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Evaluation and Modelling of Vapour-Liquid-Equilibrium and CO2 Absorption Enthalpies of Aqueous Designer Di-amines for Post Combustion Capture Processes Weiliang Luo, Qi Yang, William Owen Conway, Graeme Puxty, Paul H. M. Feron, and Jian Chen Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 31 May 2017 Downloaded from http://pubs.acs.org on June 2, 2017
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Evaluation and Modelling of Vapour-Liquid-Equilibrium and CO2
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Absorption
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Combustion Capture Processes
Enthalpies
of
Aqueous
Designer
Di-amines
for
Post
4 5
Weiliang Luo†, Qi Yang‡, William Conway§,*, Graeme Puxty§, Paul Feron§, and Jian Chen†,*
6
†
State Key laboratory of Chemical Engineering, Tsinghua University, Beijing 100084, China
7
‡
CSIRO Manufacturing, Clayton, VIC 3168, Australia
8
§
CSIRO Energy, Mayfield West, NSW 2304, Australia
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* Corresponding authors:
12
William Conway, e-mail:
[email protected]; phone: +61 2 4960 6098. ORCID: 0000-0002-7958-3872
Jian Chen, e-mail:
[email protected]; phone: +86 10 6279 8627. ORCID: 0000-0003-1695-7790
13 14
Abstract
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Novel absorbents with improved characteristics are required to reduce the existing cost
16
and environmental barriers to deployment of large scale CO2 capture. Recently, bespoke
17
absorbent molecules have been specifically designed for CO2 capture applications, and their
18
fundamental properties and suitability for CO2 capture processes evaluated. From the study,
19
two unique di-amine molecules, 4-(2-hydroxyethylamino)piperidine (A4) and 1-(2-
20
hydroxyethyl)-4-aminopiperidine (C4), were selected for further evaluation including
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thermodynamic characterisation. The solubilities of CO2 in two di-amine solutions with a
22
mass fraction of 15% and 30% were measured at different temperatures (313.15 - 393.15 K)
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and CO2 partial pressures (up to 400 kPa) by thermostatic VLE stirred cell. The absorption
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enthalpies of reactions between di-amines and CO2 were evaluated at different temperatures
25
(313.15 and 333.15 K) using a CPA201 reaction calorimeter. The amine protonation
26
constants and associated protonation enthalpies were determined by potentiometric titration.
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The interaction of CO2 with the di-amine solutions was summarized and a simple
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mathematical model established that could make a preliminary but good prediction of the
29
VLE and thermodynamic properties. Based on the analyses in this work, the two designer di-
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amines A4 and C4 showed superior performance compared to amines typically used for CO2
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capture and further research will be completed at larger scale.
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TOC/Abstract art
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1. Introduction
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Extreme weather events most likely related to the enhanced greenhouse effect have
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attracted widespread attention worldwide. The emission of CO2, the major components
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driving rapid enhancement of the greenhouse effect, has long been suspected as a major
40
cause of this issue. Carbon capture and storage (CCS), specifically post-combustion capture
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(PCC) technologies, have been identified as a crucial and key strategy in the development
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and implementation of short to intermediate term carbon reduction strategies from coal-fired
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power generation1, 2. Currently, commercial PCC processes employ amine-based chemical
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absorption to selectively bind CO2 present in industrial process flue gases, and it is this
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method, often referred to as amine scrubbing, that is becoming one of the dominant
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technology options for CO2 capture from coal-fired power plants by 20303. Aqueous
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alkanolamine solutions have been well researched for use in chemical absorption processes.
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Among the vast list of absorbents, monoethanolamine (MEA) is considered to be the most
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mature reference solvent for its high capacity and fast absorption rate with CO2. However,
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several shortcomings with MEA including high thermal and oxidative degradation rates, high
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corrosion, and notably high energy consumption for solvent regeneration, limit the large scale
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deployment of MEA processes. New and improved solvents are needed to overcome these
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issues and improve confidence in PCC technology.
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Traditionally, large amounts of experimental data are required to characterize a new
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solvent and the measurement process is very time-consuming. As a result, different criteria
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have been presented in the literature for pre-selection of molecules preceding full
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characterization and, if successful, subsequent pilot plant experimentation. Rapid CO2
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absorption rates as well as high cyclic capacity have been a critical criterion for absorbents; a
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promising amine ideally possessing both characteristics among other properties4. 2
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Furthermore, energy conservation in solvent regeneration is considered to be a brand new
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criterion in recent research5. Previous studies proved that heterocyclic amines show better
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performance in almost every field of CO2 capture6, 7, and piperazine (PZ) was even suggested
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as a substitute for MEA as the reference solvent for its superior capabilities over MEA
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including higher CO2 capture capacity, faster CO2 absorption rate, lower degradation, lower
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corrosion, and above all, lower regeneration energy consumption8. However, precipitation
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issues stemming from high concentrations of PZ solution and high CO2 loadings restrict its
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application window9. PZ derivatives with different substituents, such as 1-methylpiperazine
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(1-MPZ), 1-ethylpiperazine (1-EPZ), 1,4-dimethylpiperazine (DMPZ), have been proven to
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have better performance for CO2 capture than PZ10, 11. Based on the above, we attempted to
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utilise the robust cyclic structure of PZ or similar while minimising or eliminating
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precipitation generated in the presence of CO2 at high concentrations.
