co capture using alkanolamine/room ......guidance of prof. muhammad mazhar and dr. syed tajammul...
TRANSCRIPT
CO2 CAPTURE USING
ALKANOLAMINE/ROOM-TEMPERATURE
IONIC LIQUID BLENDS Absorption, Regeneration, and Corrosion Aspects
Thèse
Muhammad Hasib-ur-Rahman
Doctorat en génie chimique
Philosophiae Doctor (Ph.D.)
Québec, Canada
© Muhammad Hasib-ur-Rahman, 2013
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Résumé
Le réchauffement climatique, résultant essentiellement des émissions anthropiques de
dioxyde de carbone, demeure un sujet de grande préoccupation. Le captage et la
séquestration du dioxyde de carbone est une solution viable permettant de prévoir une
baisse des émissions de CO2 issues des importantes sources ponctuelles qui impliquent la
combustion des carburants fossiles. Dans cette perspective, les systèmes aqueux
d‟alcanolamines offrent une solution prometteuse à court terme pour la capture du CO2
dans les installations de production d'électricité. Cependant, ces systèmes sont confrontés à
divers accrocs opératoires tels que les limitations d‟équilibre, les grandes quantités
d‟énergie requises pour la régénération, les pertes en solvant et la corrosion prononcée des
installations, pour ne citer que ces quelques inconvénients. L‟eau étant la principale cause
de ces complications, une mesure à prendre pourrait être le remplacement de la phase
aqueuse par un solvant plus stable.
Les liquides ioniques à température ambiante, dotés d‟une haute stabilité thermique et
pratiquement non-volatils émergent en tant que candidats prometteurs. De plus, grâce à leur
nature ajustable, ils peuvent être apprêtés conformément aux exigences du procédé. La
substitution de la phase aqueuse dans les processus utilisant l‟alcanolamine par les liquides
ioniques à température ambiante ouvre une opportunité potentielle pour une capture
efficace du CO2. Un aspect remarquable de ces systèmes serait la cristallisation du produit
résultant de la capture du CO2 (c-à-d, le carbamate) au sein même du liquide ionique qui
non seulement déjouerait les contraintes d‟équilibre mais également pourvoirait une
opportunité intéressante pour la séparation des produits.
Étant donné le peu d‟information disponible dans la littérature sur la viabilité des systèmes
utilisant la combinaison d‟amine et de liquide ionique, l‟étude proposée ici a pour but
d‟apporter une meilleure compréhension sur l‟efficacité à séparer le CO2 d‟un mélange de
type postcombustion à travers une approche plus systématique. À cet effet, des liquides
ioniques à base d‟imidazolium ([Cnmim][Tf2N], [Cnmim][BF4], [Cnmim][Otf]) ont été
choisis. Deux alcanolamines, à savoir, le 2-amino-2-methyl-1-propanol (AMP) et le
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diéthanolamine (DEA) ont été examinées en détail afin d‟explorer la capture du CO2 et les
possibilités de régénération qu‟offre un système amine-liquide ionique. Les résultats ont
révélé l‟intérêt de la combinaison DEA-liquide ionique étant donné que ce système pourrait
aider à réduire de manière significative l‟écart entre les températures d‟absorption et de
régénération, promettant ainsi une perspective attrayante en termes d‟économie d‟énergie.
En outre, les liquides ioniques ont également été scrutés du point de vue de leur nature
hydrophobe/hydrophile afin d‟étudier le comportement corrosif du mélange amine-liquide
ionique au contact d‟échantillons d‟acier au carbone. Bien que l‟utilisation des liquides
ioniques hydrophiles ait aidé à abaisser la vitesse de corrosion jusqu‟à concurrence de 72%,
l‟emploi de liquides ioniques hydrophobes s‟avère plus efficace, car annulant quasiment le
phénomène de corrosion même dans un environnement riche en CO2.
Dans le cas des mélanges immiscibles comme DEA-[hmim][Tf2N], une agitation continue
s‟avère nécessaire afin d‟assurer une dispersion prolongée des gouttelettes d‟amine
émulsifiées au sein de liquides ioniques et ainsi atteindre une vitesse de capture optimale.
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Abstract
Global warming, largely resulting from anthropogenic emissions of carbon dioxide,
continues to remain a matter of great concern. Carbon capture and storage (CCS) is a viable
solution to ensure a prevised fall in CO2 emissions from large point sources involving fossil
fuel combustion. In this context, aqueous alkanolamine systems offer a promising near-
term solution for CO2 capture from power generation facilities. However, these face several
operational hitches such as equilibrium limitations, high regeneration energy requirement,
solvent loss, and soaring corrosion occurrence. The main culprit in this respect is water and,
accordingly, one feasible practice may be the replacement of aqueous phase with some
stable solvent.
Room-temperature ionic liquids (RTILs), with high thermal stability and practically no
volatility, are emerging as promising aspirants. Moreover, owing to the tunable nature of
ionic liquids, RTIL phase can be adapted in accordance with the process requirements.
Replacing aqueous phase with RTIL in case of alkanolamine based processes provided a
potential opportunity for efficient CO2 capture. The most striking aspect of these schemes
was the crystallization of CO2-captured product (carbamate) inside the RTIL phase that not
only helped evade equilibrium constraints but also rendered a worthy opportunity of
product separation.
Since there is little information available in the literature about the viability of amine-RTIL
systems, the proposed research was aimed at better understanding CO2 separation
proficiency of these fluids through a more systematic approach. Imidazolium RTILs
([Cnmim][Tf2N], [Cnmim][BF4], [Cnmim][Otf]) were chosen for this purpose. Two
alkanolamines, 2-amino-2-methyl-1-propanol (AMP) and diethanolamine (DEA) were
examined in detail to explore CO2 capture and regeneration capabilities of amine-RTIL
systems. The results revealed the superiority of DEA-RTIL combination as this scheme
could help significantly narrow the gap between absorption and regeneration temperatures
thus promising a sparkling prospect of attenuating energy needs. Furthermore, ionic liquids
were scrutinized in reference to their hydrophobic/hydrophilic nature to study the corrosion
behaviour of carbon steel in amine-RTIL media. Though hydrophilic ionic liquids helped
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decrease corrosion occurrence up to 72%, hydrophobic RTIL appeared to be the most
effective in this regard, virtually negating the corrosion phenomenon under CO2 rich
environment.
In case of immiscible blends like DEA-[hmim][Tf2N], continual agitation appeared to be a
necessity to ensure a prolonged dispersion of amine in the RTIL phase and, thereby, to
attain an optimal capture rate.
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Foreword
This PhD thesis has been divided into five chapters. The first chapter comprises the
introductory portion and it also contains a short published review [1], merged and modified
in accordance with the context of the “INTRODUCTION and OBJECTIVES” section.
Immediately after, four research articles (listed below) follow, each presented as a separate
chapter (Chapters 2-5). Out of these, three research articles ([2] to [4]) were already
published while the last one [5] is under review for publication in Separation and
Purification Technology journal. At the end, as „Appendix D‟, a short essay about corrosion
perspective regarding amine-based CO2 capture systems (i.e. aqueous amines and amine-
ionic liquid blends) has been attached that we published in „Carbon Capture Journal‟ [6].
[1] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Ionic Liquids for CO2 Capture -
Development and Progress, Chem. Eng. Process. 49 (2010) 313-322.
[2] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 Capture in Alkanolamine/Room-
Temperature Ionic Liquid Emulsions: A Viable Approach with Carbamate Crystallization
and Curbed Corrosion Behavior, Int. J. Greenhouse Gas Control 6 (2012) 246-252.
[3] M. Hasib-ur-Rahman, H. Bouteldja, P. Fongarland, M. Siaj, F. Larachi, Corrosion
Behavior of Carbon Steel in Alkanolamine/Room-Temperature Ionic Liquid based CO2
Capture Systems, Ind. Eng. Chem. Res. 51 (2012) 8711-8718.
[4] M. Hasib-ur-Rahman, F. Larachi, CO2 Capture in Alkanolamine-RTIL Blends via
Carbamate Crystallization: Route to Efficient Regeneration, Environ. Sci. Technol. 46
(2012) 11443-11450.
[5] M. Hasib-ur-Rahman, F. Larachi, Kinetic Behavior of Carbon Dioxide Absorption in
Diethanolamine/Ionic-Liquid Emulsions, Sep. Purif. Technol. Submitted February 2013.
[6] M. Hasib-ur-Rahman, F. Larachi, Corrosion in amine systems – a review, Carbon
Capture Journal, Sept - Oct 2012, 22-24.
viii
The research papers were prepared on my own and revised by my director, Prof. Faïçal
Larachi. Prof. Larachi guided and provided expertise in designing experiments and
managing data analysis during the entire research work.
Prof. Mohamed Siaj, my co-director from Department of Chemistry at Université du
Québec à Montréal, facilitated through his productive recommendations on the
characterization of CO2-captured products (carbamates) and helped correct the manuscripts
of first three publications.
Ms. Hana Bouteldja, co-author of the paper entitled, “Corrosion Behavior of Carbon Steel
in Alkanolamine/Room-Temperature Ionic Liquid based CO2 Capture Systems”, helped
perform the corrosion experiments and was involved in configuring the Chittick technique
for CO2 loading measurements. In this regard, Prof. Pascal Fongarland (supervisor of Ms.
Hana Bouteldja, and also co-author of the published work) from Ecole Centrale de Lille,
Unité de Catalyse et Chimie du Solide, France, provided some fruitful suggestions.
Some of the research outcomes were presented in the following conferences:
M. Hasib-ur-Rahman, H. Bouteldja, A.N. Khan Wardag, A. Sarvaramini, G.P.
Assima, M. Siaj, F. Larachi, Advances towards adept biomass gasification and
efficient carbon dioxide capture processes, CQMF 4th Annual Symposium at
Duchesnay, Quebec, 2011.
M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 Capture in Alkanolamine/Room-
Temperature Ionic Liquid Emulsion System, 61st Canadian Chemical Engineering
Conference, London ON, 2011.
M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Corrosion inhibition in
alkanolamine/room-temperature ionic liquid based CO2 capture systems, CAMURE
8 & ISMR 7, Naantali, Finland, 2011.
M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Alkanolamine/Ionic Liquid
Microemulsions for Efficient CO2 Capture with Diminished Corrosion
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Phenomenon”, CQMF 3rd Annual Symposium at Centre d'arts Orford, Orford QC,
2010.
M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Alkanolamine/Ionic Liquid
Microemulsions: Enhanced CO2 Capture Ability with Curbed Corrosion Behaviour,
CIGR World Congress, Québec QC, 2010.
M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 capture by alkanolamine/ionic liquid
microemulsion system equipped with micro-fluidic channels detector for in-situ
screening of the process, CQMF 2nd Annual Symposium at UQÀM, Montréal QC,
2009.
xi
Acknowledgements
Firstly my sincere appreciation goes to my affectionate parents for their kind support and
encouragement during the whole of my life. Also, I am overwhelmed with gratitude for the
consistent support of my loving wife during the challenging times of my Ph.D. studies.
Thank you so much.
This thesis would not be in good shape without the inspirational attitude of my siblings who
wisely advised in the final stretch of my education.
I would like to express my heartfelt gratitude to my supervisor, Prof. Faïçal Larachi, for his
ample guidance and help during the course of my Ph.D. through his unique thought-
provoking, supportive, and composed approach. I hope that I can pass on the research
values that he has given to me.
My co-director, Prof. Mohamed Siaj, is gratefully acknowledged for his support and many
insightful suggestions during the project progression.
I express my gratitude to Prof. Denis Rodrigue and Prof. Maria-Cornélia Iliuta for letting
use their analytical facilities.
I would also like to thank my examiners, Dr. Sylvie Fradette, Prof. Alain Garnier, and Prof.
Louis Fradette, who provided constructive feedback. It is no easy task, reviewing a thesis,
and I am grateful for their thoughtful comments.
It had been a great privilege to spend some fruitful years of my M.Phil. studies under the
guidance of Prof. Muhammad Mazhar and Dr. Syed Tajammul Hussain, at Quaid-i-Azam
University in Pakistan, who enabled me to contemplate this road. I could not have asked for
better role models, each inspirational and supportive.
Thank you my friends and colleagues. You were the sources of laughter, joy, and
encouragement, bearing the brunt of frustrations and sharing the joy of successes. The help
of the chemical engineering department technical staff, Jérome Noël, Marc Lavoie, Yann
Giroux, and Jean-Nicolas Ouellet during this research project is also appreciated.
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Finally, I acknowledge the financial support of Fonds de recherche du Québec – Nature et
technologies (FRQNT), F. Larachi Canada Research Chair “Green processes for cleaner
and sustainable energy”, the Centre québécois sur les matériaux fonctionnels (CQMF), and
the Discovery Grants to F. Larachi and M. Siaj from the Natural Sciences and Engineering
Research Council (NSERC).
