stopped-flow data at 25.0, 35.0 and 45.0 oC enables the determination of the reaction rate/equilibrium constants for CO2 and 4-AMTHP. The complete set of rate and equilibrium constants including k7, k-7, K7 and K8 together with the relevant thermodynamic properties determined here are listed in Table 3. To improve the robustness of the fitting procedure the log K9 values determined from the 1H NMR study were employed here in the calculation of K8 using microscopic reversibility where K8 = K4.K7/(K3.K2.K9). Table 3. Reaction rate and equilibrium constants determined from the current work

Temperature (oC)

25.0

35.0

Arrhenius

45.0

Ea kJ.mol-1

Eyring A

van’t Hoff

∆Ηǂ

∆Sǂ

∆ΗØ

∆SØ

kJ.mol-1

J.mol-1.K-1

kJ.mol-1

J.mol-1.K-1



k7 (M-1s-1) -1

k-7 (s )

5.95 × 103

9.53 × 103

1.28 × 104

2

1.21 × 109

30

2

82

9.94 × 10

15

27(3)

-79(11)

79.3(3)

53(1)

44

1.3 × 10

3.53 × 10

K7 (M )

135

73

36

-52(3)

-132(10)

K8

6.6

6.35

6.18

-38(4)

-2(12)

-1

Discussion Comparison of 4-AMTHP to other amines The present study reports the fundamental kinetic and equilibrium constants for the reactions of 4-AMTHP with CO2(aq) for the first time. Relevant theories can now begin to be established about the chemical behaviour of 4-AMTHP by comparing to similar chemical data for other amines including a group of cyclic amines, MEA, and NH3. Among the suite of kinetic and equilibrium constants, the k7 pathway plays an essential role in the overall suitability of amine absorbents as it describes the reactivity of amine towards CO2 and is strongly coupled to the size of the absorber column. Often, the underlying kinetic behaviour of an amine is confirmed by determining the relationship between kinetics, k7, and the basicity of the molecule, pKa value of each protonated amine, here, in a Brønsted plot. Such a plot implies that the reactivity of an amine towards CO2 is increased with its Brønsted basicity i.e. k7 increasing with pKa. Values lying away from the general trend for the majority of amines are broad indications of the presence or influence of steric hindrance and other electronic effects on the kinetic reactivity [22]. A Brønsted plot compiled from the data here and corresponding data from Conway et al.[11] is shown here in Figure 5 (a). The data for 4AMTHP is shown as a diamond in red in Figure 5 (a). It can be observed from Figure 5 (a) that the result for 4-AMTHP agrees well with other cyclic and linear amines. This strongly supports the robustness of the values determined in the current study and the trend indicates that there are no competing steric or electronic effects outside of the Brønsted effect when 4AMTHP is reacting with CO2. (a)

(b)

5.00

5.00

H N

H N

4-PIPME

PZ N H

NH2

NH2

BZA

4.00

MEA H2N

MEA

N H

NH3

H N

1-MPZ H N

OH

MORP O N

N H

CH3

CH3

NH3

2.00

4-PIPME

S

3.00

N

N H

N H

OH

TMORP

O

3.00

PIPD

H2N

4-AMTHP

OH

H N

S

PZ

N H

BZA

N H

4-AMTHP 1-MPZ

OH

TMORP MORP

PIPD H N

logk7

4.00

logk7

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 14 of 18

N H

2.00 8

8.5

9

9.5

10

10.5

11

11.5

-60

-55

-50

-45

-40

-35

-30

ΔΗprotonation(kJ.mol-1)

log K6

Figure 5. Plot of the rate constant for the formation of carbamic acid from amine and CO2(aq), logk7, vs the protonation constant of the amine, log K6, at 25.0oC (b) Plot of the rate constant for the formation of RNHCO2H from


Page 15 of 18

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


amine and CO2(aq), logk7, at 25.0oC, vs the enthalpy of the protonation of amine (∆Η). Note - diamonds represent cyclic amines and dots represents MEA and NH3.