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The broader concept of di-amines specifically designed to perform well in CO2 capture
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applications, were initially investigated in our previous work12, 13. The two parent skeleton
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amines share structural similarities to the sub-family of piperidine derivatives that, as a group
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of amines, were found to display excellent CO2 loadings in our previous study14. Selected
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functional groups were added for structural modification of the skeleton molecules, including
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carbon chains to provide connection sites for additional groups, hydroxyl groups to increase
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the boiling point and reduce volatilization, and additional amine groups to improve the
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solubility of CO2. Through a series of preliminary molecule screening tests, the research
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scope has been narrowed down to four promising candidates. Due to limited availability of
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material and costs, two molecules, designated herein as A4 and C4, were selected for further
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research.
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In this work, designer di-amines A4 and C4 were evaluated for their CO2 solubilities and
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absorption enthalpies alongside MEA for validation and comparison. Cyclic capacities were
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calculated from the solubility data as a function of temperature and CO2 partial pressure.
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Fundamental reaction mechanisms between two di-amines and CO2 were established
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according to which, a simple thermodynamic model of the di-amine-CO2-H2O vapour-liquid
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equilibrium was employed for regression of the equilibrium constants and enthalpies for the
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carbamate formation reactions. Using this model and the equilibrium constants, CO2
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solubilities and absorption enthalpies of designer di-amines were recalculated and then
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compared with the experimental data to determine the adequacy of the model for the
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prediction of CO2 equilibrium pressure.
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2. Experimental
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2.1 Chemicals
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Carbon dioxide (CO2, 99.99%) and nitrogen (N2, 99.99%) were purchased from Beijing
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Qianxi Gas Chemical Industry Co., Ltd., monoethanolamine (MEA, 99%) was purchased
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from Aladdin Industrial Corporation; they were all used without further purification. 4-(2-
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hydroxyethylamino)piperidine (A4, 99%) and 1-(2-hydroxyethyl)-4-aminopiperidine (C4,
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99%) were synthesised using the published methods13, 15. All solutions were prepared with
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deionized water. The structures of the di-amines investigated in this work are shown in Table
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1. The concentrations of solutions are expressed herein in mass fraction (% w/w).
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2.2 Vapour liquid equilibrium
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Solubility measurements were performed in a thermostatic VLE reactor shown in Figure
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1.The CO2 solubility data represents the relationship between the CO2 partial pressure and the
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total CO2 loading in the solution. The CO2 loading was determined using the difference
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between the total amount of CO2 introduced into the cell and that of unreacted CO2 in the gas
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phase using the Peng-Robinson cubic equation16. A detailed description of the stirred cell,
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calculations, and estimated uncertainty can be found in our previous work17.
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2.3 CO2 Absorption enthalpies
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Enthalpy measurements were performed using the CPA201 reaction calorimeter from
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ChemiSens AB shown in Figure 1. The true heat flow was measured by a heat transducer at
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the bottom of the 250 mL sealed reactor, and a high precision Peltier element was used to
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guarantee one-directional flow of heat. 100 mL of amine solution was placed in the reactor
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and degassed under vacuum. The experiment was performed at a constant temperature until
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equilibrium was reached. The vapour pressure of amine and water in the gas phase of the
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reactor was assumed to be constant and equal to the total equilibrium pressure before the first
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addition of CO2. CO2 was introduced batch-wise, and approximately same amount of CO2
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was added into the reactor when equilibrium was achieved after the previous addition. An
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automation script, which sought stability in pressure and true heat flow between each addition
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of CO2, was used to control the experiments. The maximum deviation in the pressure and true
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heat flow were required to be below ±0.01 bar and ±0.02 W respectively within a period of
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10 minutes, so that the equilibrium for the system was considered to be reached.
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The main source of uncertainty in the experimental data stems from the amount of CO2
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added to the reactor. The amount of CO2 added to the reactor was recorded using a BIOS 4
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flow meter with an accuracy of ±0.8%. The accuracies of the temperature sensor and the
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pressure transducer were ±2.5 K and ±1.0%, respectively. To obtain the total amount of heat
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released during the absorption of CO2, the heat flow curve was integrated over the duration of
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the loading interval. The uncertainty in the amount of heat released was estimated to be
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±2.5%.