Thank you all
xiii
Table of contents
Résumé .............................................................................................................................................................. iii
Abstract ............................................................................................................................................................. v
Foreword .......................................................................................................................................................... vii
Acknowledgements ........................................................................................................................................ xi
List of figures ................................................................................................................................................ xvii
List of tables ................................................................................................................................................... xxi
Chapter 1: Introduction and Objectives ...................................................................................................... 1
1.1. Background .................................................................................................................................... 1
1.2. Carbon dioxide capture through solvent scrubbing ............................................................ 2
1.2.1. Chemical solvents ................................................................................................................ 2
1.2.2. Degradation of amines ........................................................................................................ 4
1.2.3. Corrosion of equipment ..................................................................................................... 8
1.2.4. Corrosion inhibition ............................................................................................................ 9
1.3. Physical Solvents......................................................................................................................... 10
1.4. Ionic Liquid Solvents ................................................................................................................. 11
1.5. Ionic liquids for CO2 capture - Development and progress ............................................. 11
1.5.1. Introduction ........................................................................................................................ 13
1.5.2. CO2 capture by room-temperature ionic liquids (RTILs) ....................................... 13
1.5.3. CO2 capture by task-specific ionic liquids (TSILs) ................................................... 21
1.5.4. CO2 capture by supported ionic-liquid membranes (SILMs) ................................. 23
1.5.5. CO2 capture by polymerized ionic liquids ................................................................... 27
1.5.6. Toxicity of ILs..................................................................................................................... 28
1.5.7. Current and future developments ................................................................................. 29
1.6. Research Objectives ................................................................................................................... 32
1.7. References .................................................................................................................................... 35
Chapter 2: CO2 capture in alkanolamine/room-temperature ionic liquid emulsions: A viable
approach with carbamate crystallization and curbed corrosion behavior .............................................. 45
2.1. Introduction ................................................................................................................................. 45
2.2. Experimental ............................................................................................................................... 47
2.2.1. Materials and techniques ................................................................................................. 47
2.2.2. Crystal structure determination .................................................................................... 48
xiv
2.2.3. Electrochemical corrosion tests ...................................................................................... 48
2.3. Results and discussion ............................................................................................................... 49
2.3.1. Fate of CO2-captured product (carbamate) ................................................................ 49
2.3.2. CO2 absorption ................................................................................................................... 50
2.3.3. Characterization of crystalline product ....................................................................... 53
2.3.4. Corrosion studies ............................................................................................................... 57
2.4. Conclusions .................................................................................................................................. 61
2.5. References..................................................................................................................................... 61
Chapter 3: Corrosion behaviour of carbon steel in alkanolamine/room-temperature ionic liquid
based CO2 capture systems ........................................................................................................................... 65
3.1. Introduction ................................................................................................................................. 66
3.2. Experimental ............................................................................................................................... 67
3.2.1. Materials .............................................................................................................................. 67
3.2.2. Experimental techniques and procedure ..................................................................... 68
3.3. Results and Discussion .............................................................................................................. 70
3.3.1. Effect of amine type on corrosion of steel .................................................................... 73
3.3.2. Effect of RTIL type on corrosion behaviour ............................................................... 76
3.3.3. Effect of process temperature ......................................................................................... 79
3.3.4. Effect of gas loading .......................................................................................................... 81
3.3.5. Presence of oxygen ............................................................................................................. 82
3.3.6. Influence of water .............................................................................................................. 84
3.4. Conclusion .................................................................................................................................... 85
3.5. References..................................................................................................................................... 86
Chapter 4: CO2 capture in alkanolamine-RTIL blends via carbamate crystallization: route to
efficient regeneration ......................................................................................................................... 89
4.1. Introduction ..................................................................................................................... 90
4.2. Experimental ............................................................................................................................... 93
4.2.1. Materials .............................................................................................................................. 93
4.2.2. Procedures and techniques .............................................................................................. 94
4.3. Results and Discussion .............................................................................................................. 95
4.3.1. Maximum gas capture capacity ...................................................................................... 95
4.3.2. Nature of CO2-captured products .................................................................................. 97
4.3.3. Regeneration ability .......................................................................................................... 98
4.3.4. Amine (AMP/DEA) regeneration behavior ............................................................... 102
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4.4. Implications ............................................................................................................................... 106
4.5. References .................................................................................................................................. 107
Chapter 5: Kinetic behavior of carbon dioxide absorption in diethanolamine/ionic-liquid emulsions
........................................................................................................................................................................ 111
5.1. Introduction ............................................................................................................................... 112
5.2. Reaction mechanism in non-aqueous amines .................................................................... 113
5.3. Experimental ............................................................................................................................. 114
5.3.1. Materials ............................................................................................................................ 114
5.3.2. Setup ................................................................................................................................... 114
5.3.3. Procedure ........................................................................................................................... 115
5.4. Results and Discussion ............................................................................................................ 116
5.4.1. Impact of variation in amine concentration .............................................................. 117
5.4.2. CO2 volume ratio in the gaseous mixture .................................................................. 119
5.4.3. Influence of agitation speed ........................................................................................... 121
5.4.4. Effect of temperature variation .................................................................................... 122
5.5. Conclusion .................................................................................................................................. 123
5.6. References .................................................................................................................................. 124
Chapter 6: Conclusions and recommendations ...................................................................................... 129
6.1. General conclusions ................................................................................................................. 129
6.2. Future work recommendations ............................................................................................. 131
Appendix A: Supporting Information (Chapter 2) ................................................................................. 133
Appendix B: Supporting Information (Chapter 3) ................................................................................. 143
Appendix C: Supporting Information (Chapter 4) ................................................................................. 149
Appendix D: Corrosion in amine systems – a review ........................................................................... 159
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List of figures
Figure 1. 1. Reaction scheme including reactions between amine, CO2/carbonate and protons. ....... 3
Figure 1. 2. Oxidative degradation mechanism of aqueous MEA. .................................................... 4
Figure 1. 3. Amine degradation: ■ thermal; ■ CO2 induced. ............................................................. 6
Figure 1. 4. Effect of SO2 on MEA degradation. ............................................................................... 7
Figure 1. 5. Effect of corrosion inhibitor (NaVO3) on MEA degradation. ......................................... 7
Figure 1. 6. Effect of various parameters on the corrosion rate of carbon steel C1020 in aqueous
MEA (basal conditions: MEA conc. 5 kmol/m3; gas loading 0.4 mol CO2/mol MEA; 80 °C
temperature). ....................................................................................................................................... 9
Figure 1. 7. Some cations and anions constituting ionic liquids (ILs). ............................................ 13
Figure 1. 8. Solubilities of CO2, C2H4, C2H6, CH4, Ar and O2 in [bmim][PF6] at 25 °C.................. 14
Figure 1. 9. CO2 solubility in [emim][Tf2N] and [emim][PF6]. ....................................................... 17
Figure 1. 10. Proposed mechanism for chemical absorption of CO2 by the TSIL. .......................... 19
Figure 1. 11. [hmim][Tf2N]-MEA solution: (a) fresh sample; (b) on CO2 exposure; showing
precipitated MEA-carbamate. ........................................................................................................... 20
Figure 1. 12. Proposed mechanism for CO2 capture by [pabim][BF4]. ............................................ 21
Figure 1. 13. Molar CO2 loads in solvent volume (for MEA/MDEA, consider aqueous solution
volume): data for ionic liquids at 30 °C [50]; data for MEA and MDEA at 40 °C........................... 23
Figure 1. 14. Proposed setup for CO2 separation by SILM in a coal-fired power plant. ................. 24
Figure 1. 15. Proposed mechanisms of CO2 capture: (a, b) without water; (c) with water. ............. 25
Figure 2. 1. DEA/RTIL system: (a-c) (without surfactant) after CO2 capture; d) (with surfactant)
before and after CO2 capture. ............................................................................................................ 50
Figure 2. 2. CO2 capture capacity profiles of DEA/RTIL system (surfactant stabilized emulsions;
30% w/w) at atmospheric pressure and 25 °C. ................................................................................. 51
Figure 2. 3. CO2 absorption isotherms for DEA/[hmim][Tf2N] surfactant stabilized emulsions
obtained at 25°C. ............................................................................................................................... 52
Figure 2. 4. Basic structural unit in DEA-carbamate (C9H22N2O6) crystal. ..................................... 54
Figure 2. 5. Hydrogen bonding pattern in the compound (DEA-carbamate). H atoms not
participating in hydrogen bonding are omitted for clarity. ............................................................... 55
Figure 2. 6. 13
C NMR spectrum of crystalline carbamate (retaining traces of [hmim][Tf2N]) taken in
DMSO-d6 solvent. ............................................................................................................................ 56
Figure 2. 7. FTIR analysis of crystalline product (DEA-carbamate). .............................................. 57
xviii
Figure 2. 8. Tafel plots for carbon steel electrode in aqueous DEA under different environments: a)
CO2 bubbling at 25 °C, b) CO2+O2 bubbling at 25 °C, c) CO2 bubbling at 60 °C, d) CO2+O2
bubbling at 60 °C. .............................................................................................................................. 58
Figure 2. 9. Corrosion rate of carbon steel 1020 in: a) RTIL pure, CO2+O2+H2O(vap.) bubbling at 60
°C, b) DEA/RTIL emulsion, CO2+O2+H2O(vap.) bubbling at 60 °C, c) DEA(aq), CO2+O2 bubbling at
60 °C. ................................................................................................................................................. 60
Figure 2. 10. SEM micrographs of working electrode specimen. In DEA/RTIL emulsion (15%
w/w): a) Fresh surface; b) after electrochemical corrosion test. In DEAaq. (15% w/w): c) Fresh
surface; d) after electrochemical corrosion test. ................................................................................ 60
Figure 3. 1. Experimental setup for electrochemical corrosion tests. ............................................... 69
Figure 3. 2. MEA-RTIL fluid showing solid carbamate, after CO2 bubbling: 1) MEA+[bmim][BF4];
2) MEA+[emim][BF4]; 3) MEA+[emim][Otf]. ................................................................................. 71
Figure 3. 3. Thermogravimetric evolution of CO2 absorption for MEA-RTIL mixtures (MEA: 5
kmol/m3) at 25 °C. ............................................................................................................................. 72
Figure 3. 4. SEM micrographs of steel electrode surface before and after electrochemical
polarization runs at 25 °C under CO2(15%)+O2(5%)+N2 atmosphere in: a) MEA (aqueous); b)
MEA+[bmim][BF4]; c) MEA+Water+[bmim][BF4]. ........................................................................ 73
Figure 3. 5. Linear polarization curves of carbon steel 1020 at 25 °C: a) in aqueous alkanolamines;
b) in alkanolamine+[bmim][BF4] mixtures. ...................................................................................... 74
Figure 3. 6. Effect of RTIL type on polarization behavior of carbon steel 1020 at 25 °C. .............. 76
Figure 3. 7. EDX analysis of steel electrode surface: a) freshly polished surface; b,c,d) after
electrochemical corrosion tests in MEA+[bmim][BF4], MEA+[emim][BF4], MEA+[emim][Otf]
blends respectively. ........................................................................................................................... 79
Figure 3. 8. Comparison of temperature effect on steel corrosion in aqueous as well as RTIL based
media. ................................................................................................................................................ 80
Figure 3. 9. EDX scan of steel electrode surface after electrochemical corrosion test in
MEA+[bmim][BF4] at 60 °C. ............................................................................................................ 80
Figure 3. 10. CO2 loading effect on steel corrosion at 25 °C in a) aqueous MEA b)
MEA+[bmim][BF4] mixture. ............................................................................................................. 82
Figure 3. 11. Effect of O2 concentration in flue gas on corrosion of steel a) in aqueous MEA; b) in
MEA+[bmim][BF4] mixture. ............................................................................................................. 83
Figure 3. 12. Effect of water content in CO2 capture medium on corrosion of steel. ....................... 84
Figure 4. 1. The simplified process flow diagram of alkanolamine-RTIL based CO2 capture
process. .............................................................................................................................................. 93
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Figure 4. 2. CO2 absorption isotherm for alkanolamine-[hmim][Tf2N] systems obtained at
atmospheric pressure and 35 °C temperature. ................................................................................... 96
Figure 4. 3. Evaporation profiles of amines (in amine-RTIL blends) at 35 °C under N2. ................ 96
Figure 4. 4. Packing diagrams: a) AMP-carbamate; b) DEA-carbamate. ........................................ 97
Figure 4. 5. a) FTIR spectra, and b) 13
C NMR spectra: AMP (fresh amine), AMPC (AMP-
carbamate) and RAMP (regenerated AMP). ................................................................................... 100
Figure 4. 6. a) FTIR spectra, and b) 13
C NMR spectra of DEA (fresh amine), DEAC (DEA-
carbamate) and RDEA (regenerated DEA). .................................................................................... 101
Figure 4. 7. DSC/TG profiles of AMP-carbamate: Thermal behavior observed under N2 atmosphere
at heating rate of 5 °C/min. ............................................................................................................. 102
Figure 4. 8. DSC/TG curves of DEA-carbamate: Thermal behavior under N2 atmosphere, using
heating rate of 5 °C/min. ................................................................................................................. 103
Figure 4. 9. QMS monitoring of carbamates‟ decomposition by measuring positive ion current m/z
= 44 (CO2) under N2 atmosphere (100 mL/min. flow rate) at 5 °C/min heating rate. .................... 104
Figure 4. 10. TG profiles of carbamates: Thermal behavior under CO2 atmosphere, using heating
rate of 5 °C/min. .............................................................................................................................. 106
Figure 5. 1. Experimental set-up scheme: A) Gas inlet; B) Gas outlet (A & B connect to a gas
reservoir via closed loop system); C) CO2 probe; D) Injection port; E) Thermocouple; F) Rotor-
stator homogeniser; G) Absorption cell; Hi) Heating bath inlet; Ho) Heating bath outlet. ............. 115
Figure 5. 2. CO2-captured product (carbamate) precipitation in DEA-[hmim][Tf2N]: a) immediately
after CO2 bubbling; b) 24 hours later. ............................................................................................. 116
Figure 5. 3. Influence of [DEA] molar concentration on absorption rate with respect to initial CO2
vol% in the gaseous mixture, at 33 °C and 3000 rpm agitation speed: a) 2M DEA in [hmim][Tf2N];
b) 1M DEA in [hmim][Tf2N]; c) 0.5M DEA in [hmim][Tf2N]. Smoothed lines show trends. ....... 119
Figure 5. 4. Influence of initial CO2 volume ratio (in gaseous mixture) on absorption rate w.r.t.
[DEA], at 33 °C and 3000 rpm agitation speed: a) 10 vol% CO2; b) 5 vol% CO2; c) 2.5 vol% CO2.
Smoothed lines show trends. ........................................................................................................... 121
Figure 5. 5. Influence of agitation on CO2 absorption rate (2M DEA in [hmim][Tf2N]; 10 vol%
CO2; 33 °C). Smoothed lines show trends. ..................................................................................... 122
Figure 5. 6. Effect of temperature on CO2 capture rate (1M DEA in [hmim][Tf2N]; 10 vol% CO2;
3000 rpm). Smoothed lines show trends. ........................................................................................ 123
xxi
List of tables
Table 1. 1. Initial rates of oxidative degradation of MEA under different operating conditions. ...... 5
Table 1. 2. Toxicity of absorption solvents and conventional inhibitors. ........................................ 10
Table 1. 3. Physical solvent processes. ............................................................................................. 11
Table 1. 4. Henry‟s constants (bar, at 25 °C) for gases in different organic solvents. ..................... 15
Table 1. 5. Henry‟s constants for CO2 in different ionic liquids. ..................................................... 15
Table 1. 6. Henry‟s law constants of CO2 in ionic liquids. .............................................................. 16
Table 1. 7. CO2 solubility data in [emim][MDEGSO4]. ................................................................... 18
Table 1. 8. Viscosity values for different compositions of tri-iso-butyl(methyl)phosphonium
tosylate/water mixtures. .................................................................................................................... 19
Table 1. 9. Viscosities and water content of the ionic liquids, at 25 °C. .......................................... 26
Table 1. 10. Summary of gas absorption capacities (at 592.3 mmHg & 22 °C) and glass transition
temperatures of poly(ionic liquid)s. .................................................................................................. 27
Table 1. 11. Permeability, solubility and diffusivity values in: a) styrene-based poly(ionic liquid)s;
b) acrylate-based poly(ionic liquid)s, at 20 °C. ................................................................................. 28
Table 1. 12. Lethal concentrations (LC50) of different ionic liquids to fresh water snail (Physa
acuta) in 96-hour acute toxicity exposures. ....................................................................................... 29
Table 1. 13. Summary of CO2 capture by ionic liquids. ................................................................... 31
Table 2. 1. Density (ρ) and viscosity (η) values measured at 25 °C. ................................................ 50
Table 2. 2. Crystallographic data ...................................................................................................... 53
Table 2. 3. Relevant hydrogen bonding parameters [bond distances (Å) and angles (°)]. ............... 54
Table 2. 4. Corrosion rates of carbon steel 1020 .............................................................................. 59
Table 3. 1. Summary of process parameters/conditions. .................................................................. 67
Table 3. 2. Viscosity values of the ionic liquids used. ..................................................................... 72
Table 3. 3. Effect of amine type on corrosion parameters at 25 °C. ................................................. 75
Table 3. 4. Effect of RTIL type on corrosion rate of carbon steel 1020 at 25 °C. ............................ 76
Table 3. 5. Effect of process temperature on corrosion rate of carbon steel 1020. .......................... 81
Table 3. 6. Corrosion rate of steel in aqueous MEA and MEA+[bmim][BF4] blends at different CO2
loadings and 25 °C. ........................................................................................................................... 82
Table 3. 7. Effect of oxygen presence/absence on corrosion rate of carbon steel 1020 at 25 °C. .... 84
Table 3. 8. Influence of water content in the gas capture fluid on corrosion of steel at 25 °C. ........ 85
Table 5. 1. Viscosities of the capture fluid components at three temperatures. ............................. 122
Chapter 1
1
Introduction and Objectives
1.1. Background
In the energy era driven by greenhouse gas (predominantly CO2) constraints, there are
mounting concerns over the alarming situation of global warming phenomenon, being
intensified brusquely by anthropogenic activities. Fossil fuel based power plants are the
largest among the stationary sources, accounting for approximately 78.6% whereas
refineries and oil & gas processing facilities share about 6.34% of total carbon dioxide
emissions. The Intergovernmental Panel on Climate Change (IPCC) perceives that by the
year 2100 there may be a rise of 1.9 °C in the global temperature [1,2]. This has turned
carbon dioxide capture and sequestration into an extensively investigated topic nowadays.
By and large, there are three major approaches for CO2 capture: chemical/physical
absorption/adsorption; membrane separation, and cryogenic distillation [3]. Cryogenic
distillation, being exceedingly expensive, is not considered feasible regarding the flue gas
purification. Though absorption processes involving chemical solvents (often aqueous
alkanolamines) are being used widely, these put forth a number of limitations that include
insufficient capture capacity, evaporation/degradation of costly reagents and thermal
stability problems, equipment corrosion and high energy consumption during regeneration.
Likewise, physical solvents such as methanol, poly(ethylene glycol) dimethyl ether, (as
well as membrane technology), require higher concentrations of the acid gas in the feed
stream at elevated pressures and lower temperatures. For natural gas sweetening and post-
combustion capture, physical solvents and membranes may not be the efficient tools due to
small concentrations of CO2 and ambient pressures [1,4]. All these discrepancies have to be
overcome and replaced by more efficient and less costly systems.
Ionic liquids (ILs) are being proposed as an alternative for CO2 capture with special
emphasis on their stability, tunabe chemistry, and negligible volatility with considerable
CO2 solubility [5]. Just like common physical solvents, these necessitate feed gas at high
pressure. To surpass the efficiency of industrially well-established alkanolamines systems,
researchers are investigating the abilities of functionalized ionic liquids in bulk form or
Chapter 1
2
through supportive membranes. Nevertheless, to take advantage of CO2 capture capabilities
of both ionic liquids and alkanolamines, combinations of these two might be a better
option. However, to cope with the problem of high viscosities of these ionic liquid fluids,
supported ionic liquid membranes or polymerized ionic liquids are also being probed as an
alternative mechanism.