Another vitally important parameter involves the enthalpy of the amine protonation reaction, ∆Η, which has been suggested to contribute up to ~60% of the total reaction enthalpy required for solvent regeneration [23]. Similarly to the Brønsted plot, a plot of the kinetic constant, log k7, against the enthalpy of amine protonation, ∆H, is shown in Figure 5 (b). The ideal amine will have both a high reactivity towards CO2 (k7) and large negative enthalpy for amine protonation (∆Η), thus this ideal amine should appear in the top left-hand corner of Figure 5 (b). It can be observed from Figure 5 (b) that the reaction rate of 4AMTHP is similar to MEA and other cyclic amines presented here (similar position along yaxis) and higher than NH3 (lower values on the y-axis respectively). Interestingly, 4-AMTHP has the second largest negative enthalpy (∆Η) of the cyclic amines presented here (only BZA exceeding 4-AMTHP of the amines presented) positioning it further to a more desirable position to the left on the x-axis. From the data here it appears that the intramolecular location of the amine group external to the ring structure is favourable given the majority of the cyclic molecules with intermolecular amine groups (secondary amines with amine located within the ring) appear to have lower protonation enthalpies. This observation is an interesting insight which should be considered in the design of new amine molecules for CO2 capture. Interestingly, NH3 also has a more exothermic protonation enthalpy than 4-AMTHP making a blend of 4-AMTHP with NH3 a potentially attractive option. Conclusions Characterisation of a new heterocyclic absorbent 4-aminomethyltetrahydropyran (4AMTHP) for its kinetic and equilibrium behaviour when reacting with CO2 has been investigated. The corresponding rate constants for the reversible formation of carbamic acid together with equilibrium constants describing the stability of the carbamate, and the protonation of the amine, reported here for the first time to our knowledge. Thermodynamic analysis of the resulting constants using the Eyring, Arrhenius, and van’t Hoff relationships has revealed the activation energies, enthalpies, and entropies for the reactions allowing a comparison to the industrial standard monoethanolamine (MEA). Compared to the benchmark CO2 absorption solvent MEA, 4-AMTHP has the advantages of a similar CO2 absorption rate and a larger negative energy which will ideally benefit the energy required to regenerate CO2. These chemical properties make 4-AMTHP a promising alternative for MEA.



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Results from this study contribute to the development of both an advanced 4-AMTHP absorbent on outright or blended with other amines for CO2 capture. References 1.

2. 3. 4.

5. 6.

7.

8.

9.

10.

11.

12.

13.

14.

15.

Chakraborty, A., G. Astarita, and K. Bischoff, CO2 absorption in aqueous solutions of hindered amines. Chemical Engineering Science, 1986. 41(4): p. 997-1003, https://doi.org/10.1016/0009-2509(86)87185-8. Yeh, J.T., H.W. Pennline, and K.P. Resnik, Study of CO2 absorption and desorption in a packed column. Energy Fuels, 2001. 15(2): p. 274-278, DOI: 10.1021/ef0002389. Rochelle, G.T., Amine scrubbing for CO2 capture. Science, 2009. 325(5948): p. 16521654, DOI: 10.1126/science.1176731. Stéphenne, K., Start-up of world's first commercial post-combustion coal fired CCS project: contribution of Shell Cansolv to SaskPower Boundary Dam ICCS Project. Energy Procedia, 2014. 63: p. 6106-6110, https://doi.org/10.1016/j.egypro.2014.11.642. Jenkins, J., A CASE STUDY OF THE PETRA NOVA CARBON CAPTURE PROJECT. 2015. Freeman, S.A., et al., Carbon dioxide capture with concentrated, aqueous piperazine. Energy Procedia, 2009. 1(1): p. 1489-1496, https://doi.org/10.1016/j.egypro.2009.01.195. Rochelle, G., et al., Aqueous piperazine as the new standard for CO2 capture technology. Chemical Engineering Journal, 2011. 171(3): p. 725-733, https://doi.org/10.1016/j.cej.2011.02.011. Conway, W., et al., Reactions of CO2 with aqueous piperazine solutions: formation and decomposition of mono-and dicarbamic acids/carbamates of piperazine at 25.0° C. J.Phys.Chem. A, 2013. 117(5): p. 806-813, DOI: 10.1021/jp310560b. Conway, W., et al., CO2 absorption into aqueous solutions containing 3piperidinemethanol: CO2 mass transfer, stopped-flow kinetics, 1H/13C NMR, and vapor–liquid equilibrium investigations. Ind.Eng.Chem.Res., 2014. 53(43): p. 1671516724, DOI: 10.1021/ie503195x. Jackson, P., Experimental and theoretical evidence suggests carbamate intermediates play a key role in CO2 sequestration catalysed by sterically hindered amines. Structural Chemistry, 2014. 25(5): p. 1535-1546, DOI: 10.1007/s11224-014-0431-5. Conway, W., et al., Toward rational design of amine solutions for PCC applications: the kinetics of the reaction of CO2 (aq) with cyclic and secondary amines in aqueous solution. Environ.Sci.Technol., 2012. 46(13): p. 7422-7429, DOI: 10.1021/es300541t. Richner, G., et al., Thermokinetic properties and performance evaluation of benzylamine-based solvents for CO2 capture. Chemical Engineering Journal, 2015. 264: p. 230-240, https://doi.org/10.1016/j.cej.2014.11.067. Conway, W., et al., Rapid CO2 absorption into aqueous benzylamine (BZA) solutions and its formulations with monoethanolamine (MEA), and 2-amino-2-methyl-1propanol (AMP) as components for post combustion capture processes. Chemical Engineering Journal, 2015. 264: p. 954-961, https://doi.org/10.1016/j.cej.2014.11.040. Fernandes, D., et al., Protonation constants and thermodynamic properties of amines for post combustion capture of CO2. Journal of Chemical Thermodynamics, 2012. 51: p. 97-102, https://doi.org/10.1016/j.jct.2012.02.031. Fernandes, D., et al., Investigations of primary and secondary amine carbamate stability by 1 H NMR spectroscopy for post combustion capture of carbon dioxide.