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2.4 Amine protonation constants
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The protonation constants of A4 and C4 (Eq. (4) and (5) below) was determined
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between 288.15 and 318.15 K by potentiometric titration adapted from the method of
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Fernandes, et al18. The amine solution was initially acidified by addition of standardised
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hydrochloric acid (HCl) and back titrated to high pH with sodium hydroxide (NaOH).
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Titrations were performed using a 665 Dosimat automated burette system in a jacketed
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titration vessel connected to a Julabo ED-5 heating circulator. The pH of the solution was
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measured with a Metrohm combined micro-pH glass electrode. The electrode signal was
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acquired and amplified by a National Instruments NI-DAQ 7 board running with Matlab
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software on a personal computer. The electrode signal, initially recorded in mV, was
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transformed into the pH scale using the Hyperquad2008 software19. At each temperature, the
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electrode was calibrated by titrating HCl solutions of known concentration with standardised
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NaOH solution.
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3. Modelling
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Absorption of CO2 into aqueous di-amine solutions is described by the series of
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equilibrium reactions in Eq. (1) - (9). CO2 reacts with water to form bicarbonate/carbonate
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according to the reactions of Eq. (1) - (3). Eq. (1) is bicarbonate/carbonate protonation
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equilibria, Eq. (2) is the reaction of CO2 and water to form bicarbonate and Eq. (3) represents
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water autoprotolysis. The equations describing the equilibrium constants and their
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temperature dependence for these reaction were taken from previous work20.
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K1 H+ +CO32- ← →HCO3∆H1
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K2 H+ +HCO3- ← →H2O+CO2 ∆H2
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K3 H+ +OH- ← →H2O ∆H3
(1) (2) (3)
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Incorporating the di-amine into the model, the first and second protonation of the di-
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amine, K4 and K5, are described by Eq. (4) and (5).
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K4 H+ +Am← →AmH+ ∆H4
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K5 H+ +AmH+ ← →AmH22+ ∆H5
(4) (5)
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Formation of the mono and di-carbamate, K6 and K7, together with the corresponding
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protonation of the mono and di-carbamate, K8 and K9, are described in Eq. (6) - (9).
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Carbamate protonation has a negligible impact on the prediction of chemical speciation for
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typical mono-amines due to the weakly basic nature of the carbamate carboxyl group.
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However, in the case of di-amines the often strongly basic nature of the unreacted amine
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group plays a more significant role in the chemistry and must be included in the chemical
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model. Formation of the doubly protonated carbamate and doubly protonated di-carbamate is
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unlikely under the conditions here, hence these reactions have been omitted from the model.
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K6 HCO3- +Am← →AmCOO− +H2O ∆H6
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K7 HCO3- +AmCOO- ← → Am ( COO ) 2 +H2O ∆H7
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K8 H+ +AmCOO- ← →HAmCOO ∆H8
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K9 H + +Am ( COO )2 ← → HAm ( COO ) 2 ∆H 9
2-
2-
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(6) (7) (8) -
(9)
Since tertiary amines are unable to react directly with CO2, as is the case for the second amine group in C4, Eq. (7) and (9) are not required for C4.
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To fully determine the speciation in the system of equilibria, equations defining the
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equilibrium constants and mass balance of the system are required. The equilibrium constants
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are defined according to Eq. (10), or Eq. (11) if water is involved, and the mass balance by
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Eq. (12) - (15).
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Ki =
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cP γP γP ⋅ = K i0 ⋅ γ R1 ⋅ γ R 2 cR 1 ⋅ c R 2 γ R 1 ⋅ γ R 2
(10)
Ki K0 cP γP γP = ⋅ = i ⋅ cH 2 O cR1 ⋅ c R 2 γ R 1 ⋅ γ R 2 cH 2 O γ R 1 ⋅ γ R 2
(11)
(
total c AM = cAM + c AMH + + cAMCOO - + cHAMCOO + cAMH 2+ + cAM ( COO )2- + cHAM ( COO )2
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(
total cCO = cCO 2 + cCO 2- + cHCO - + cAMCOO - + cHAMCOO +2 cAM 2 3
3
2
( COO )2-2
2
+ 2 cHAM
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c Htotal = c H + + c H C O - + 2 c C O 2 + c AM H + + 2 c AM H 2 + + c H AM C O O + c H A M + 3
2
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( C O O )-2
+ cH 2O
(14) 184
total cOH + cH2O − = c OH−
(15)
cP and cR represent the concentration of the product species and each reactant
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where
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respectively for each equilibrium reaction, and the concentrations in parentheses were
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included in the case of A4, but not in C4. Activity coefficients were estimated using the
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simple Specific Interaction Theory (SIT) as described from Puxty and Maeder20. The
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regression of interaction parameters between ion pairs was found to not improve the model
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fitting results and as such were not included in the SIT based calculation of activity
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coefficients.