1.2. Carbon dioxide capture through solvent scrubbing
1.2.1. Chemical solvents
Large point sources of CO2 include fossil fuel-based power/hydrogen production plants,
synthetic fuel industries and natural gas production facilities. Use of a gas capture process
depends on the concentration/partial pressure of CO2 in the feed gas. Natural gas
processing and post combustion capture, where CO2 concentrations are in the range of 2 -
65 vol% and 3 - 15 vol% respectively, mainly involves well established amine based
systems [1,6]. While pre-combustion capture employ physical solvent scrubbers where CO2
is present in proportions greater than 15% under appreciably high pressure. Amines that
gained much consideration in CO2 capture include monoethanolamine (MEA),
diethanolamine (DEA), and methyldiethanolamine (MDEA). In addition to the above stated
alkanolamines, there are certain propriety formulations composed of aqueous solutions of
blended amines along with certain additives like corrosion inhibitors, buffers, foam
depressants, etc [7,8].
Primary/secondary alkanolamines capture CO2 through carbamate formation at lower
temperatures (~ 40 °C) and stripped at higher temperature (≥100 °C). The most accepted
reaction mechanism (1.1) was proposed by Caplow [9].
2 2 2 2
2 2 2
CO RNH RNH CO
RNH CO B RNHCO BH
(1.1)
Since amides are very weak bases, their protonation in aqueous media is not considered
favourable. So the zwitterions concentration should be insignificant.
A more rational mechanism (given below) considering direct interaction of amine with CO2
(dissolved) is the formation of carbamic acid followed by deprotonation [10].
Chapter 1
3
2 2 2
2 2
CO RNH RNHCO H
RNHCO H RNHCO H
(1.2)
For aqueous amine solutions used in gas capture systems, the above mechanisms alone are
not sufficient. Interactions of amine with carbonic acid, bicarbonate/carbonate species
(Figure 1.1) have to be considered while evaluating the process efficiency.
Figure 1. 1. Reaction scheme including reactions between amine, CO2/carbonate and
protons (reproduced from [10]).
One of the negative aspects of primary/secondary alkanolamines is their low CO2 loading
capacity (~50 mol%). Tertiary amines like MDEA possess double the loading capacity
(~100 mol%). However, as tertiary alkanolamines have no labile hydrogen, carbamate
formation is not possible and the feasible hydrolytic mechanism (1.3) is less favourable
kinetically. Most propriety industrial chemical solvents include both primary/secondary and
tertiary alkanolamines (blended amines) in order to enhance the capture capacity [11, 12].
2 2 3
1 2 3 1 2 3
( )CO aq H O HCO H
R R R N H R R R NH
(1.3)
RNHCO2H RNHCO2ˉ
H2CO3 HCO3ˉ
CO2 (aq)
+H
2N
R
-H
2O
+H
2O
-
H2N
R
+H
2N
R
-H
2O
+H
2O
-
H2N
R
H+
H+
+H2NR
-H2NR
+H2O -H2O
+OHˉ
-OHˉ
CO32ˉ
H+
pH
RNH3+ RNH2
H+
Chapter 1
4
In spite of the industrial importance of aqueous amine systems in acid gas capture
processes; these pose a number of drawbacks including:
Equilibrium limitations
Amine evaporation/degradation
Corrosion of equipment
High regeneration costs
CO2 capture facility increases the energy consumption up to 50%, mostly consumed in
solvent regeneration, thus greatly reducing the power plant efficiency [1].
1.2.2. Degradation of amines
During recycling of amine-based CO2 absorption systems, one of the major causes of amine
losses is degradation phenomenon that not only causes reduction in gas capture capacity but
also boosts the corrosion of the equipment and adds to the toxicity of the environment.
Most alkanolamines, especially MEA, degrades quite fast in the presence of oxygen (Figure
1.2) [13]. Degradation phenomenon results in loss of almost 2.2 kg of MEA per tonne of
CO2 captured. Thus disposal and make-up of the degraded solvent considerably heave the
costs [14].
CH2CH2OH:N
H
H
Fe3+
CH2CH2OH+.N
H
H
C CH2OH:N
H
H H
Fe3+
CH CH2OH:NHCH CH2OH:N
H
HOOMEA
CH CH2OH:N
H
HOOH
H2O H2O
CH
O
H2
+ NH3
CH
O
CH2OH
NH3+Imine CH CH2OH:NH
Imine
Peroxide
Peroxide Radical
O2
-H+
Hydroxyacetaldehyde Formaldehyde
Imine RadicalAminium RadicalMEA
H2O
-H2O2
Figure 1. 2. Oxidative degradation mechanism of aqueous MEA (adapted from [13]).
Chapter 1
5
No single mechanism can be used to generalize the degradation phenomenon of
alkanolamines under different operating conditions. This is quite obvious from the analysis
of degradation products at different temperatures. At 100 °C, 1,2-ethanediol, 1,3,5-triazine,
N-butylformamide, 1,4,7,10,13,16-hexaoxacyclooctadecane and 1,2,3,6-tetrahydro-1-
nitropyridine are the degradation products while at 120 °C, methylpyrazine, 7-
oxabicyclo[2.2.1]hept-5-en-2-one, 1-propanamine, ethylamine, 1,3,5-triazine, and 3,3-(1,2-
ethanediyl)bis(syndone) are obtained in case of MEA-H2O-O2-CO2 system [13].
Degradation of amines results in the production of heat stable salts that are impossible to
regenerate under the prevailing conditions of solvent regeneration in gas capture unit.
Increase in temperature, O2 partial pressure, MEA initial concentration, and CO2 loading
significantly raise the rate of oxidative degradation (Table 1.1) [15].
Table 1. 1. Initial rates of oxidative degradation of MEA under different operating
conditions.*
Initial rate
[kmol/(hm3)]
Temperature
(°C)
Initial MEA concentration
(kmol/m3)
O2 concentration
(mol/m3)
0.044 160 2 3.994
0.065 160 3 3.994
0.082 160 4 3.994
0.208 160 11 3.994
0.104 170 3 4.293
0.117 170 4 4.293
0.380 170 8 4.293
0.431 170 10 4.293
0.007 120 4 3.154
0.056 140 4 3.500
0.070 170 3 3.305 *adapted from [15]
CO2 is also detrimental to alkanolamines as higher gas loading is found to increase the
degradation process. Studies have shown resemblance among MEA, DEA, and MDEA
degradation in the presence of CO2; amines, oxazolidinones and imidazolidinones being the
main products. Thermal degradation phenomenon is negligible compared to CO2/O2
induced decay (Figure 1.3) [16].
Chapter 1
6
Figure 1. 3. Amine degradation: ■ thermal; ■ CO2 induced (adapted from [16]).
Certain other impurities like SO2, in the flue gas also have adverse effect on amine
degradation (Figure 1.4). Corrosion inhibitors used in amine systems also found to trigger
solvent degradation, see Figure 1.5 [17].
0
10
20
30
40
50
60
70
80
90
100
0.34
951
1.36
89
2.33
01
3.34
95
4.33
98
5.35
92
6.34
95
7.33
98
8.33
01
9.34
95
10.34
11.33
■Thermal degradation: 4 mol.kg-1
amine, 140 °C, 15 days
■ CO2 induced degradation: 4 mol.kg-1
amine, 140 °C, 15 days, 2MPa CO2
(%)
MD
EA
DM
AE
AM
P
ME
A
MA
E
DE
A
HE
ED
A
DM
P
TM
ED
A
N,N
-diM
ED
A
N,N
,N’-
triM
ED
A
N,N’-
diM
ED
A
Chapter 1
7
Figure 1. 4. Effect of SO2 on MEA degradation [17].
Figure 1. 5. Effect of corrosion inhibitor (NaVO3) on MEA degradation [17].
0
0.001
0.002
0.003
0.004
0.005
0.006
0.007
0 50 100 150
Time (hr)
Rate
of
deg
rad
ati
on
(m
ol/
L.h
r)
6% O2
6% O2 + 6 ppm SO26% O2 + 11 ppm SO2
0
1
2
3
4
5
6
0 50 100 150 200 250
Time (hr)
ME
A c
on
cen
trati
on
(m
ol/
L)
with NaVO3
without NaVO3
Chapter 1
8
Moreover, regeneration of alkanolamine to enable smooth recycling is not cost effective.
Among alkanolamines, MEA is the most efficient CO2 capture agent with high CO2-
captured product (carbamate) stability which correspondingly accounts for high energy
requirements during regeneration, a major drawback leading to hiking costs of the process
[18].
1.2.3. Corrosion of equipment
All amine treating plants face corrosion problems. Bottom of absorbers and regenerators,
pumps and valves are more vulnerable to corrosion due to high gas loading and elevated
temperatures. Corrosion is also a major concern in the safety of plants, causing weakening
of the equipment that may lead to explosion of pressure vessels. High CO2 loading,
increased concentration of amine/O2, elevated temperatures and higher solution velocities
all cause corrosion rate to accelerate (Figures 1.6). Corrosion is the result of anodic (iron
dissolution) and cathodic (reduction of oxidizers present in the solution) electrochemical
reactions [19]. Dissociation of water/protonated amine, hydrolysis of CO2 and amine
regeneration reactions provide oxidizing species enabling corrosion process to continue
[20]. Most significant anodic/cathodic reactions leading to corrosion are:
2
2
3 3 2
2 2
2
2 2 2
2 2 2
Fe Fe e
HCO e CO H
H O e OH H
(1.4)
Chapter 1
9
Figure 1. 6. Effect of various parameters on the corrosion rate of carbon steel C1020 in
aqueous MEA (basal conditions: MEA conc. 5 kmol/m3; gas loading 0.4 mol CO2/mol
MEA; 80 °C temperature) [19].
Other impurities like SO2 also gear up wear and tear of the equipment [21]. Presence of
SO2 speeds up the corrosion phenomenon through the formation of hydrogen ions as shown
below:
2 2 3
2
3 3
2122 2 2 42
SO H O H HSO
HSO H SO
SO O H O H SO
(1.5)
Or SO2 may react with O2 and iron causing direct corrosion of steel:
2 2 4
4 2 2 2 3 2 2 4
2 4 2 4 2
4 6 2 4
4 4 2 4 4
Fe SO O FeSO
FeSO O H O Fe O H O H SO
H SO Fe O FeSO H O
(1.6)
1.2.4. Corrosion inhibition
By employing suitable inhibitors, corrosion phenomenon may be effectively suppressed up
to 80%. A number of corrosion inhibitors, based on arsenic, antimony, vanadium, copper
Chapter 1
10
(like NaVO3, CuCO3) are being used in order to control and prevent corrosion that not only
adds to the capital cost but most of these are toxic and hazardous to life as well.
Degradation also plays its role in this regard as this phenomenon not only depletes the
active CO2 capturing species but the resulting products also enhance the corrosion rate by
lowering the inhibition ability. Presence of certain salts like NaCl greatly lowers the
inhibitor efficiency which might be due to the attack of Clˉ ions on the passive film [22].
Presence of heat stable salts (acetate, formate, oxalate, etc) increase the corrosion rate,
probably by introducing additional oxidizing agents [23].
A number of organic (including thiourea, salicylic acid) and inorganic (vanadium,
antimony, copper, cobalt, tin and sulfur compounds) inhibitors have been exploited.
Sodium metavanadate (NaVO3) is the most trusted in amine based CO2 capture plants that
can reduce corrosion rate to less than 1 mpy (0.0254 mm/year). In spite of their successful
use, the probable consequences of inhibitors‟ toxicity (more toxic than absorption solvents,
Table 1.2) on human health and environment are of great worry. The more strict regulations
in case of toxic/hazardous substances in very near future may limit the use of such
compounds due to high disposal costs [24].
Table 1. 2. Toxicity of absorption solvents and conventional inhibitors.*
Chemical LD50-orala (mg/kg)
Mouse Rat
Absorption solvents
Monoethanolamine (MEA) 700 1720
Diethanolamine (DEA) 3300 710
Conventional inhibitors
Vanadium pentaoxide 23 10
Sodium metavanadate 74.6 98
Ammonium metavanadate 25 58.1 *adapted from [24];
a LD50 (lethal dose) is the dose large enough to kill 50% of a sample of animals under
test.
1.3. Physical Solvents
To cope with the problems of higher regeneration energy requirements, degradation and
corrosion posed by chemical solvents (aqueous alkanolamines), physical solvents (Table
1.3) have been employed where there is higher CO2 concentration found in the feed gas
Chapter 1
11
(such as pre-combustion capture). But these face their own downsides i.e. prerequisite of
higher CO2 concentrations, elevated pressures, and refrigeration/cooling of the solvent/feed
gas. Moreover, most physical solvents are also liable for dissolution of heavier
hydrocarbons in reasonable quantities [11].
Table 1. 3. Physical solvent processes.*
Process name Solvent Licensor
Fluor solvent Propylene carbonate (PC) Fluor Corporation
SELEXOL Dimethyl ether of polyethylene glycol
(DMPEG)
Dow Chemical Company
Purisol N-Methyl-2-pyrrolidone (NMP) Lurgi
Rectisol Methanol Lurgi
Sulfinol Sulfolane and MDEA/DIPA
(Mixed physical/chemical solvent)
Jacobs
*reproduced from [11]
1.4. Ionic Liquid Solvents
An exciting new class of solvents known as ionic liquids (ILs), entirely composed of ions,
are being synthesized and investigated for diverse applications such as organic/inorganic
reactions, catalysis, metal extraction, gas separations, etc. The prime advantage of using
ionic liquids is that these have no detectable vapor pressure and hence don‟t contribute to
atmospheric pollution. Also, owing to the availability of numerous constituent ion pairs,
thousands of binary ionic liquids are potentially possible and by choice application specific
solvent can be synthesized [25,26].
In the following pages, from the perspective of carbon dioxide capture, the literature has
been reviewed to have a thorough knowhow about the use of these novel species.
1.5. Ionic liquids for CO2 capture - Development and progress*
Abstract/Résumé
Innovative off-the-shelf CO2 capture approaches are burgeoning in the literature, among
which, ionic liquids seem to have been omitted in the recent Intergovernmental Panel on
Climate Change (IPCC) survey. Ionic liquids (ILs), because of their tunable properties,
* M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Chem. Eng. Process. 49 (2010) 313-322.
Chapter 1
12
wide liquid range, reasonable thermal stability, and negligible vapor pressure, are emerging
as promising candidates rivaling with conventional amine scrubbing. Due to substantial
solubility, room-temperature ionic liquids (RTILs) are quite useful for CO2 separation from
flue gases. Their absorption capacity can be greatly enhanced by functionalization with an
amine moiety but with concurrent increase in viscosity making process handling difficult.
However this downside can be overcome by making use of supported ionic-liquid
membranes (SILMs), especially where high pressures and temperatures are involved.
Moreover, due to negligible loss of ionic liquids during recycling, these technologies will
also decrease the CO2 capture cost to a reasonable extent when employed on industrial
scale. There is also need to look deeply into the noxious behavior of these unique species.
Nevertheless, the flexibility in synthetic structure of ionic liquids may make them
opportunistic in CO2 capture scenarios.
Des approches de capture du CO2 innovantes sont en plein essor comme le révèle la
littérature actuelle, parmi lesquelles, les liquides ioniques semblent avoir été omis dans la
récente revue du GIEC (Intergovernmental Panel on Climate Change, IPCC). Les liquides
ioniques (ILs), en raison de leurs propriétés ajustables, large gamme de liquide, stabilité
thermique et pression de vapeur négligeable, apparaissent comme des candidats
prometteurs rivalisant avec les amines dans les contacteurs gaz-liquide classiques. En
raison de la solubilité importante de gaz, les liquides ioniques à température ambiante
(RTILs) sont très utiles pour la séparation du CO2 des gaz de combustion. Leur capacité
d'absorption peut être grandement améliorée par fonctionnalisation avec un groupement
amine, mais avec une augmentation concomitante de la viscosité rendant le contrôle du
procédé difficile. Toutefois cet inconvénient peut être surmonté par la mise en oeuvre de
membranes à base de liquides ioniques, en particulier lorsque les pressions et températures
élevées sont impliquées. En outre, en raison de la perte négligeable de liquides ioniques
lors du recyclage, ces technologies permettront aussi de réduire le coût du captage du CO2
dans une mesure raisonnable lorsqu'elles sont utilisées à l'échelle industrielle. Il est
également nécessaire d‟examiner attentivement le caractère toxique de ces espèces.
Néanmoins, la souplesse de la structure de synthèse des liquides ioniques peut les rendre
abordables dans les scénarios de capture du CO2.