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16.

17.

18. 19.

20.

21.

22.

23.

Journal of Chemical Thermodynamics, 2012. 54: p. 183-191, https://doi.org/10.1016/j.jct.2012.03.030. Wang, X., et al., Kinetics of the Reversible Reaction of CO2 (aq) with Ammonia in Aqueous Solution. J.Phys.Chem.A, 2011. 115(24): p. 6405-6412, DOI: 10.1021/jp108491a. Wang, X., et al., Comprehensive study of the hydration and dehydration reactions of carbon dioxide in aqueous solution. J.Phys.Chem.A, 2009. 114(4): p. 1734-1740, DOI: 10.1021/jp909019u. Harned, H.S. and S.R. Scholes Jr, The Ionization Constant of HCO3- from 0 to 50oC. J.Am.Chem.Soc., 1941. 63(6): p. 1706-1709, DOI: 10.1021/ja01851a058. Maeda, M., et al., Estimation of salt and temperature effects on ion product of water in aqueous solution. Bulletin of the Chemical Society of Japan, 1987. 60(9): p. 32333239, https://doi.org/10.1246/bcsj.60.3233. Conway, W., et al., Comprehensive Kinetic and Thermodynamic Study of the Reactions of CO2 (aq) and HCO3– with Monoethanolamine (MEA) in Aqueous Solution. J.Phys.Chem.A, 2011. 115(50): p. 14340-14349, DOI: 10.1021/jp2081462. Manov, G.G., et al., Values of the constants in the Debye—Hückel equation for activity coefficients. J.Am.Chem.Soc., 1943. 65(9): p. 1765-1767, DOI: 10.1021/ja01249a028. Conway, W., et al., Toward the understanding of chemical absorption processes for post-combustion capture of carbon dioxide: electronic and steric considerations from the kinetics of reactions of CO2 (aq) with sterically hindered amines. Environ.Sci.Technol., 2012. 47(2): p. 1163-1169, DOI: 10.1021/es3025885. McCann, N., M. Maeder, and M. Attalla, Simulation of enthalpy and capacity of CO2 absorption by aqueous amine systems. Ind.Eng.Chem.Res., 2008. 47(6): p. 2002-2009, DOI: 10.1021/ie070619a.



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For table of contents use only O C O O NH2

O

k7

H N

k-7 K8

O

OH

O H N

O

+ H O

Kinetic and equilibrium reactions leading to the formation of 4-AMTHP carbamic acid (green) and carbamate (pink) – an example of the main chemical pathways employed during rapid CO2 absorption by primary amines for sustainable carbon capture utilisation and storage (CCUS) from industrial flue gases.


Page 18 of 18

Kinetic and Equilibrium Reactions of a New Heterocyclic Aqueous 4

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