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In addition to the chemistry model, the Henry’s coefficient of CO2 in solution is required
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to relate the free aqueous CO2 concentration to CO2 partial pressure. In lieu of experimental
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data for these specific aqueous amines, the value for water was used which has previously
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been shown to be a reasonable approach20.
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With established knowledge of the equilibrium constants for the series of CO2 hydration
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reactions available from the literature21, and the corresponding amine reactions measured or
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estimated in this work, the chemical speciation can be determined by solving the system of
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nonlinear simultaneous equations consisting of the equilibrium constants and mass balance
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equations for the concentrations of all species using the Newton-Raphson method22. From the
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calculated speciation, the free aqueous CO2 concentration in solution was determined, and
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then the CO2 partial pressure above the solution can be calculated using the Henry’s
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coefficient. Conversely, equilibrium constants for the di-amine reactions in Eq. (4) - (9) can
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be determined from regression of the experimental CO2 partial pressures and known total
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concentrations from the VLE measurements. The Newton-Gauss-Levenberg/Marquardt
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method was used as described from Puxty et al.23. The objective function to be minimised in
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the regression was the sum of the squared relative error (SSQRE) between calculated and
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experimental CO2 partial pressure:
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exp cal pCO -pCO 2 ,i 2 ,i SSQRE = ∑ exp pCO 2 ,i i =1 n
n
2
(16)
exp is the total number of data points, pCO2 ,i is the experimental CO2 partial pressure
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where
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cal and pCO2 ,i is the calculated CO2 partial pressure from the model. The relative error, instead of
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the absolute error, was used since the data varies over orders of magnitude. The agreement
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between calculated and experimental data has been expressed as the average absolute relative
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deviation (AARD):
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exp cal 1 n pCO2 ,i -pCO2 ,i AARD = ∑ × 100% exp n i =1 pCO 2 ,i
(17)
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The temperature dependence of the amine reaction equilibrium constants was described
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using the form of the Van’t Hoff equation of Eq. (18). The regressed parameters for each
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equilibrium reaction were the zero ionic strength reference state equilibrium constant at
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0 313.15 K, lg K j ,313.15K , and the reaction enthalpy, ∆Hj .
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lg K j = lg K 0j ,313.15K −
∆H j
1 1 ⋅ − 2.3026 R T 313.15
(18)
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4. Results and discussion
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4.1 Vapour-Liquid-Equilibrium (VLE)
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CO2 solubilities in 30% MEA solutions from 313.15 to 393.15 K were initially
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evaluated to validate the experimental VLE setup for measurements of CO2 solubilities.
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Experimental CO2 partial pressure data as a function of CO2 loading is available in Table S1
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in Supporting Information. The experimental VLE data together with that from literature24, 25
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for comparison are shown in Figure 2. Good agreement was observed with the data from
229
literature at different temperatures, hence confirming the validity and suitability of the VLE
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setup for measurements of CO2 solubilities.
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CO2 solubilities in 15% and 30% aqueous solutions of the designer di-amines A4 and C4
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were measured at temperatures of 313.15, 333.15, 363.15, and 393.15 K. CO2 partial
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pressures as a function of CO2 loading in solutions of A4 and C4 in various concentrations
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are available in Tables S2 - S5 in Supporting Information. Graphical representations of CO2
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solubilities of A4 and C4 solutions are shown in Figure 3 with fitted trends from the
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modelling results.
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Cyclic loading and cyclic capacity were evaluated to assess the capability for CO2
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capture between different solutions. Cyclic loading, in terms of mole CO2 per mole amine,
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represents the difference between CO2 loadings after absorption and after desorption. Cyclic
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capacity, in terms of mole CO2 per kilogram solution, describes the total amount of CO2
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removed by regenerating one kilogram of solution in desorption process for one cycle. 8
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According to the analysis results in our previous work10, the cyclic loading and cyclic
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capacity of different solutions in an absorber-stripper cycle can be approximately evaluated
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by the equilibrium solubility of CO2 corresponding with the simulation conditions
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(absorption at 15 kPa, 313.15 K and desorption at 15 kPa, 393.15 K). A comparison of the
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results is shown in Figure 4.
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For the designer di-amines, the CO2 loadings in absorption and desorption of 15%
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solution are marginally higher than those of 30% solution, demonstrating that a lower amine
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concentration could provide a better opportunity for the reactive amine groups to contact with
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free CO2 in dilute solution. Cyclic loading of the 15% solution is similar to that of 30%
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solution, while the cyclic capacity of 15% solution is only half that of 30% solution as a
252
result of lower molality. Overall, higher amine concentrations would be favoured.