Chapter 1
13
1.5.1. Introduction
Recent concept of using ionic liquids (Figure 1.7) for CO2 capture is gaining interest due to
their unique characteristics, i.e., wide liquid range, thermal stability, negligible vapor
pressure, tunable physicochemical character and high CO2 solubility. An important
drawback much discussed in the case of ILs is their high viscosity. However, by choosing
an appropriate combination of cation and anion, the viscosities can be adjusted over an
acceptable range of <50 cP to >10,000 cP. For CO2 capture at high temperatures and high
pressures, such as in integrated gasification combined cycle (IGCC) pre-combustion
capture, IL viscosity is less of a concern for its sharp decrease at elevated temperatures,
though thermodynamics of CO2 absorption untowardly dictates poor abatement
performances. Therefore, paths pursued in recent research works include the use of ionic
liquids for carbon dioxide capture involving room-temperature ionic liquids (RTILs), task-
specific ionic liquids (TSILs) or supported ionic-liquid membranes (SILMs) [27-30].
Figure 1. 7. Some cations and anions constituting ionic liquids (ILs).
1.5.2. CO2 capture by room-temperature ionic liquids (RTILs)
Considerable research work is being done showing high carbon dioxide solubility in certain
RTILs, especially in those having imidazolium-based cations. Depending on their thermal
stability and CO2 selectivity (in general over nitrogen and smaller hydrocarbons), ILs are
stronger candidates (for CO2 capture) compared to certain conventional solvents such as
methanol, ethanol, and acetone [31]. RTILs portray a typical behavior of a physical solvent;
y
Chapter 1
14
that is, increase in CO2 partial pressure results in linear increase in the gas solubility
whereas temperature exerts opposing effect on CO2 absorption in the RTIL [32]. The
solubility of carbon dioxide, ethylene, ethane, methane, argon, oxygen, carbon monoxide,
hydrogen and nitrogen in 1-n-butyl-3-methylimidazolium hexafluorophosphate,
[bmim][PF6] in the temperature range between 10 and 50 °C and pressures up to 13 bar
proves the superiority of IL over various organic solvents like heptane, cyclohexane,
benzene, ethanol and acetone (see Figure 1.8 and Table 1.4). Dissolution enthalpy and
entropy values suggest stronger interaction of CO2 with the IL, [bmim][PF6]. The relatively
higher solubility of CO2 may be attributed to its quadrupole moment and dispersion forces.
Owing to their negligible volatility and thermal stability under the explicit conditions, ILs
are unlikely to contaminate the gas stream. Raeissi and Peters verified the thermal stability
of 1-n-butyl-3-methylimidazolium bis[trifluoromethylsulfonyl]imide, [bmim][Tf2N], by
conducting the gas capture experiments in the temperature range of 40–177 °C and
pressures up to 140 bar. Even after keeping at 177 °C for more than 10 h, the ionic-liquid
conferred reproducible results for CO2 solubility [33,34]. Mass transfer of the gas is of
much importance especially where gas is to undergo a chemical interaction. Hence, during
fabrication of an appropriate ionic liquid, drawbacks posed by high viscosity must be
addressed [34].
Figure 1. 8. Solubilities of CO2, C2H4, C2H6, CH4, Ar and O2 in [bmim][PF6] at 25 °C
(adapted with permission from [34]).
0
0.05
0.1
0.15
0.2
0.25
0 2 4 6 8 10 12 14
Pressure (bar)
Mo
le F
rac
tio
n
CO2 C2H4
C2H6 CH4
Ar O2
Chapter 1
15
Table 1. 4. Henry‟s constants (bar, at 25 °C) for gases in different organic solvents.*
[bmim][PF6] heptane cyclohexane benzene ethanol acetone
CO2 53.4 84.3 133.3 104.1 159.2 54.7
C2H4 173 44.2a - 82.2 166.0 92.9
C2H6 355 31.7 43.0 68.1 148.2 105.2
CH4 1690 293.4 309.4 487.8 791.6 552.2
O2 8000 467.8 811.9 1241.0 1734.7 1208.7
Ar 8000 407.4 684.6 1149.5 1626.1 1117.5
CO nondetect 587.7 1022.5 1516.8 2092.2 1312.7
N2 nondetect 748.3 1331.5 2271.4 2820.1 1878.1
H2 nondetect 1477.3 2446.3 3927.3 4902.0 3382.0 *adapted with permission from [34];
a for ethylene in hexane
The experimental and simulation studies have shown that CO2 is significantly soluble in
alkylimidazolium-based ILs. The origin of this high solubility could be related more to the
anion moiety that enhances interactions by favoring peculiar distributions of CO2 molecules
around the cation [35]. Alkyl-side chain length of the imidazolium cation of the ILs also
affects CO2 solubility to a certain extent (Table 1.5). Fluorine substituted side chains
greatly augment the uptake of CO2 compared to the corresponding non-substituted side
chains but at the expense of an increase in viscosity [36-39].
Table 1. 5. Henry‟s constants for CO2 in different ionic liquids.*
Ionic Liquid HCO2 (bar)
C3mimTf2N 37 ± 7
C3mimTf2N with constant-density gas 39 ± 1
C3mimPF6 52 ± 5
C4mimTf2N 37 ± 3
C4mimTf2N with 2.7 wt% polyethylenimine 38 ± 3
C6mimTf2N 35 ± 5
C8mimTf2N 30 ± 1
C8mimTf2N with 20% relative humidity 30 ± 2
C8mimTf2N with 40% relative humidity 27 ± 4
C8F13mimTf2N 4.5 ± 1
C8mimTf2N (58 mol%)/C8F13mimTf2N (42 mol%) 15 ± 1
1,4-dibutyl-3-phenylimidazolium bis(trifluoromethylsulfonyl)imide 63 ± 7
1-butyl-3-phenylimidazolium bis(trifluoromethylsulfonyl)imide 180 ± 17 *adapted with permission from [39]
The nature of anion seems to have a stronger influence on gas solubility than that of the
cation. Ionic liquids possessing [Tf2N] anion show higher CO2 solubility among
imidazolium-based RTILs (Table 1.6). A number of factors like free volume, size of the
counter ions, and strength of cation-anion interactions within the ionic liquid structure seem
Chapter 1
16
to govern CO2 solubility in RTILs. Higher gas solubility with increase in alkyl-side chain
may be the result of increased free volume available for CO2 with corresponding decrease
in cation–anion interactions [40,41]. The thermal stability and negligible volatility of
RTILs make them quite viable for prolonged use. Hou and Baltus found that even after
regenerating the ionic liquid six times, by N2 purging followed by vacuum application at 70
°C, there was practically no change in gas capture capacities [40].
Table 1. 6. Henry‟s law constants of CO2 in ionic liquids.*
Ionic Liquid HCO2 (bar)
10 ˚C 20 ˚C 25 ˚C 30 ˚C 40 ˚C 50 ˚C
[bmim][Tf2N] 28 ± 2 30.7 ± 0.3 34.3 ± 0.8 42 ± 2 45 ± 3 51 ± 2
[pmmim][Tf2N] 29.6 ± 0.6 34 ± 3 38.5 ± 0.9 40.4 ± 0.6 46 ± 3 53 ± 2
[bmpy][Tf2N] 26 ± 1 31.2 ± 0.1 33 ± 1 35 ± 2 41 ± 4 46 ± 1
[perfluoro-hmim][Tf2N] 25.5 ± 0.2 29.2 ± 0.4 31 ± 2 32 ± 2 36 ± 4 42 ± 2
[bmim][BF4] 41.9 ± 0.2 52 ± 2 56 ± 2 63 ± 2 73 ± 1 84 ± 4 *adapted from [40]
The equilibrium pressure not only depends on temperature but also on CO2 concentration.
At 60 bar, CO2 solubility in 1-ethyl-3-methylimidazolium
bis[trifluoromethylsulfonyl]imide, [emim][Tf2N], is found to be 60 mol%. When compared
with 1-ethyl-3-methylimidazolium hexafluorophosphate, [emim][PF6], the gas is found
more soluble in IL with [Tf2N]− anion. The difference is further pronounced at higher CO2
mole fraction (Figure 1.9). Such data is confirming the effect of anion on CO2 interaction
with IL [42]. Fluoroalkyl group enhances CO2 solubility, thus making [emim][Tf2N] more
efficient for CO2 capture.
Chapter 1
17
Figure 1. 9. CO2 solubility in [emim][Tf2N] and [emim][PF6] (adapted from [42]).
Room-temperature ionic liquids can be effectively used for hydrogen purification with high
selectivity for CO2/H2 separation. Selectivity of the ionic liquid, [bmim][PF6], for CO2/H2
mixtures constituting 45–50 wt% H2 is in the range of 30-300. Selectivity drops at higher
temperature but enhances with pressure increase [43-45]. Hence this setup may be
employed in CO2 capture from pre-combustion power plants. A pressure-swing
adsorption/desorption method can be employed for H2 purification by RTILs. CO2 showed
good solubility in 1-ethyl-3-methylimidazolium 2-(2-methoxyethoxy)ethylsulfate,
[emim][MDEGSO4] at 30 °C in the pressure range of 8.54–67 bar, and expectably
increasing with pressure rise (Table 1.7). Pyrrolidinium and ammonium based RTILs like
1-n-butyl-1-methylpyrrolidinium bis(trifluoromethylsulfonyl)amide ([bmpy][Tf2N]) and
trimethyl(butyl)ammonium bis(trifluoromethyl)sulfonyl)imide ([N(4)111][Tf2N]) have also
been investigated for H2 purification showing CO2 absorption capacity comparable to
imidazolium-based RTILs in the temperature range of 20-140 °C [46,47]. Regarding
H2S/CO2 selectivity, H2S was found almost three times more soluble than CO2 in 1-(2-
hydroxyethyl)-3-methylimidazolium tetrafluoroborate ([hemim][BF4]). However, owing to
the greater concentration of CO2 in the flue gases, higher partial pressure of CO2 diminishes
0
100
200
300
400
500
600
700
800
900
0 0.1 0.2 0.3 0.4 0.5 0.6
P (
ba
r)
x (CO2)
[emim][PF6]
[emim][Tf2N]
Chapter 1
18
this advantage. This observance illustrates that RTILs can be efficiently tailored to remove
H2S and CO2 concurrently [48,49].
Table 1. 7. CO2 solubility data in [emim][MDEGSO4].*
P/bar mCO2 a/(molCO2 . kgIL
-1) P/bar mCO2
a/(molCO2 . kgIL
-1)
30 ˚C 40 ˚C
8.540 0.3850 8.650 0.3301
14.72 0.6654 14.97 0.5713
28.67 1.3239 28.88 1.1162
42.30 2.0404 42.81 1.7053
55.21 2.7357 56.62 2.2899
62.30 3.0936 63.50 2.5606
50 ˚C 60 ˚C
8.420 0.2743 8.470 0.2380
15.12 0.4911 15.21 0.4257
29.38 0.9587 29.61 0.8235
43.59 1.4509 43.95 1.2359
57.32 1.9205 57.70 1.6254
65.20 2.1710 66.36 1.8551
70 ˚C
8.560 0.2110
15.22 0.3737
29.87 0.7171
44.27 1.0655
58.68 1.4097
67.10 1.6008 *adapted from [44];
a with buoyancy correction
The viscosity of common RTILs is quite high, [bmim][BF4] (79.5 cP) is found to be 40
times more viscous compared to 30% MEA (monoethanolamine) solution at the same
temperature (33 °C) [50]. To cope with the viscosity constraints, RTILs may be mixed with
some common organic solvents or water. Addition of water (IL aqueous solutions) helped
overcome viscosity problems as shown in Table 1.8 [51]. However, inclusion of such
liquids will come with their drawbacks as well. Besides, the advantage comes at the
expense of a decrease in gas capture capability. This is evident from the behavior of an
ionic liquid [Choline][Pro] (Figure 1.10) examined in pure form as well as after mixing
with polyethylene glycol (PEG 200) at temperatures 35–80 °C and ambient pressure [52].
Gas solubility decreased with increasing amount of PEG 200, under constant temperature
and pressure conditions. This is explicable because of the low CO2 solubility in PEG 200.
However, to enhance the rate of both absorption and desorption, addition of an appropriate
Chapter 1
19
amount of PEG 200 has been found favorable. This may be due to the decrease in viscosity
and/or solvent role of PEG 200.
Figure 1. 10. Proposed mechanism for chemical absorption of CO2 by the TSIL (adapted
from [52]).
Table 1. 8. Viscosity values for different compositions of tri-iso-
butyl(methyl)phosphonium tosylate/water mixtures.*
Mass fraction IL ± 0.0001/(w/w) η±σa (cP)
0.0000 0.89
0.1250 1.65±0.08
0.2500 2.6±0.1
0.3750 4.0±0.2
0.5000 6.9±0.3
0.6250 11.6±0.5
0.7500 23.0±0.7
0.8720 68.0±2.0
1.0000 1320±13 *adapted from [51];
a standard deviations
Another more workable option, in case of alkanolamine systems, may be the replacement
of aqueous medium with some stable and non-volatile room-temperature ionic liquid in
order to combine the advantages of both, i.e., negligible vapor pressure, higher thermal
stability and lower heat capacity of ionic liquids, and fast capture kinetics and low viscosity
of certain alkanolamines [53]. Switching the CO2 capture product (carbamate in this case)
into a foreign phase would pull the equilibrium-limited CO2 absorption towards higher CO2
conversion values, unlike in conventional aqueous amine solutions with soluble carbamate
salt (Figure 1.11). Thus, it can be inferred that to take advantage of useful properties of ILs,
Chapter 1
20
amine-IL solutions need to be investigated more deeply as potential replacement solvents
for aqueous amine scrubbing systems.
Figure 1. 11. [hmim][Tf2N]-MEA solution: (a) fresh sample; (b) on CO2 exposure;
showing precipitated MEA-carbamate (reprinted with permission from [53]).
Regarding natural gas purification, certain hygroscopic imidazolium-based ionic liquids
like [bmim][PF6], [C8mim][BF4] and [C8mim][PF6] have the ability to dehydrate the gas
stream as well [54-56]. Also, the presence of water along with acetate ion in some ionic
liquids akin to [hmim][acetate] and [bmim][acetate] may facilitate the capture phenomenon
through weak bonding with CO2 [57]. Diminished corrosion of the equipment, almost one-
third the heat capacity of (especially imidazolium-based) RTILs, compared to the aqueous
systems, may help rationalize the large scale application of these unique species for CO2
capture [20,53,58–60].
In short, room-temperature ionic liquids especially imidazolium-based RTILs may be
employed in natural gas/hydrogen purification or in CO2 capture from fossil fuel based
power plants. Regarding regeneration, room-temperature ionic liquid based materials may
be easily recovered either by pressure sweep process coupled with vacuum treatment, by
applying heat or by bubbling nitrogen through the absorbent [50,52]. However, task-
Chapter 1
21
specific ionic liquids or RTILs mixed with amine bearing species require temperature
sweep regeneration involving vacuum heating [39].
1.5.3. CO2 capture by task-specific ionic liquids (TSILs)
As discussed above, CO2 is sufficiently soluble in room-temperature ionic liquids (RTILs).
However, the CO2 capture ability can be significantly enhanced by introducing basic
character in the ILs. Functionalization of ionic liquids with a suitable moiety (like amine)
may be opted in this regard [61,62]. CO2 absorption ability of TSILs can reach up to
threefold that of the corresponding RTILs. The enhanced effect of pressure in case of
TSILs was observed by the fact that there was a steady increase in gas loading with rise in
pressure, providing evidence both for chemical as well as physical sorption. The effect is
not so apparent in case of aqueous amine solutions which possess stoichiometric limitations
[50]. Reversible sequestration of CO2 has been achieved by attaching primary amine
moiety to an imidazolium cation, without any decrease in the ionic-liquid stability. For five
consecutive cycles of gas absorption/desorption, the regenerated TSIL ([pabim][BF4]) did
not show any loss of efficiency. [pabim][BF4] exhibits better CO2 capture competence
compared to [hmim][PF6], owing to chemical capture phenomenon in the former. The TSIL
when exposed to CO2 for 3 h at room temperature and pressure, the mass gain was 7.4%
which corresponds to 0.5 molar uptake of CO2 (maximum theoretical value for CO2 capture
as amine carbamate). The proposed mechanism of interaction between CO2 and
[pabim][BF4] is shown in Figure 1.12. The inclusion of water in the ionic liquid was found
to increase the CO2 holding capacity which might be due to the formation of additional
bicarbonate species [63,64].
Figure 1. 12. Proposed mechanism for CO2 capture by [pabim][BF4] (adapted from [63]).
Chapter 1
22
In spite of the tunable approach towards TSILs, these functionalized species exhibit much
higher viscosities as compared to the corresponding RTILs or other commercially available
CO2 scrubbing solutions, posing too serious complications to be applicable on an industrial
scale. CO2 capture by TSILs causes a sharp increase in viscosity, resulting into a gel-like
material [65]. This drawback may be avoided by utilizing mixtures of TSILs and RTILs or
TSILs may be adsorbed onto porous membranes.
Comparison of CO2 capture by ionic liquids with that by conventional aqueous amine
solutions (30 wt% MEA/MDEA) illustrates that the absorption activities of ionic liquids
resembles that of common physical solvents (Figure 1.13). Nonetheless, CO2 absorption
ability increases significantly on functionalization of ionic liquid with primary amine
moiety. Task-specific ionic liquids, [Amim][BF4] and [Am-im][DCA], perform like
chemical solvents at low pressures (≤1 bar). However, at higher pressures, they pursue the
performance of room-temperature ionic liquid, [bmim][BF4]. On the other hand, aqueous
amine solutions accomplish the maximum capacity at about 2 bar and any further increase
in pressure does not seem feasible. Whereas functionalized ionic liquids (TSILs) carry on
steady CO2 absorption with ascending pressure even beyond the stoichiometric limit
[50,66]. This behavior shows that TSILs possess both chemical as well as physical tools for
gas capture.
Chapter 1
23
Figure 1. 13. Molar CO2 loads in solvent volume (for MEA/MDEA, consider aqueous
solution volume): data for ionic liquids at 30 °C [50]; data for MEA and MDEA at 40 °C
[66].
1.5.4. CO2 capture by supported ionic-liquid membranes (SILMs)
A number of studies have been performed to explore the prospects of supported ionic-liquid
membranes involving RTILs or TSILs or both in CO2 capture applications. To take
advantage of thermal/chemical stability and essentially no volatility, and to deal with the
limitations due to viscosity and also to increase the contact area between gas and ionic
liquid, supported ionic liquids may prove a better choice in CO2 separation from flue gases.
RTIL, [bmim][Tf2N], supported on porous alumina membrane revealed very encouraging
results in favor of CO2 separation ability [67]. The SILM with [bmim][Tf2N] shows higher
CO2/N2 selectivity of 127 than that with [C8F13mim][Tf2N] (72). Furthermore, the
fluorinated ionic liquid is much more viscous than [bmim][Tf2N] that tends to cause a
decrease in CO2 diffusivity. A proposed process diagram regarding the application of SILM
in a coal-fired power plant is shown in Figure 1.14. SILMs may compete economically
with commercial amine scrubbing provided permeance and selectivity are optimized. Ionic
liquids like [bmim][PF6] adsorbed to a porous (ceramic or zeolite) material may be
Chapter 1
24
employed for CO2 separation by introducing pressurized gas on one side and collecting the
CO2-depleted gas downstream of the porous medium [68].
Figure 1. 14. Proposed setup for CO2 separation by SILM in a coal-fired power plant
(adapted from [67]).
In another study [69,70], [bmim][BF4] was adsorbed onto polyvinylidene fluoride (PVDF)
polymeric membrane. The mass ratio of IL/membrane in SILMs was kept 0.5–2.0. With the
increase of IL content, the permeability coefficient was seen to increase abruptly. Rise in
temperature resulted in a corresponding increase in membrane free volumes caused by
increased mobility of polymeric chains. This development stimulated simultaneous increase
in permeability. However, the selectivity for CO2 decreased when compared with CH4. This
is because CH4 show more diffusion selective property than solubility selective property
and so its solubility is more affected by membrane structure. The rise in pressure
demonstrates a positive effect on selectivity. Through optimization of operating conditions,
25-45 CO2/CH4 selectivity was achieved. The solubility behavior of CO2, H2, CO and CH4
in two ionic liquids, [bmim][Tf2N] and [emim][Tf2N] makes their usage interesting as
separation membranes [71]. The solubility of CO2 in the two ionic liquids reaches up to 60
mol% compared to that of H2 that remains up to 7 mol% at 90 bar. The pressure increase
has little effect on H2, CO and CH4 solubility compared to that of CO2. However, taking
Chapter 1
25
into account the economics of the capture process, optimum conditions of temperature and
pressure need to be set.
Amino-acid based ionic liquids supported on porous silica show fast CO2 capture compared
to the gas absorption into the corresponding pure ionic liquids. Experiments with supported
TSIL reveal 50 mol% CO2 capture capacity through carbamate formation with reference to
ionic liquid amount. However, in presence of small amounts of water (~1 mass%), the
capture capacity reaches equimolar ratio as shown in Figure 1.15 (a)-(c). In the latter case,
the capture results into carbonate formation [72]. Similarly, the imidazolium, pyridinium,
pyrrolidinium, phosphonium, ammonium, and guanidinium based ionic liquids can be
adsorbed to polymeric materials for gas separations, especially for CO2, NOx and SOx [73].
a)
b)
Figure 1. 15. Proposed mechanisms of CO2 capture: (a, b) without water; (c) with water
(reproduced with permission from [72]).
c)
Chapter 1
26
Supported ionic-liquid membranes (especially bearing amine functionalized TSILs) possess
high selectivity and stability, also diminishing negative impact due to high viscosity of
TSILs (Table 1.9). In a porous polytetrafluoroethylene (PTFE) membrane with non-
functionalized ionic liquids like [C4mim][Tf2N], the permeation of gas is by solution-
diffusion mechanism whereas SILMs with adsorbed TSILs like [C3NH2mim][CF3SO3] or
[C3NH2mim][Tf2N] demonstrate much higher CO2 permeation, mediated by chemical
interaction with amine moiety. The studied SILMs possess high stability, confirmed by
continuous use for 260 days without any detectable loss in performance [30,74].
Nevertheless, increase in temperature has a negative effect on permeation of CO2 as high
temperature prevents the interaction between CO2 and amine moiety. Temperature rise
above 85 °C results in corresponding decrease in CO2 solubility as well as carbamate
stability, and diffusion phenomenon starts to dominate [75]. Even so, combining SILMs
with TSILs may possibly be a better choice for CO2 separation at elevated temperatures and
pressures [76]. In case of hydrophilic composite membranes, presence of moisture in flue
gas affects the CO2 separation performance. Moist feed seems to increase permeability up
to 35-fold without any detectable loss in CO2/H2 or CO2/N2 selectivity as compared to dry
feed [74]. The capabilities of amine-functionalized TSILs based on beta-hydroxy amines,
aryl amines and tertiary amines may prove greatly supportive in this regard for proficient
reversible CO2 uptake [77].
Table 1. 9. Viscosities and water content of the ionic liquids, at 25 °C.*
Abbreviation Molecular Structure Water content
(%)
Viscosity
(cP)
[C3NH2mim][CF3SO3]
11.4 3760
[C3NH2mim][Tf2N]
5.7 2180
[C4mim][Tf2N]
1.8 70
*adapted from [30]
Development of more efficient and cost-effective SILMs requires in-depth study to probe
the role of anion/cation in optimization of molar volume of constituent ionic liquids that
should lead to the fabrication of more stable, more selective, more permeable but thin
membranes [78].
Chapter 1
27
1.5.5. CO2 capture by polymerized ionic liquids
One of the negative aspects of SILMs is the leaching of the liquid through membrane pores
as the pressure drop surpasses the liquid stabilizing forces within the matrix. Membranes
made up of polymerizable ionic liquids may be a better option for CO2 separation [5]. CO2
absorption experiments with ionic-liquid polymers demonstrate their superiority over
RTILs [79]. 1-[2-(Methacryloyloxy)ethyl]-3-butyl-imidazolium tetrafluoroborate,
[MABI][BF4]; 1-(p-vinylbenzyl)-3-butyl-imidazolium tetrafluoroborate, [VBBI][BF4]; 1-
(p-vinylbenzyl)-3-butyl-imidazolium hexafluorophosphate, [VBBI][PF6]; 1-(p-
vinylbenzyl)-3-butylimidazolium o-benzoicsulphimide, [VBBI][Sac]; 1-(p-vinylbenzyl)-3-
butyl-imidazolium trifluoromethane sulfonamide, [VBBI][Tf2N]; 1-(p-vinylbenzyl)-3-
methyl-imidazolium tetrafluoroborate, [VBMI][BF4] ionic-liquid polymers were found
remarkably fit for CO2 capture (Table 1.10).
Table 1. 10. Summary of gas absorption capacities (at 592.3 mmHg & 22 °C) and glass
transition temperatures of poly(ionic liquid)s.*
Poly(ionic liquid)s or
Ionic Liquid
Tg (°C) CO2 Absorption Capacity (mol %)
P[VBBI][PF6] 85 2.8
P[VBBI][BF4] 78 2.27
P[VBBI][Sac] 40 1.55
P[VBBI][Tf2N] 3 2.23
P[VBMI][BF4] 110 3.05
P[MABI][BF4] 54 1.78
P[EIBO][BF4] 33 1.06
[bmim][BF4] - 1.34 *adapted from [79]
In contrast to RTILs, the poly(ionic liquids) with PF6− show higher efficiency as compared
to those with BF4− or Tf2N
− anions. Moreover sorption/desorption rates of the polymerized
ionic liquid is quite fast as compared to RTILs. The bulk absorption phenomenon appears
to govern the capture progress [80-82]. RTILs containing polymerizable entities show
higher permeability, solubility and diffusivity values for CO2, as given in Table 1.11 [83].
Chapter 1
28
Table 1. 11. Permeability, solubility and diffusivity values in: a) styrene-based poly(ionic
liquid)s; b) acrylate-based poly(ionic liquid)s, at 20 °C.*
a)
Styrene CO2 N2 CH4
Pa S
b D
c P S D P S D
Methyl 9.2±0.
5
4.0±0.1 1.7±0.1 0.29±0.01 N/A N/A 0.24±0.01 0.21±0.05 0.88±0.16
Butyl 20±1 4.4±0.3 3.5±0.4 0.67±0.02 N/A N/A 0.91±0.06 0.55±0.07 1.28±0.20
Hexyl 32±1 3.9±0.1 7.7±0.4 1.4±0.1 0.1±0.01 11±2 2.3±0.1 0.57±0.03 3.10±0.15
b)
Acrylate CO2 N2 CH4
P S D P S D P S D Methyl 7.0±0.4 3.6±0.1 1.5±0.1 0.23±0.02 N/A N/A 0.19±0.02 0.17±0.04 0.89±0.20
Butyl 22±1 4.5±0.4 3.6±0.4 0.71±0.06 N/A N/A 0.97±0.08 0.59±0.09 1.27±0.09
*adapted with permission from [83]; a Permeability in Barrers;
b Solubility in cubic centimeters gas (STP) per
cubic centimeter polymer atmosphere; c Diffusivity in squared centimeters per second x 10
8
As the length of the alkyl chain increases, gas permeability and diffusivity increases
considerably. However, styrene-based polymer with methyl group shows higher CO2
permeability than the corresponding acrylate-based polymer. The CO2 solubility was found
quite high in both types of poly(RTILs) but lower than that for poly(RTILs)-PEG
copolymers [84]. Polymerizable ionic liquids exhibit high CO2 capture capacity and
selectivity with respect to N2, O2 or CH4 [85]. These polymerized structures can capture
almost double the amount of CO2 compared to the corresponding RTILs. The efficiency of
these polymeric structures can be enhanced further by modifying monomers with
appropriate entities like oligo(ethylene glycol) or nitrile-containing alkyl groups [86]. By
incorporating an appropriate amount of RTIL and consequently introducing free ion pairs
into the poly(RTIL) membranes, CO2 permeability and CO2/N2 selectivity may be
increased up to about 300–600% and 25% respectively [87,88]. Presence of longer alkyl
chains on the cations of poly(ionic liquids) may pose steric hindrance between CO2-cation
interaction. Moreover, the shrinkage of the microvoid volume, resulting from plasticization
and rigidity due to cross-linking, might cause a decrease in CO2 sorption capacity [89].
1.5.6. Toxicity of ILs
Due to negligible volatility, ionic liquids are not supposed to contaminate air, yet most of
these, being water soluble, may pollute the hydrosphere via industrial effluents or
Chapter 1
29
accidental leakages. Though considerable work is being done regarding the physical,
thermodynamic, kinetic or engineering aspects; comparatively much less data is available
about toxicology of ionic liquids. Bernot et al. [90] investigated the toxic behavior of
certain imidazolium- and pyridinium-based ionic liquids. ILs with longer alkyl side chains
showed higher toxicity (Table 1.12). Similar behavior was found by Wells and Coombe
about the role of alkyl-side chain length [91]. Nevertheless, in order to fully understand the
role of cation/anion towards toxicity more sturdy analysis is needed [92].
Table 1. 12. Lethal concentrations (LC50) of different ionic liquids to fresh water snail
(Physa acuta) in 96-hour acute toxicity exposures.*
Ionic Liquid Alkyl chain length
(carbon atoms)
LC50a
(mg/dm3)
[omp]Br 8 1.0
[omim]Br 8 8.2
[hmim]Br 6 56.2
[bmim]PF6 4 123.3
Tetrabutyl phosphonium Br 4 208.0
[hmp]Br 6 226.7
[bmim]Br 4 229.0
[bmp]Br 4 325.2
Tetrabutyl ammonium Br 4 580.2 *adapted from [90];
a LC50 is the concentration large enough to kill 50% of a sample of animals under test
The studies about the toxic nature of ionic liquids reveal that they cannot be classified as
green media without proper evaluation. Prior to large scale employment, this very aspect
need more extensive investigations so that truly green ionic liquids could evolve for the
purpose by taking advantage of their tunable nature [93].
1.5.7. Current and future developments
This brief survey on the current trends on the ionic-liquid mediated CO2 capture suggests
that CO2 capture by ionic liquids is feasible. A variety of ionic-liquid techniques involving
RTILs, TSILs or SILMs can be employed for CO2 capture, extending from low to high
temperature applications. Some of the benefits/downsides discussed in this section are
presented in Table 1.13.
At present, the lack of availability of inexpensive and diverse ionic liquids is the major
cause of hesitation in employing ionic liquid systems for large scale CO2 capture. Also, in
Chapter 1
30
spite of numerous studies on CO2 solubility and its selectivity, systems mimicking
industrial effluents, where the presence of water or other foreign molecules can affect CO2
transfer, yet requires in depth investigations before industrial-scale implementation of ionic
liquids is sought. Selection of an appropriate combination of the constituent ion pair (cation
+ anion) of ionic liquids, particularly in the context of viscosity and gas absorption kinetics,
needs to be further scrutinized. Aspects related to the gas capture at higher temperatures
and higher pressures, and subsequent regeneration without any appreciable loss and/or
degradation as well as toxicological concerns call for intense analysis to take advantage of
long-lasting cyclic use of IL-based scrubbers. An appropriate balance between cost and
performance is crucial in order for these approaches to take any helm as commercially
viable CO2 capture technologies.
Though few pilot projects for evaluating the ability of ionic liquids in a wider scope are in
progression, gas capture data is not available. Ion Engineering Company, founded by
scientists of Colorado University, possesses demonstration facility and intended to use the
knowhow of ionic liquids for industrial-scale sweetening of natural gas and flue gas CO2
separation [94-96]. Nevertheless, by taking advantage of the tunable nature of ionic liquids,
more meticulous efforts are needed to make them well-adapted and efficient enough for
adequately capturing CO2 from large point sources.
Chapter 1
31
Table 1. 13. Summary of CO2 capture by ionic liquids. Type Examples Equilibrium time Advantages Drawbacks References
Bulk RTILs [hmpy][Tf2N] (32.8a);
[hmim][Tf2N] (31.6a);
[bmim][ Tf2N] (33.0a);
[bmim][PF6] (53.4a);
[bmim][BF4] (59.0a);
[C6H4F9mim][Tf2N] (28.4a)
>90 min, depending
upon viscosity
Negligible vapor pressure; thermally
stable (even after multiple
absorption/desorption experiments, no
detectable loss in mass occurred [17]);
highly CO2-philic; CO2 capture >90%
High viscosity, so
mass transfer a
major concern;
Longer time to
reach equilibrium
[31]
Bulk TSILs [Amim][BF4];
[Pabim][BF4];
[Am-Im][DCA];
[Am-im]þ[BF4]
≥180 min Functionalization increases the CO2 load
almost three fold; CO2 loading continue
to increase with rise in CO2 pressure; gas
load reached up to 0.5, comparable to
standard amine scrubber.
Extremely high
viscosity ≥2000 cP,
undergo further
increase by CO2
complexation;
Much longer
equilibrium time;
exceptionally long
regeneration time
≥24 hours.
[46-49]
RTILs
based
SILMs
[bmim][BF4] + PVDF _ Extremely low volatility prevents solvent
loss; CO2/CH4 selectivity 25-45; CO2/N2
selectivity ≥127; CO2/H2 selectivity <10;
better at low temperatures
Higher
temperatures result
in decrease of
selectivity
[53]
TSILs
based
SILMs
[C3NH2mim][CF3SO3] + PTFE;
[C3NH2mim][Tf2N] + PTFE;
[H2NC3H6mim][Tf2N] + Cross-linked
Nylon 66
_ CO2/CH4 selectivity reached 100-120;
CO2/H2 selectivity >15.