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Considering the two designer di-amines, CO2 loadings after absorption and desorption in
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A4 are both higher than those in C4 at same concentration, indicating that A4 possesses a
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greater affinity for CO2 than C4 (assuming physical CO2 absorption does not play a
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significant role in the absorption capacity). This result is consistent with the molecular
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structure characteristics that there are two secondary amine groups in the A4 molecule, while
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there is one primary and one tertiary amine group in the C4 molecule. Primary and secondary
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amine groups are able to directly react with CO2 while tertiary amine cannot. Similarly, both
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cyclic loading and cyclic capacity of A4 are marginally higher than C4, making A4 the more
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promising molecule.
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MEA was used as a comparative example to evaluate the capture performance of these
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two designer di-amines. With the same mass fraction, it can be seen that cyclic loadings of
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A4 and C4 are both at least twice as high as that of MEA, which is likely related to the high
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CO2 loading after absorption (74% and 52% higher than that of MEA for A4 and C4
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respectively). The CO2 loading after desorptionis a different situation with A4 17% higher
267
and C4 37% lower than that of MEA. However, the cyclic capacity of A4 and C4 are 14%
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and 12% lower than that of MEA respectively, resulting from the higher molecular weight
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and therefore lower molality of the designer di-amines when at equal mass fraction to MEA.
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A straight chain di-amine solution, N-(2-aminoethyl)ethanolamine26 (AEEA, 30%) and a
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straight chain tri-amine solution, diethylenetriamine17 (DETA, 20%) were used to investigate
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how the molecular structure affects the CO2 capture performance of amines. The effective
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average cyclic loadings assigned to each amine group of the designer di-amines (0.357 and
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0.349 for 30% A4 and C4 respectively) are higher than those of AEEA and DETA (0.287 and
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0.280 for 30% AEEA and 20% DETA respectively), indicating that the designer di-amines 9
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utilize each reactive amine group more effectively. Among all these amines, both the
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designer di-amines possess a lower CO2 loading at high temperature, enhancing the high
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cyclic loading. It was postulated that this phenomenon could be due to the steric hindrance
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and electron distribution introduced by the ring in the molecule similar to piperidine and
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piperazine. It is understandable that the sterical hindrance of A4 makes its 4-carbamate
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unstable, hence benefiting bicarbonate formation. Additionally, the tertiary amine group of
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C4 also promotes bicarbonate formation. AEEA has a similar chemical structure to the
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designer di-amines, however, its arrangement is one of a straight chain instead of the ring.
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The CO2 loading after absorption in AEEA is approaching those of A4 and C4, however, the
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CO2 loading after desorption is significantly higher. Thus, assumptions surrounding the
286
importance of the ring structure in improving the desorption performance of amines were
287
judged to be plausible. DETA has an extra primary amine group, replacing the hydroxyl
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group in AEEA, and still has a high cyclic capacity despite the modification. The hydroxyl
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group in the designer di-amines was initially designed to increase the boiling point of the
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amine, however, the capture performance from DETA indicates that the hydroxyl group in
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the designer di-amines could be replaced by another amine group to further improve the
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capture performance.
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4.2 CO2 absorption enthalpies
294
Differential and integral CO2 absorption enthalpies have been determined using a
295
CPA201 reaction calorimeter. The differential heat of CO2 absorption was determined by
296
dividing the heat released from each addition of CO2 by the amount of CO2 absorbed with
297
each addition. The integral heat of absorption was determined by dividing the total amount of
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heat released by the total amount of CO2 absorbed up to the equilibrium point. With
299
increasing CO2 loading the amount of heat released following CO2 addition became smaller
300
which is consistent with the amines becoming saturated with CO2.
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CO2 absorption enthalpies in 30% MEA solutions at 313.15 and 353.15 K were initially
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determined to validate the experimental enthalpy setup. A summary of differential and
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integral enthalpies is available in Table S6 in Supporting Information. The experimental
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enthalpies comparing the data from literature27-29 in MEA solutions are shown in Figure 5.
305
High consistency for heat curves was observed, from the figures, thus confirming reliability
306
and validity for measurements of CO2 absorption enthalpies.
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CO2 absorption enthalpies in 15% and 30% aqueous solutions of A4 and C4 were
308
measured at 313.15 K and 333.15 K. The measured differential enthalpies of CO2 absorption 10
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at various CO2 loading, as well as the integral values by integration, are given in Table S7
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and S8 in Supporting Information. Graphical representations are shown in Figure 6. There
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was obvious similarity with MEA in the CO2 absorption enthalpy for the designer di-amines,
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where CO2 absorption enthalpies decreased with increasing CO2 loading. However, unlike
313
MEA there was an approximately linear decrease for the integral enthalpies at the full range
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of CO2 loadings, while the integral enthalpies of MEA remained stable before reaching the
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critical loading point of 0.5. This indicates the ongoing reactions in solutions of the designer
316
di-amines include more than one dominant reaction. According to the previous work in which
317
the reaction products in designer di-amine solutions with CO2 were determined13, carbamate
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was the primary product at low CO2 loading, and carbonate/bicarbonate with increasing CO2
319
loading. As a result, the reaction enthalpies should be composed of the formation enthalpies
320
of carbamate and carbonate/bicarbonate, and the former is higher than the latter. This
321
phenomenon indicates the reaction mechanism of designer di-amines and CO2 at low loading
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is similar to the mechanism of primary amines with CO2, however, the formation of
323
carbonate/bicarbonate should be taken into account as CO2 loading increases.