Selectivity
increased till 85 °C
and then decreased
with rise in
temperature
[57]
Poly(ionic
liquid)s
P[VBBI][BF4] (26.0a);
P[MABI][BF4] (37.7a);
P[VBBI][Tf2N];
P[VBTMA][BF4] (3.7a: 22 °C);
P[MATMA][BF4] (5.4a: 22 °C)
<60 min Highly selective CO2 absorption,
compared to N2 & O2 (both showed
negligible absorption); Much faster CO2
sorption; Poly(RTIL)s captured twice the
CO2 compared to their liquid
counterparts.
_ [20,60]
a Henry‟s law constant (bar) for CO2 (at 25 °C except where otherwise stated)
Chapter 1
32
1.6. Research Objectives
Till now, in spite of some serious drawbacks (like equilibrium limitations, high energy
consumption during regeneration, corrosion of equipment, solvent loss, etc.), aqueous
alkanolamine based acid gas scrubbing is the most accomplished approach in industry. To
make alkanolamine systems safer and more affordable, it seems attractive to substitute
water with some stable solvent as most of the drawbacks of alkanolamine processes are
incited by the aqueous phase.
In this regard, room-temperature ionic liquids can be promising contenders. These unique
species are rightly considered as designer solvents that possess some unique characteristics
such as negligible volatility, good thermal stability, wide liquid range, etc. Besides, these
have shown significant affinity for CO2.
Accordingly, the overall goal of this project is to develop a scheme by coupling the
advantages of both alkanolamines and RTILs to render alkanolamine based systems more
productive for efficient carbon dioxide capture. This tactic will also help avoid the
induction of chemical functionality, and accompanying drawbacks (as in case of TSILs), to
ionic liquids required to remove low concentrations of CO2 as is encountered in post-
combustion flue gases.
Blending an alkanolamine with a room-temperature ionic liquid may provide a potential
formulation with less problems and enhanced stability of the process; and this concept has
directed us towards the following objectives:
To explore apposite ionic liquids (preferentially hydrophobic room-temperature
ionic liquids), to be used in combination with alkanolamines, that can make the
CO2-captured product (carbamate/carbonate) precipitate out thus enabling the
chemical absorption to continue at higher rate by overcoming the equilibrium
limitations;
To contrive an amine-RTIL combination that can serve alleviate corrosion
occurrence as well as can suppress amine losses due to evaporation/degradation;
Chapter 1
33
To reach a peculiar possibility of regenerating a smaller volume (through
separation of precipitated CO2-captured product) thus promising less energy
consumption.
In the light of above stated goals, four sections (Chapters 2-4) presenting the theoretical
background and experimental results specific to each fundamental study are provided in
this thesis:
In Chapter two CO2 absorption rate in alkanolamine/RTIL emulsions comprising
diethanolamine (DEA) dispersed in hydrophobic 1-alkyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide [Cnmim][Tf2N] was monitored using thermogravimetric
analyzer, whereas carbon steel 1020 was selected to examine the corrosive behavior of the
capture fluid. Chemical nature of the CO2-captured product (carbamate crystals) was
verified through X-ray crystallography.
Chapter three was intended to investigate the nature (hydrophilicity) of ionic liquids with
reference to the CO2 absorption behavior as well as corrosion phenomenon. The influence
of alkanolamine type was also evaluated. The main purpose of the study, phrased in the
first two sections (Chapter 2-3), was to find an apposite alkanolamine-RTIL combination
that would help address the drawbacks of current amine processes adeptly.
Chapter four was apportioned for regeneration studies. Two alkanolamines, 2-amino-2-
methyl-1-propanol (a primary amine), diethanolamine (a secondary amine), and one
hydrophobic ionic liquid, [hmim][Tf2N], were employed to get solid carbamates (CO2-
captured products). The aim was to find an amine-RTIL pair that can help narrow the gap
between CO2 absorption and stripping temperatures. Thermogravimetric analyzer coupled
with quadrupole mass spectrometer, differential scanning calorimetry, 13
C NMR, and ATR-
FTIR techniques were used to investigate the thermal behavior and regeneration
mechanism.
Finally in the 5th chapter kinetic aspects of CO2 absorption in DEA-[hmim][Tf2N] system
was studied. A stirred-cell reactor fitted with a CO2 probe was used to monitor the gas
Chapter 1
34
absorption behavior by varying amine concentration, gas partial pressure, agitation speed,
and temperature.
Acronyms
AMP 2-amino-2-methylpropan-1-ol
[BIEO][BF4] (1-butylimidazolium-3)methylethylene oxide tetrafluoroborate
[bmim][BF4] 1-n-butyl-3-methylimidazolium tetrafluoroborate
[bmim][PF6] 1-n-butyl-3-methylimidazolium hexafluorophosphate
[bmim][Tf2N] or
C4mimTf2N
1-n-butyl-3-methylimidazolium bis[trifluoromethylsulfonyl]imide
C3mimPF6 1-propyl-3-methylimidazolium hexafluorophosphate
C3mimTf2N 1-propyl-3-methylimidazolium bis[trifluoromethylsulfonyl]imide
C6mimTf2N or
[hmim][Tf2N]
1-n-hexyl-3-methylimidazolium bis[trifluoromethylsulfonyl]imide
C8F13mimTf2N 1-methyl-3-(3,3,4,4,5,5,6,6,7,7,8,8,8-tridecafluorooctyl)-
imidazolium bis[trifluoromethylsulfonyl]imide
cP Centipoise
DEA Diethanolamine
DIPA Diisopropanolamine
DMAE N,N-dimethylethanolamine
DMP N,N‟-dimethylpiperazine
DMPEG Dimethyl ether of polyethylene glycol
[emim][Tf2N] 1-ethyl-3-methylimidazolium bis[trifluoromethylsulfonyl]imide
HEEDA N-(2-hydroxyethyl)ethylenediamine
[hmim][PF6] 1-n-hexyl-3-methylimidazolium hexafluorophosphate
hr Hour
ILs Ionic liquids
IPCC Intergovernmental Panel on Climate Change
MAE N-methylethanolamine
MEA Monoethanolamine
Chapter 1
35
MDEA N-methyldiethanolamine
mmpy Millimetre per year
mpy Milliinch per year
NMP N-methyl-2-pyrrolidone
N,N-diMEDA N,N-dimethylethylenediamine
N,N‟-diMEDA N,N‟-dimethylethylenediamine
N,N,N‟-triMEDA N,N,N‟-trimethylethylenediamine
[pabim][BF4] 1-(3-aminopropyl)-3-butylimidazolium tetrafluoroborate
PC Propylene carbonate
PEG Polyethylene glycol
[perfluoro-hmim][Tf2N] 1-(3,4,5,6-perfluorohexyl)-3-methylimdazolium
bis(trifluoromethylsulfonyl)imide
[pmmim][Tf2N] 1,2-dimethyl-3-propylimidazolium bis(trifluoromethylsulfonyl)imide
PTFE Polytetrafluoroethylene
PVDF Poly vinylidene fluoride
RTILs Room-temperature ionic liquids
SILMs Supported ionic liquid membranes
TMEDA N,N,N‟,N‟-tetramethylethylenediamine
TSILs Task-specific ionic liquids
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Chapter 2
45
CO2 capture in alkanolamine/room-temperature ionic liquid emulsions: A
viable approach with carbamate crystallization and curbed corrosion
behavior*
Abstract/Résumé
By making use of alkanolamine/room-temperature ionic liquid emulsions, it has been found
practicable to capture carbon dioxide up to stoichiometric maximum (0.5 mole of CO2 per
mole of diethanolamine) through crystallization of CO2-captured product (DEA-carbamate)
by avoiding equilibrium limitations. This enabled easy separation of a reasonably smaller
(solid carbamate) volume, thus offering cost effective regeneration. The scanning electron
microscopy (SEM) and electrochemical corrosion studies further revealed that inclusion of
ionic liquid helped suppress corrosion to an extent as low as 0.31 milli-inch per year.
Grâce à l‟utilisation de mélanges d‟alcanolamine/liquide ionique à température ambiante, il
a été possible de capturer du CO2 jusqu‟au maximum stoechiométrique (0.5 mol de CO2 par
mol de diéthanolamine) par le biais de la cristallisation du produit formé (carbamate de
DEA) en évitant les limitations d'équilibre. Ceci a permis une séparation aisée d‟un volume
raisonnablement petit de carbamate solide, offrant ainsi une régénération rentable. La
microscopie électronique à balayage (MEB) et des études de corrosion électrochimiques
ont par la suite révélé que l'inclusion du liquide ionique a contribué à supprimer la
corrosion à des valeurs aussi basses que 0,31 millième de pouce par année.
2.1. Introduction
In the energy future driven by greenhouse gas (predominantly CO2) constraints, there are
mounting concerns over global warming phenomenon being intensified brusquely by
anthropogenic activities. Fossil-fuel based power plants are the largest among stationary
sources accounting for approximately 78.6% of carbon dioxide emissions. It is perceived
that by the year 2100 there may be a rise of 1.9 °C in the global temperature [1,2]. This has
* M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Int. J. Greenhouse Gas Control 6 (2012) 246-252.
Chapter 2
46
turned carbon dioxide capture and sequestration into an extensively investigated topic
nowadays. Carbon dioxide capture processes based on aqueous alkanolamines are the most
widely used on industrial scale. Nonetheless, these technologies pose a number of
drawbacks, including: equilibrium limitations [3,4], high regeneration energy penalty [5],
degradation/evaporation of amines [6-8], low gas loadings [9,10], and corrosion of
equipment [11]. In order to lessen the severity of solvent degradation as well as corrosion
phenomenon gas loading is kept low [3]. In addition, certain additives like corrosion
inhibitors, antifoaming agents are also used to alleviate the process snags [12-14]. This
practice not only increases cost but also supplements to toxicity. To deal with such
concerns, proprietary physical solvent processes like rectisol and selexol are being
employed [10]. Even so, these require high partial pressures of feed gas as well as
refrigeration of the gas/solvent.
In order to bring versatility and robustness to the CO2 capture systems, a great deal of work
is being done in exploitation of ionic liquids in carbon dioxide capture. Because of their
unique characteristics, i.e., wide liquid range, thermal stability, negligible vapor pressure,
tunable physicochemical nature, and quite reasonable CO2 solubility, ionic liquids are
considered as green alternates [15]. However, to fully explore capabilities of these
promising fluids more knowledge is needed to come up with cost-efficient practicable CO2
capture methods realistically implementable in industry.
Looking into the current information available about utilization of ionic liquids in CO2
capture, these alone, like common physical solvents, do not appear competitive enough
when compared to gas capture efficiency of aqueous alkanolamine systems. In order to
make a new process more cost effective, it must possess higher ability of attenuating the
drawbacks faced in current state-of-the-art technologies. Coupling advantages of
commodity alkanolamines with those of room-temperature ionic liquids (RTILs) might
provide a better route regarding global efficacy and stability [16].
In the present study, CO2 capture behavior of emulsions comprising immiscible
alkanolamine dispersed in hydrophobic RTIL continuous phase [diethanolamine (DEA)/1-
alkyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide emulsions] was examined.
We opted to use immiscible phases to ascertain:
Chapter 2
47
Easy separation of the solid product
Diminished exposure of CO2 capture sites to concomitant water vapours in flue gas,
through induction of hydrophobic barrier (RTIL continuous phase) between feed
gas and amine droplets (dispersed phase).
A systematic series of experiments have made it possible to investigate the fate of gas-
captured product (carbamate) as well as the corrosive action of the fluid on carbon steel.
The gas absorption profiles were obtained by thermogravimetric analysis and the
precipitated carbamate was analyzed using single crystal X-ray technique. Fourier
transform infrared spectroscopy (FTIR) as well as 13
C NMR methods were employed to
further complement the characterization. Linear polarization and Tafel plots were used to
probe corrosion behaviour. This study may elaborate valuable facts about one of several
cost-effective options to be practicable for industrial scale CO2 capture.
2.2. Experimental
2.2.1. Materials and techniques
1-alkyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ionic liquids (99% purity)
were provided by IoLiTec Inc. while diethanolamine (DEA) and Triton® X-100 were
purchased from EMD Chemicals. All chemicals were used as received without further
purification. Carbon dioxide, nitrogen, oxygen and argon (≥99% purity) were obtained
from Praxair Canada Inc.
Emulsification was carried out using Omni homogenizer (Omni International) fitted with
rotor-stator generator. In case of surfactant stabilized emulsion, [hmim][Tf2N] and Triton®
X-100 were mixed first in 2:3 ratio (w/w). And then, after addition of DEA, stirring was
continued for 3-4 minutes at a speed of 6000 rpm.
CO2 absorption plots were obtained by thermogravimetric analyzer (Perkin-Elmer Diamond
TG/DTA) isothermally under 100% carbon dioxide atmosphere. The product
characterization was done by single crystal X-ray diffraction and 13
C NMR means. The FT-
IR spectra were recorded with a Nicolet Magna 850 spectrometer (Thermo Scientific,
Madison, WI) employing attenuated total reflectance (ATR) technique. Ultrafoam™ 1200e
Chapter 2
48
pycnometer was used to determine density of crystalline solid while AR-G2 rheometer (TA
Instruments) with parallel plate geometry was used for viscosity analysis. Water content
measurements were done by Karl Fischer titrator (784 KFP Titrino, Metrohm AG).
2.2.2. Crystal structure determination
Crystallographic data measurements were made at 200 K on a Bruker APEX II area detector
diffractometer equipped with Mo-Kα monochromated radiation (λ = 0.71073 Å). APEX 2
and SAINT programs were used for retrieving cell parameters and data collection (APEX2
Version 2.0-2, 2005; SAINT Version 7.07a, 2003) [17,18]. Data were corrected for Lorentz
and polarization effects. Face-indexed and multiscan absorption corrections were
performed using XPREP and SADABS programs, respectively (XPREP Version 2005/2,
2005; SADABS Version 2004/1, 2004) [19,20]. The structure was solved and refined by
full-matrix least-squares against F2 using SHELXS-97 and SHELXL-97 programs
(Sheldrick, 1997) [21]. Refinement of all non-hydrogen atoms was done with anisotropic
thermal parameters. The hydrogen atoms were placed at geometrically idealized positions
using a riding model (SHELXTL Version 6.12, 2001) [22]. Neutral atom scattering factors
were taken from International Tables for Crystallography, Vol C, 1992 [23]. This crystal
structure gives a satisfactory checkCIF report.
2.2.3. Electrochemical corrosion tests
Bio-Logic VSP potentiostat was used to evaluate corrosion occurrence utilizing a rotating
disc electrode assembly. The electrochemical tests were conducted by bubbling either pure
CO2 or mixture of CO2+O2 (1:1 ratio by volume) after passing through deionised water in
order to saturate with water vapours.
Carbon steel 1020 was used as working electrode to study the corrosive behavior of
aqueous DEA as well as that of DEA/RTIL emulsion. The setup was comprised of three
electrode assembly, i.e., platinum counter electrode, silver/silver chloride (Ag/AgCl/
sat.KCl) reference electrode, and carbon steel working electrode. A disc shaped working
electrode having surface area of 0.196 cm2 was mounted in Teflon cap. The experiments
were carried out in 100 cm3 volume corrosion cell using an oil bath for temperature control
Chapter 2
49
and a condenser to minimize evaporation of the experimental fluid. For each run, working
electrode surface was successively polished with 600 grit SiC paper and alumina (5 µm
particle size) suspension respectively, followed by sequential degreasing with acetone and
rinsing with deionized water. Each time, during an electrochemical test, bubbling of
respective gas/gaseous mixture was initiated at a flow rate of 70 cm3/min under
atmospheric pressure, one hour prior to the commencement of polarization run. After the
accomplishment of required conditions a computer controlled potentiostat was used to carry
out linear polarization resistance (LPR) measurements starting from a cathodic potential of
-250 mV to an anodic potential of +250 mV (versus open circuit potential) with a scan rate
of 0.166 mVs-1
. In all the experiments, 500 rpm of rotation speed was maintained for the
working electrode. During this practice, the influence of O2 as well as that of temperature
on corrosion rate was assessed for aqueous solutions of DEA. And then the most severe
conditions tested in case of aqueous DEA were used to evaluate the corrosion behavior of
carbon steel in DEA/RTIL emulsions.