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4.3 Amine protonation constants
325
As both amines are di-amines two protonation constants were determined for each amine
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at each measurement temperature. The results from potentiometric titration and linear
327
regression are tabulated in Table 2.
328
4.4 Modelling regression
329
95 and 92 points of experimental VLE data, covering an expansive range of CO2
330
loadings and temperatures, were used to regress equilibrium constants for the chemical model
331
of A4 and C4 respectively. Thus any chemical model regressed by this data set can be
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broadly considered to be representative of the chemical and physical behaviour relevant to
333
the CO2 absorption process with these di-amines.
334
As the CO2-H2O chemistry has been well characterised, the parameters for Eq. (1) to (3)
335
were fixed at the values used in literature21. The parameters for the protonation reactions of
336
Eq. (4) and (5) were fixed at the values given in Table 2. The regressed parameters were the
337
equilibrium constants at zero ionic strength and infinite dilution, and enthalpies for carbamate
338
formation, Eq. (6) and (7) and carbamate protonation, Eq. (8) and (9). It was found that the
339
parameters of the di-carbamate formation, Eq. (7) and di-carbamate protonation, Eq. (9),
340
were not defined by the data sets. This indicates that these reactions did not occur to a
341
significant extent and could be excluded from the model. This was expected for amine C4, as 11
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one amine is tertiary, however for amine A4, it was not previously known if di-carbamate
343
formation would be significant. Thus, the resulting chemical models only included Eq. (6)
344
and (8) resulting in 4 regressed parameters. Inclusion of interaction parameters between ionic
345
species did not result in an improved fit (reduced AARD) and therefore, were not included in
346
the model (that is the SIT model was used without interaction parameters). The regressed
347
parameter values are also given in Table 2. The parameter of particular interest is the stability
348
0 constant of carbamate formation, lg K 6,313.15K . The value for A4 is smaller than C4,
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indicating a less stable carbamate. This is consistent with the speciation determined by NMR
350
in Yang, et al.13 which showed significantly more carbamate formation for C4. Also the
351
enthalpy of protonation for the C4 carbamate is smaller than the A4 carbamate. This is
352
consistent with protonation of a tertiary amine group rather than a primary/secondary amine
353
group18.
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A parity plot of CO2 partial pressures is shown in Figure 7 with a random error
355
distribution. The model calculated VLE values using the regressed parameters are shown in
356
Figure 3. As can be seen, the AARD of the model for C4 was 11.9%, which was good enough
357
to represent the behaviour in the absorption process and showed agreement across the full
358
dataset. Meanwhile, the AARD for A4 was 17.3%, which is larger but still acceptable for a
359
simple model.
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The model calculated integral enthalpies are shown in Figure 6. The modelling results
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for A4 show good agreement with the experimental results below a CO2 loading of 0.5 and
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the agreement is good for C4 below a loading of 0.8. Above these loadings the model
363
systematically over estimates the absorption enthalpy with a largest deviation of ~20 kJ/mol
364
CO2. The model also predicts that the enthalpy profiles for A4 cross for both temperatures. It
365
is unclear why this is the case, however it may be related to the slow mass transfer at high
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CO2 loading resulting in small heat flows over extended time periods that are difficult to
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measure. Additionally, uncertainty in the Henry coefficient may also be a factor as it has a
368
larger influence upon the estimated enthalpy as the CO2 loading increases. In order to
369
improve the model, the measured enthalpy data could be used in the regression of the
370
equilibrium constants and reaction enthalpies. However, combining such disparate data sets
371
in a model regression is a challenging task requiring complex statistical analysis beyond the
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scope of this work.
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The calculated results of VLE data were in good agreement with the experimental data,
374
and the predicted absorption enthalpies were well match to measured data at low CO2
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loadings. These two designer di-amines were found to have beneficial performance
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characteristics for CO2 capture from flue gases. Relevant research of their properties at high
377
CO2 loading is still in progress.
378 379 380 381
Supporting Information Nomenclature, CO2 solubility data in Table S1 - S5, CO2 reaction enthalpy data in Table S6 - S8.
382 383
Acknowledgements
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Financial supports from National Science and Technology Support Program of China
385
(No.2015BAC04B01), the National Natural Science Foundation of China (No.51134017), are
386
greatly appreciated. Financial supports from Australia-China Joint Coordination Group (JCG)
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Partnership Fund is also appreciated. Huaneng Clean Energy Research Institute is gratefully
388
acknowledged for providing access to their experimental instruments.
389 390
Reference
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1. Leung, D. Y. C.; Caramanna, G.; Maroto-Valer, M. M., An overview of current status of carbon dioxide capture and storage technologies. Renewable Sustainable Energy Rev. 2014, 39, 426-443. 2. MacDowell, N.; Florin, N.; Buchard, A.; Hallett, J.; Galindo, A.; Jackson, G.; Adjiman, C. S.; Williams, C. K.; Shah, N.; Fennell, P., An overview of CO2 capture technologies. Energy Environ. Sci. 2010, 3, (11), 1645-1669. 3. Rochelle, G. T., Amine scrubbing for CO2 capture. Science 2009, 325, (5948), 1652-1654. 4. Ma’mun, S.; Svendsen, H. F.; Hoff, K. A.; Juliussen, O., Selection of new absorbents for carbon dioxide capture. Energy Convers. Manage. 2007, 48, (1), 251-258. 5. Zhang, X.; Fu, K.; Liang, Z.; Rongwong, W.; Yang, Z.; Idem, R.; Tontiwachwuthikul, P., Experimental studies of regeneration heat duty for CO2 desorption from diethylenetriamine (DETA) solution in a stripper column packed with Dixon ring random packing. Fuel 2014, 136, 261-267. 6. Conway, W.; Wang, X.; Fernandes, D.; Burns, R.; Lawrance, G.; Puxty, G.; Maeder, M., 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), 7422-7429. 7. Robinson, K.; McCluskey, A.; Attalla, M. I., An ATR-FTIR study on the effect of molecular structural variations on the CO2 absorption characteristics of heterocyclic amines, part II. Chemphyschem 2012, 13, (9), 2331-41. 8. Rochelle, G.; Chen, E.; Freeman, S.; Van Wagener, D.; Xu, Q.; Voice, A., Aqueous piperazine as the new standard for CO2 capture technology. Chem. Eng. J. 2011, 171, (3), 725-733. 9. Li, H.; Li, L.; Nguyen, T.; Rochelle, G. T.; Chen, J., Characterization of piperazine/2aminomethylpropanol for carbon dioxide capture. Energy Procedia 2013, 37, 340-352. 13
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416 417 418 419 420 421 422 423 424 425 426 427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455 456 457 458 459 460 461 462
10. Li, H.; Moullec, Y. L.; Lu, J.; Chen, J.; Marcos, J. C. V.; Chen, G., Solubility and energy analysis for CO2 absorption in piperazine derivatives and their mixtures. Int. J. Greenh. Gas Con. 2014, 31, 25-32. 11. Li, H.; Le Moullec, Y.; Lu, J.; Chen, J.; Valle Marcos, J. C.; Chen, G.; Chopin, F., CO2 solubility measurement and thermodynamic modeling for 1-methylpiperazine/water/CO2. Fluid Phase Equilib. 2015, 394, 118-128. 12. Conway, W.; Yang, Q.; James, S.; Wei, C.-C.; Bown, M.; Feron, P.; Puxty, G., Designer amines for post combustion CO2 capture processes. Energy Procedia 2014, 63, 1827-1834. 13. Yang, Q.; Puxty, G.; James, S.; Bown, M.; Feron, P.; Conway, W., Toward intelligent CO2 capture solvent design through experimental solvent development and amine synthesis. Energy Fuels 2016, 30, (9), 7503-7510. 14. Puxty, G.; Rowland, R.; Allport, A.; Yang, Q.; Bown, M.; Burns, R.; Maeder, M.; Attalla, M., Carbon dioxide postcombustion capture: A novel screening study of the carbon dioxide absorption performance of 76 amines. Environ. Sci. Technol. 2009, 43, (16), 6427-6433. 15. Yang, Q.; James, S. N.; Ballard, M. J.; Bown, M.; Feron, P.; Puxty, G. D. Gas capture process. WO2012142668A1, 2012. 16. Peng, D.-Y.; Robinson, D. B., A new two-constant equation of state. Ind. Eng. Chem. Fundam. 1976, 15, (1), 59-64. 17. Luo, W.; Guo, D.; Zheng, J.; Gao, S.; Chen, J., CO2 absorption using biphasic solvent: Blends of diethylenetriamine, sulfolane, and water. Int. J. Greenh. Gas Con. 2016, 53, 141148. 18. Fernandes, D.; Conway, W.; Wang, X.; Burns, R.; Lawrance, G.; Maeder, M.; Puxty, G., Protonation constants and thermodynamic properties of amines for post combustion capture of CO2. J. Chem. Thermodyn. 2012, 51, 97-102. 19. Gans, P.; O'Sullivan, B., GLEE, a new computer program for glass electrode calibration. Talanta 2000, 51, (1), 33-37. 20. Puxty, G.; Maeder, M., A simple chemical model to represent CO2-amine-H2O vapourliquid-equilibria. Int. J. Greenh. Gas Con. 2013, 17, 215-224. 21. Edwards, T. J.; Maurer, G.; Newman, J.; Prausnitz, J. M., Vapor-liquid equilibria in multicomponent aqueous solutions of volatile weak electrolytes. AlChE J. 1978, 24, (6), 966976. 