2.3. Results and discussion
2.3.1. Fate of CO2-captured product (carbamate)
To know the behaviour of product (carbamate), CO2 was bubbled through 30% w/w
DEA/RTIL emulsions for 2 hours at 25 °C and atmospheric pressure amid continued rotor-
stator stirring (2000 rpm). The gas capture resulted in precipitation of carbamate in each of
three categories involving [emim][Tf2N], [bmim][Tf2N] or [hmim][Tf2N] ionic liquids
(Table 2.1). In case of DEA/[emim][Tf2N] and DEA/[hmim][Tf2N] schemes, the solid
phase rose to the surface rather promptly thus making it quite easy to be separable,
promising considerably lesser volume to regenerate, as established in Figure 2.1(a-c). This
trend depicts that RTIL hydrophobicity as well as density difference (between solid and
liquid phases) are responsible for carbamate crystals to easily move out of the liquid as a
supernatant solid. Yet, hydrophobic nature of the ionic liquid appeared to be dominating
factor in segregation of solid product from the fluid phase, which was quite evident from
the carbamate orientation in [hmim][Tf2N] based system, in spite of markedly higher
viscosity of RTIL (compared to that of [emim][Tf2N]) and minor density difference
between the solid/liquid phases. Nevertheless, in surfactant (Triton®
X-100) stabilized
Chapter 2
50
emulsion, the carbamate product remained dispersed transforming emulsion into
suspension (Figure 2.1d).
Figure 2. 1. DEA/RTIL system: (a-c) (without surfactant) after CO2 capture; d) (with
surfactant) before and after CO2 capture.
2.3.2. CO2 absorption
Isothermal absorption of CO2 in surfactant (Triton® X-100) stabilized DEA/RTIL emulsion
was carried out using thermogravimetric analyzer. The results demonstrated the prospect of
maximum gas loading capacity (0.5 mole of CO2 per mole of DEA) of this novel scheme,
without undergoing any momentous effect of equilibrium restraint. Regarding the gas
capture rate, thermogravimetric analysis does not show much variation in CO2 uptake array
in case of three emulsion types. The slight disparity seemed to arise from difference in
viscosities of the three ionic liquids (Table 2.1; Figure 2.2).
Table 2. 1. Density (ρ) and viscosity (η) values measured at 25 °C.
Abbreviation Name ρ (g/cm3) η (cP)
[emim][Tf2N] 1-ethyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide
1.52 34.1
[bmim][Tf2N] 1-butyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide
1.44 49.6
[hmim][Tf2N] 1-hexyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide
1.37 73.5
C9H22N2O6* DEA-carbamate (CO2-captured product) 1.36 -
DEA Diethanolamine 1.09 469 *empirical formula
a) b) (a)
DEA/[emim][Tf2N]
(b) DEA/[bmim][Tf2N]
(c) DEA/[hmim][Tf2N]
(d) DEA/[hmim][Tf2N]
Chapter 2
51
Figure 2. 2. CO2 capture capacity profiles of DEA/RTIL system (surfactant stabilized
emulsions; 30% w/w) at atmospheric pressure and 25 °C.
However, as shown in Figure 2.3, increase in diethanolamine (DEA) ratio from 15% to
30% (w/w) resulted in relatively slower kinetics of the process. This behaviour was
expected due to decreased diffusivity [24], owing to greater proportion of more viscous
DEA.
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150 200 250 300 350
Mo
le r
ati
o C
O2/D
EA
Time (min.)
[EMIM][Tf2N]
[BMIM][Tf2N]
[HMIM]
30% DEA/[EMIM][Tf2N]
30% DEA/[BMIM][Tf2N]
30% DEA/[HMIM][Tf2N]
Chapter 2
52
Figure 2. 3. CO2 absorption isotherms for DEA/[hmim][Tf2N] surfactant stabilized
emulsions obtained at 25°C.
The CO2 capture by diethanolamine (DEA) possibly involves a fairly rational mechanism
(1) comprising direct interaction of amine with CO2 forming zwitterion followed by
abstraction of proton, thus consuming a second amine molecule to act as a counter ion to
induce stability to carbamate [25-27].
The crystallization of the carbamate product enabled the process to reach completion
avoiding any equilibrium limitations specifically faced in aqueous amine systems.
Furthermore, separation of carbamate solid would provide an imperative opportunity in
reducing regeneration costs.
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150 200 250 300 350
Mo
le r
ati
o C
O2/D
EA
Time (min.)
15% DEA
30% DEA
15% DEA/[HMIM][Tf2N]
30% DEA/[HMIM][Tf2N]
CO2 + RR'NH RR'NH+CO2
-
RR'NH+CO2
- + RR'NH RR'NCO2-+ RR'NH2
+(1)
Chapter 2
53
2.3.3. Characterization of crystalline product
Superior hydrophobicity [28], relatively higher CO2 solubility by virtue of longer alkyl
side-chain [29,30], as well as reasonably good separation of solid carbamate from liquid
phase, eased the selection of [hmim][Tf2N] as continuous phase in the process for further
evaluation specifically in DEA-carbamate characterization and corrosion studies.
As there is no involvement of water in the gas capturing fluid, single crystal analysis
established that there was no question of bicarbonate or carbonate species (typically found
in aqueous amine systems). CO2 absorption occurred only through carbamate formation
resulting in 50 mol % mass increase (w.r.t. DEA) as confirmed by the thermogravimetric
analysis. The crystallographic information is summarized in Table 2.2.
Table 2. 2. Crystallographic data
DEA Carbamate
Empirical formula C9H22N2O6
Moiety formula C5H10NO4, C4H12NO2
Formula weight (M) 254.29
Temperature 200(2) K
Crystal dimensions 0.47x0.11x0.09 mm3
Crystal system Monoclinic
Space group Pn
Unit cell dimensions a = 10.6841(7) Å α= 90°
b = 4.6017(3) Å β= 99.8990°(10)
c = 12.8334(8) Å γ= 90°
Unit cell volume 621.56(7) Å3
No. of formula units in unit cell (Z) 2
F(000) 276
θ range for data collection 2.30° to 27.00°
Completeness to θ = 27.0° 99.8 %
Reflections collected 6808
Independent reflections 1356 [R(int)=0.0183]
Observed reflections 1329 [I>2σ(I)]
R indices (all data) 0.0249
Final R indices [I>2σ(I)] 0.0244
Density (calculated) 1.359 g/cm3
Absorption coefficient 0.113 mm-1
hkl range -13≤h≤13, -5≤k≤5, -16≤l≤16
Refinement method Full-matrix least-squares on F2
Data/restraints/parameters 1356/2/158
Goodness-of-fit on F2 1.266
Chapter 2
54
The basic structural unit is composed of protonated-DEA cation and DEA-carbamate anion,
as shown in Figure 2.4.
Figure 2. 4. Basic structural unit in DEA-carbamate (C9H22N2O6) crystal.
The single crystal X-ray structure analysis has shown that the packing mode is monoclinic
with Pn space group. The average lengths of both O(1)-C(1) and O(2)-C(1) bonds (1.2695Å
and 1.2764Å, respectively; Table A.1 in Appendix A) in carbamate anion are quite identical
depicting the occurrence of delocalization. Attachment of CO2- moiety (captured CO2)
caused to decrease the N(1)-C(2) and N(1)-C(4) bond lengths (1.4596Å and 1.4633Å)
compared to the respective bonds in counter cation (protonated amine), also evident in 13
C
NMR spectra owing to dissimilar environments.
Table 2. 3. Relevant hydrogen bonding parameters [bond distances (Å) and angles (°)].
D―H···A d(D―H) d(H···A) d(D···A) ∠(D―H···A) Symmetry operators*
O(3)―H(3)···O(4) 0.84 1.96 2.7947(19) 176.3 x-1/2,-y+1,z-1/2
O(4)―H(4)···O(5) 0.84 1.92 2.7362(17) 163.2 x,y-1,z
O(5)―H(5)···O(1) 0.84 1.84 2.6650(15) 168.5 x+1/2,-y+2,z+1/2
O(6)―H(6)···O(1) 0.84 1.88 2.7136(16) 173.6 x+1/2,-y+2,z+1/2
N(2)―H(2A)···O(2) 0.92 1.94 2.8133(16) 158.3 x,y+1,z
N(2)―H(2B)···O(2) 0.92 1.90 2.7949(16) 163.4 x,y+1,z
D: donor atom; A: acceptor atom; *Symmetry operators used to generate equivalent acceptor atoms
Chapter 2
55
Ionic interactions as well as intensive hydrogen bonding in the crystalline carbamate make
impossible for the ionic liquid (1-alkyl-3-methylimidazolium bis(trifluoromethylsulfonyl)
imide) to dissolve it. The bond distances as well as bond angles, elaborating on the
hydrogen bonding configuration, are listed in Table 2.3. Figure 2.5 shows the hydrogen
bonding pattern involving two hydrogen bonds for each oxygen of CO2- moiety in
carbamate anion (see also Figure A.1 in Appendix A). One oxygen is an acceptor of
hydrogen bonding from OH of two cations while the other acquires hydrogen bonds from
NH2 of two different cations, one of these cations being the same involved in hydrogen
bonding with first oxygen of the CO2- moiety. A terminal OH of the anion is hydrogen
bonded to the equivalent site of neighboring anion while the second terminal oxygen (of
OH moiety) forms two hydrogen bonds; with terminal OH of another anion as well as with
terminal OH of a cation. Likewise, cation bears five hydrogen bonds. Out of these, two are
created between cationic NH2 and CO2- moieties of two different anions whereas additional
two are formed by one of terminal oxygens (of OH group) with respective CO2- and OH
moieties of two neighboring anions. The remaining terminal OH forms only one hydrogen
bond with CO2- moiety of a nearby anion.
Figure 2. 5. Hydrogen bonding pattern in the compound (DEA-carbamate). H atoms not
participating in hydrogen bonding are omitted for clarity.
Chapter 2
56
The 13
C NMR spectrum (taken in DMSO-d6) of crystalline carbamate displays four peaks
in the range of 50.84-61.35 ppm. Out of these, two comparatively more intense peaks (one
at 50.843 ppm and another at 58.588 ppm) arise from CH2-CH2 carbons of protonated
amine (DEAH+) while two low intensity signals at 51.331 and 61.347 ppm originate from
ethylene carbons of carbamate derivative. A low intensity resonance at 162.57 ppm
confirms the emergence of carbamate carbon resulting from CO2 capture, as shown in
Figure 2.6 (see also Figures A.2-A.5 in Appendix A). FT-IR technique further validates the
existence of carbamate moiety appearing as carbonyl stretching frequency at 1654.68 cm-1
in Figure 2.7 (see also Figures A.6-A.8, Appendix A).
Figure 2. 6. 13
C NMR spectrum of crystalline carbamate (retaining traces of [hmim][Tf2N])
taken in DMSO-d6 solvent.
PPM 180.0 160.0 140.0 120.0 100.0 80.0 60.0 40.0 20.0
162.5
700
61.3
468
58.5
883
51.3
312
50.8
435
61.35 – 50.84
16
2.5
7
OO-
OHN
OH
NH2
+OH OH
Chapter 2
57
Figure 2. 7. FTIR analysis of crystalline product (DEA-carbamate).
2.3.4. Corrosion studies
Tafel analysis was accomplished using the extrapolation mode to determine corrosion
current (icorr) which in turn enabled to calculate the corrosion rate, CR:
Where CR is in milli-inches per year (mpy), icorr is the corrosion current in Amperes, W is
equivalent mass of metal specimen in gram per equivalent, ρ is the density of metal in
g/cm3 and A is the area (in contact with experimental fluid) of the rotating disc working
electrode in cm2.
Figure 2.8 presents the Tafel plots generated by performing anodic polarization runs for
aqueous amine solutions under different environments. At lower pH (~ 8) resulting from
CO2 absorption, high temperatures as well as presence of oxygen adjoined to detrimental
approach towards corrosion of steel. By increasing the temperature from 25 °C to 60 °C
5(1.29 10 ) corri WCR
A
Chapter 2
58
(through 35 °C only), the corrosion rate augmented by more than three-fold. Elevated
temperature facilitated fast distribution of corrosion products whereas inclusion of oxygen
increased the concentration of oxidizing species, escalating the chances of iron oxidation
and thus accelerating the corrosion process.
Figure 2. 8. Tafel plots for carbon steel electrode in aqueous DEA under different
environments: a) CO2 bubbling at 25 °C, b) CO2+O2 bubbling at 25 °C, c) CO2 bubbling at
60 °C, d) CO2+O2 bubbling at 60 °C.
Potentiodynamic experiments exhibited towering corrosion rate in case of aqueous
diethanolamine (15% w/w) rendering the addition of corrosion inhibitors a mandatory
activity that not only adds to the cost but also makes the solvent more toxic [12]. The major
anodic and cathodic electrochemical reactions occurring in aqueous amine systems during
corrosion phenomenon are written below [31].
a) Anodic reaction
Fe → Fe2+
+ 2e- (oxidation of iron) (2)
b) Cathodic reactions
2H2O + 2e- → 2OH
- + H2↑ (3)
0
0.5
1
1.5
2
2.5
3
3.5
-900 -800 -700 -600 -500
log
I (
µA
)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
Scan rate: 0.166 mV/s
ab
cd
Chapter 2
59
2HCO3- + 2e
- → 2CO3
2- + H2↑ (4)
O2 + 2H2O + 4e- → 4OH
- (5)
c) Corrosion products
Fe2+
+ 2OH- → Fe(OH)2 (6)
Fe2+
+ CO32-
→ FeCO3 (7)
By replacing aqueous part with hydrophobic room-temperature ionic liquid, [hmim][Tf2N],
it has been possible to reduce corrosion virtually to negligible (Table 2.4; Figure 2.9).
Exclusion of water truncated the probable oxidizers mainly responsible for cathodic
reactions (equations 3-5) in aqueous media. This behavior suggests that the RTIL,
[hmim][Tf2N], was stable under the investigated conditions and did not take part in any of
the corrosion-related electrochemical reactions. Thus RTIL not only enabled carbamate
product to crystallize out but also made it possible to evade the addition of costly and toxic
corrosion inhibitors.
Table 2. 4. Corrosion rates of carbon steel 1020*
*Density: 7.86 g.cm-3
; Composition (weight %): 0.20% carbon, 0.50% manganese, 0.04% phosphorus, 0.05%
sulfur, balanced by iron. #Just prior to the start of gas bubbling.
SEM (scanning electron microscope) micrographs of the working electrodes‟ surfaces
before and after electrochemical corrosion tests under CO2-O2-H2O(vap.) atmosphere at 60
°C further confirmed the absence of corrosion in case of DEA/RTIL emulsion. Though in
aqueous DEA, deterioration of steel is quite evident in Figure 2.10.
Medium Environment Temperature Corrosion
Potential
(mV)
Corrosion
Current
(µA)
Corrosion
Rate
(mpy)
Water content
(% w/w)
Before
electrochemical
run#
After
electrochemical
run
DEA (aq)
, 15% CO2 25 °C -729 40.95 95.60 - -
DEA (aq)
, 15% CO2+O
2 25 °C -604 80.37 187.62 - -
DEA (aq)
, 15% CO2 60 °C -766 124.15 289.82 - -
DEA (aq)
, 15% CO2+O
2 60 °C -688 137.87 321.84 - -
RTIL (Pure) CO2+O
2+H
2O
(vap.) 60 °C 116 0.05 0.11 0.02 0.32
DEA/IL
emulsion
CO2+O
2+H
2O
(vap.) 60 °C -157 0.14 0.31 0.12 0.73
Chapter 2
60
Figure 2. 9. Corrosion rate of carbon steel 1020 in: a) RTIL pure, CO2+O2+H2O(vap.)
bubbling at 60 °C, b) DEA/RTIL emulsion, CO2+O2+H2O(vap.) bubbling at 60 °C, c)
DEA(aq), CO2+O2 bubbling at 60 °C.
Figure 2. 10. SEM micrographs of working electrode specimen. In DEA/RTIL emulsion
(15% w/w): a) Fresh surface; b) after electrochemical corrosion test. In DEAaq. (15% w/w):
c) Fresh surface; d) after electrochemical corrosion test.