22. Maeder, M.; Neuhold, Y.-M., Practical data analysis in chemistry. Elsevier Science: 2007; Vol. 26, p 350. 23. Puxty, G.; Maeder, M.; Hungerbühler, K., Tutorial on the fitting of kinetics models to multivariate spectroscopic measurements with non-linear least-squares regression. Chemometrics Intellig. Lab. Syst. 2006, 81, (2), 149-164. 24. Lee, J. I.; Otto, F. D.; Mather, A. E., Equilibrium between carbon dioxide and aqueous monoethanolamine solutions. J. Appl. Chem. Biotechnol. 1976, 26, (1), 541-549. 25. Jou, F.-Y.; Mather, A. E.; Otto, F. D., The solubility of CO2 in a 30 mass percent monoethanolamine solution. Can. J. Chem. Eng. 1995, 73, (1), 140-147. 26. Ma'mun, S.; Jakobsen, J. P.; Svendsen, H. F.; Juliussen, O., Experimental and modeling study of the solubility of carbon dioxide in aqueous 30 Mass % 2-((2aminoethyl)amino)ethanol solution. Ind. Eng. Chem. Res. 2006, 45, (8), 2505-2512. 27. Kim, I.; Svendsen, H. F., Heat of absorption of carbon dioxide (CO2) in monoethanolamine (MEA) and 2-(aminoethyl)ethanolamine (AEEA) solutions. Ind. Eng. Chem. Res. 2007, 46, (17), 5803-5809. 28. Lee, J. I.; Otto, F. D.; Mather, A. E., The solubility of H2S and CO2 in aqueous 14
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monoethanolamine solutions. Can. J. Chem. Eng. 1974, 52, (6), 803-805. 29. Jou, F.-Y.; Otto, F. D.; Mather, A. E., Vapor-Liquid Equilibrium of Carbon Dioxide in Aqueous Mixtures of Monoethanolamine and Methyldiethanolamine. Ind. Eng. Chem. Res. 1994, 33, (8), 2002-2005.
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Table 1. The structures of di-amines A4 and C4. Di-amines
Structure
MW (g.mol-1)
A4
144.22
C4
144.22
470
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Table 2. Amine protonation constants at zero ionic strength as determined from potentiometric titration and
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the values of
lg K 0j ,313.15K and ∆Hj from Eq. (4), (5), (6), and (8) determined by regression. Experimental
Regression
lg K 0j ,313.15K
∆H j (kJ/mol)
Eq.
288.15 K
298.15 K
308.15 K
318.15 K
(4)
10.56
10.39
10.11
9.95
10.04±0.02
37.0±2.9
(5)
7.42
7.24
7.04
6.86
6.95±0.01
33.0±0.9
(6)
-
-
-
-
-0.01±0.36
-16.2±12.4
(8)
-
-
-
-
9.95±0.32
-66.7±20.4
A4
AARD
17.3%
(4)
10.04
9.73
9.48
9.15
9.31±0.02
51.2±2.4
(5)
7.07
6.89
6.71
6.55
6.63±0.01
30.5±0.3
(6)
-
-
-
-
0.62±0.06
-14.8±2.7
(8)
-
-
-
-
8.86±0.05
-24.0±3.0
11.9%
C4
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Figure Captions
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Figure 1. Schematic diagram of (a) the thermostatic VLE stirred cell, (b) the CPA201 reaction calorimeter.
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Figure 2. Comparison of CO2 solubilities in 30% MEA solutions from this work (Solid) and with literature
477
data (hollow) at temperatures of (a) 313.15 and 373.15 K, (b) 333.15 and 393.15 K.
478
Figure 3. Experimental data (points) and model calculated data (lines) of CO2 solubilities in A4 and C4
479
solutions from 313.15 to 393.15 K.
480
Figure 4. Equilibrium loadings, cyclic loadings, average cyclic loadings, cyclic capacities in various amine
481
solutions.
482
Figure 5. Differential and integral CO2 absorption enthalpies in 30% MEA solutions with comparison of
483
literature data.
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Figure 6. Differential CO2 absorption enthalpies in 15% and 30% A4 and C4 solutions.
485
Figure 7. The parity plot of model calculated versus measured CO2 partial pressures.
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Figure 1.
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Figure 2.
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493 494
Figure 3.
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496 497
Figure 4.
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Figure 5.
501
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503 504
Figure 6.
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Figure 7.
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