-5
-4
-3
-2
-1
0
1
2
3
4
-1000 -800 -600 -400 -200 0 200 400
log
I (
µA
)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
Scan rate: 0.166 mV/s
ab
c
a) b)
c) d)
X500 10µm X500 10µm
X500 10µm X500 10µm
Chapter 2
61
2.4. Conclusions
In order to accomplish a more efficient scheme for CO2 capture, we have been able to
devise a process by combining advantages of both immiscible alkanolamine (superior CO2
capture efficiency) and hydrophobic room-temperature ionic liquid (excellent thermal
stability and practically no volatility). Scheming emulsions with RTIL as continuous phase
bearing dispersed alkanolamine droplets may offer a potential opportunity with less CO2
capture cost and enhanced process stability. This has been quite evident from our
experimental results for CO2 capture and corrosion rate measurements. Enabling carbamate
(CO2-captured product) to crystallize out of the continuous phase, it has been possible to
run the process at higher rates reaching maximum gas loading capacity, thus avoiding
equilibrium limitations - a major obstacle in case of aqueous alkanolamines. The
insolubility of the product also offers the advantage of regenerating a smaller volume with
less energy consumption. Negligible corrosion phenomenon further helps establish the
benefit of alkanolamine/RTIL emulsions. Though stabilization (through surfactant addition)
of emulsion was required for time-consuming thermogravimetric/electrochemical
experimentation, carbamate separation and consequently amine regeneration appeared to be
far easier and hence cost-effective without the use of Triton® X-100. In addition,
hydrophobic barrier of RTIL continuous phase might help eliminate the dehydrating step
during subsequent regeneration of amine and stripping of pure CO2 from thermal heating of
the recovered solid carbamate cake.
2.5. References
[1] C. Stewart, M-A. Hessami, A study of methods of carbon dioxide capture and
sequestration–the sustainability of a photosynthetic bioreactor approach, Energ. Convers.
Manage. 46 (2005) 403-420.
[2] B. Metz, O. Davidson, H. de Coninck, M. Loos, L. Meyer, Eds., IPCC Special Report
on Carbon Dioxide Capture and Storage, Prepared by Working Group III of the
Intergovernmental Panel on Climate Change, Cambridge University Press, New York,
2005, pp. 51-74.
Chapter 2
62
[3] F. Barzagli, F. Mani, M. Peruzzini, Continuous cycles of CO2 absorption and amine
regeneration with aqueous alkanolamines: a comparison of the efficiency between pure and
blended DEA, MDEA and AMP solutions by 13
C NMR spectroscopy, Energy Environ. Sci.
3 (2010) 772-779.
[4] S. Bishnoi, G.T. Rochelle, Absorption of carbon dioxide into aqueous piperazine:
reaction kinetics, mass transfer and solubility, Chem. Eng. Sci. 55 (2000) 5531-5543.
[5] R. Idem, M. Wilson, P. Tontiwachwuthikul, A. Chakma, A. Veawab, A. Aroonwilas, D.
Gelowitz, Pilot plant studies of the CO2 capture performance of aqueous MEA and mixed
MEA/MDEA solvents at the University of Regina CO2 Capture Technology Development
Plant and the Boundary Dam CO2 Capture Demonstration Plant, Ind. Eng. Chem. Res. 45
(2006) 2414-2420.
[6] A. Bello, R.O. Idem, Pathways for the formation of products of the oxidative
degradation of CO2-loaded concentrated aqueous monoethanolamine solutions during CO2
absorption from flue gases, Ind. Eng. Chem. Res. 44 (2005) 945-969.
[7] B.R. Strazisar, R.R. Anderson, C.M. White, Degradation pathways for
monoethanolamine in a CO2 capture facility, Energy Fuels 17 (2003) 1034-1039.
[8] J. Davis, G.T. Rochelle, Thermal degradation of monoethanolamine at stripper
conditions, Energy Procedia 1 (2009) 327-333.
[9] F. Mani, M. Peruzzini, P. Stoppioni, CO2 absorption by aqueous NH3 solutions:
speciation of ammonium carbamate, bicarbonate and carbonate by a 13
C NMR study, Green
Chem. 8 (2006) 995-1000.
[10] A.L. Kohl, R.B. Nielsen, Gas Purification, 5th ed. Gulf Publishing Company,
Houston, Texas, 1997.
[11] N. Kladkaew, R. Idem, P. Tontiwachwuthikul, C. Saiwan, Corrosion behavior of
carbon steel in the monoethanolamine-H2O-CO2-O2-SO2 system: products, reaction
pathways, and kinetics, Ind. Eng. Chem. Res. 48 (2009) 10169-10179.
Chapter 2
63
[12] A. Veawab, P. Tontiwachwuthikul, A. Chakma, Investigation of low-toxic organic
corrosion inhibitors for CO2 separation process using aqueous MEA solvent, Ind. Eng.
Chem. Res. 40 (2001) 4771-4777.
[13] S. Zhou, X. Chen, T. Nguyen, A.K. Voice, G.T. Rochelle, Aqueous ethylenediamine
for CO2 capture, Chem. Sus. Chem. 3 (2010) 913-918.
[14] X. Chen, S.A. Freeman, G.T. Rochelle, Foaming of aqueous piperazine and
monoethanolamine for CO2 capture, Int. J. Greenh. Gas Control 5 (2011) 381-386.
[15] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Ionic liquids for CO2 capture- Development
and progress, Chem. Eng. Process. 49 (2010) 313-322.
[16] D. Camper, J.E. Bara, D.L. Gin, R.D. Noble, Room-temperature ionic liquid-amine
solutions: Tunable solvents for efficient and reversible capture of CO2, Ind. Eng. Chem.
Res. 47 (2008) 8496-8498.
[17] APEX2 Version 2.0-2, Bruker AXS Inc., Madison, WI, USA, 2005.
[18] SAINT Version 7.07a, Bruker AXS Inc., Madison, WI, USA, 2003.
[19] XPREP Version 2005/2, Bruker AXS Inc., Madison, WI, USA, 2005.
[20] SADABS Version 2004/1, Bruker AXS Inc., Madison, WI, USA, 2004.
[21] G.M. Sheldrick, SHELXS-97 and SHELXL-97, Programs for the refinement of crystal
structures, University of Göttingen, Germany, 1997.
[22] SHELXTL Version 6.12, Bruker AXS Inc., Madison, WI, USA, 2001.
[23] A.J.C. Wilson, Ed., International Tables for Crystallography, Vol. C, Kluwer
Academic Publishers, Dordrecht, 1992, pp. 219-222, 500-502.
[24] S.S. Moganty, R.E. Baltus, Diffusivity of carbon dioxide in room-temperature ionic
liquids, Ind. Eng. Chem. Res. 49 (2010) 9370-9376.
Chapter 2
64
[25] M. Caplow, Kinetics of carbamate formation and breakdown, J. Am. Chem. Soc. 90
(1968) 6795-6803.
[26] P.V. Danckwerts, The reaction of CO2 with ethanolamines, Chem. Eng. Sci. 34 (1979)
443-446.
[27] P.S. Kumar, J.A. Hogendoorn, G.F. Versteeg, P.H.M. Feron, Kinetics of the reaction
of CO2 with aqueous potassium salt of Taurine and Glycine, AIChE J. 49 (2003) 203-213.
[28] M.G. Freire, C.M.S.S. Neves, P.J. Carvalho, R.L. Gardas, A.M. Fernandes, I.M.
Marrucho, L.M.N.B.F. Santos, J.A.P. Coutinho, Mutual solubilities of water and
hydrophobic ionic liquids, J. Phys. Chem. B 111 (2007) 13082-13089.
[29] J.L. Anderson, J.K. Dixon, J.F. Brennecke, Solubility of CO2, CH4, C2H6, C2H4, O2,
and N2 in 1-hexyl-3-methylpyridinium bis(trifluoromethylsulfonyl)imide: Comparison to
other ionic liquids, Acc. Chem. Res. 40 (2007) 1208-1216.
[30] D. Almantariotis, T. Gefflaut, A.A.H. Padua, J.-Y. Coxam, M.F. Costa Gomes, Effect
of fluorination and size of the alkyl side-chain on the solubility of carbon dioxide in 1-
alkyl-3-methylimidazolium bis(trifluoromethylsulfonyl)amide ionic liquids, J. Phys. Chem.
B 114 (2010) 3608-3617.
[31] I.R. Soosaiprakasam, A. Veawab, Corrosion and polarization behavior of carbon steel
in MEA-based CO2 capture process, Int. J. Greenh. Gas Control 2 (2008) 553-562.
Chapter 3
65
Corrosion behaviour of carbon steel in alkanolamine/room-temperature
ionic liquid based CO2 capture systems*
Abstract/Résumé
To address the drawbacks of aqueous alkanolamine based state-of-the-art technology for
industrial scale carbon dioxide capture, among a number of options; alkanolamine/room-
temperature ionic liquid (RTIL) systems are also being tested as a likely replacement.
These new schemes seem to be a better alternative to hamper corrosion occurrence.
Omission of the aqueous phase marks abolition of probable oxidizing species mainly
responsible for corrosion in water-based chemical absorption processes. In the present
study, corrosion phenomenon in amine/room-temperature ionic liquid blends comprising
alkanolamine/s (monoethanolamine, 2-amino-2-methyl-1-propanol, diethanolamine, N-
methyldiethanolamine) and hydrophilic room-temperature ionic liquid ([bmim][BF4],
[emim][BF4], and [emim][Otf]) has been investigated by systematically probing the effect
of amine/RTIL type, process temperature, CO2 loading, presence/absence of oxygen in flue
gas as well as the influence of water content. The analytical techniques exercised in this
regard include linear polarization resistance (LPR), scanning electron microscopy (SEM),
and energy-dispersive X-ray spectroscopy (EDX).
Pour éviter les inconvénients des technologies de capture du dioxyde de carbone à l‟échelle
industrielle basé sur l‟utilisation d‟alcanolamines aqueuses, parmi un nombre d'options, les
systèmes alcanolamine/liquide ionique à température ambiante (RTIL) sont également
testés comme substituts potentiels. Ces nouveaux systèmes se révèlent être une alternative
intéressante pour endiguer l‟apparition de la corrosion. L‟absence de phase aqueuse marque
la pénurie en espèces oxydantes qui sont principalement responsables de la corrosion dans
les procédés d‟absorption chimique utilisant de l‟eau. Dans la présente étude, le phénomène
de corrosion au sein d‟un mélange amine/liquide ionique à température ambiante
comprenant des alcanolamines (monoéthanolamine, 2-amino-2-methyl-1-propanol,
* M. Hasib-ur-Rahman, H. Bouteldja, P. Fongarland, M. Siaj, F. Larachi, Ind. Eng. Chem. Res. 51 (2012)
8711-8718.
Chapter 3
66
diéthanolamine, et/ou N-méthyldiéthanolamine) et liquide ionique à température ambiante
hydrophile ([bmim][BF4], [emim][BF4] ou [emim][Otf]) a été étudié en testant
systématiquement l'effet du type de mélange amine/RTIL, la température du procédé, la
concentration en CO2, la présence/absence d‟oxygène dans les gaz de combustion, de même
que l‟influence de la présence d‟eau. Les techniques d‟analyse utilisées à cet égard
comprennent la polarisation linéaire, la microscopie électronique à balayage (MEB) et la
spectroscopie aux rayons X à dispersion d'énergie.
3.1. Introduction
As fossil fuels are supposed to sustain as a major energy source at least until the middle of
the 21st century, global warming largely resulting from anthropogenic emissions of carbon
dioxide remains a matter of great concern [1]. Carbon dioxide capture and storage is a
viable solution to ensure a prevised fall in CO2 emissions from large point sources
involving fossil fuel combustion. In this regard, aqueous alkanolamine systems offer a
promising near-term solution, particularly, in natural gas sweetening and post-combustion
capture from flue gases containing low CO2 concentrations [1,2]. However, these face some
severe operational hitches. Corrosion is one of the major impediments in this regard,
principally due to the presence of water phase [3]. Corrosion products not only trigger the
catalytic degradation of amine but also incite deterioration of plant equipment. A number of
factors like amine concentration, elevated process temperature, and high gas loading can
cause amplification of the corrosion phenomenon [4]. Various types of corrosion inhibitors
such as compounds of copper and vanadium are being used to prevent equipment decay [5-
7]. However, most of these species are not only toxic to both life and the environment but
also add to the cost. Substituting the aqueous phase with a more stable counterpart in the
case of amine based processes may be a better alternative.
Room-temperature ionic liquids (RTILs), generally possessing a tunable nature, greater
thermal stability, and practically no volatility even at elevated temperatures, are emerging
as promising aspirants [8-15]. Significant work is being done to explore the viability of
amine/RTIL blends in CO2 capture facilities, thus combining advantages of both the
counterparts [16,17]. This approach has been found to evade equilibrium limitations owing
to carbamate precipitation in case of primary/secondary alkanolamine based systems,
Chapter 3
67
enabling stoichiometric-maximum gas loading. Likewise, the inclusion of RTIL is
reckoned to alleviate corrosion since in their pure form RTILs can reduce corrosion rates
well below 1 mpy (milli-inch per year) [18]. Thus the use of amine/RTIL blends can
address the CO2 capture problem more efficiently.
In the present work we have probed the corrosion behavior of carbon steel 1020 in
alkanolamine-RTIL mixtures under diverse process parameters and compared the results
with corresponding aqueous amines under alike experimental conditions (summarized in
Table 3.1). Electrochemical linear polarization technique and Tafel fit were used to
determine the corrosion rates. Scanning electron micrographs were taken to observe the
corrosive effect of the media on the working electrode (carbon steel specimen) surface,
while EDX has provided vital information about the role of RTIL in curbing corrosion.
Table 3. 1. Summary of process parameters/conditions.
Parameter Condition
Amine type Monoethanolamine (MEA), 2-amino-2-methyl-1-
propanol (AMP), Diethanolamine (DEA), N-
Methyldiethanolamine (MDEA)
RTIL type 1-butyl-3-methylimidazolium tetrafluoroborate
[bmim][BF4], 1-ethyl-3-methylimidazolium
tetrafluoro borate [emim][BF4], 1-ethyl-3-
methylimidazolium trifluoromethanesulfonate
[emim][Otf]
Amine concentration (kmol/m3) 5.0
Fluid temperature (°C) 25, 60
CO2 loading (mol CO2 per mol amine) Aqueous MEA MEA-RTIL
0.23, 0.42, 0.53 0.24, 0.35, 0.50
O2 content in simulated flue gases (volume %) 0.0, 5.0, 10.0
Water content in RTIL based test fluid (kmol/m3) 0.0, 5.0
Gas flow rate 100 ml/min.
Working electrode rotation speed (rpm) 500
3.2. Experimental
3.2.1. Materials
1-butyl-3-methylimidazolium tetrafluoroborate, 1-ethyl-3-methylimidazolium
tetrafluoroborate, and 1-ethyl-3-methylimidazolium trifluoromethanesulfonate ionic liquids
(99% purity) were purchased from IoLiTec Inc. Monoethanolamine, 2-amino-2-methyl-1-
propanol, diethanolamine and N-Methyldiethanolamine were provided by Sigma-Aldrich
Chapter 3
68
Canada Ltd. All the chemicals were used without further purification. Carbon dioxide,
nitrogen, oxygen and argon (≥99.8% purity) were obtained from Praxair Canada Inc.
3.2.2. Experimental techniques and procedure
3.2.2.1.CO2 capture studies
CO2 absorption profiles were obtained by a thermogravimetric analyzer (Perkin-Elmer
Diamond TG/DTA) isothermally under 100% carbon dioxide atmosphere at 25 °C. A 100
mL/min gas flow rate was exercised during CO2 absorption quantifications using mass flow
controllers. MEA-RTIL samples containing 5 kmol/m3 of amine were used to estimate the
influence of RTIL type on gas loading capacity of the media. Mass uptake was measured to
calculate the CO2 capture capability of amine-RTIL mixtures in terms of molar absorption
ratio.
3.2.2.2.Corrosion studies
A Bio-Logic VSP potentiostat was used to determine corrosion current, and accordingly
corrosion rates, using a rotating disk electrode assembly. The role of RTIL was investigated
through scanning electron microscopy (SEM) and EDX analysis. The Oakton® pH meter
was used to monitor changes in pH of the aqueous amines during corrosion experiments.
Electrochemical setup. Electrochemical experiments were carried out using a setup with a
three-electrode configuration, i.e., platinum counter electrode, silver/silver chloride
(Ag/AgCl/sat. KCl) reference electrode, and a working electrode, as illustrated in Figure
3.1. Carbon steel 1020 having a chemical composition of 0.20% carbon, 0.50% manganese,
0.04% phosphorus, 0.05% sulfur, and balanced by iron was used as the specimen working
electrode to study the corrosive behavior of aqueous alkanolamines as well as that of
amine/RTIL solutions. A disk shaped working electrode having an exposure area of 0.196
cm2 was mounted in a Teflon cap. The experiments were performed in a 100 cm
3 volume
corrosion cell and each time 80 cm3 of test fluid with a specific composition was utilized.
Temperature was controlled using an oil bath and a condenser was engaged to minimize
evaporation of the experimental fluid.
Chapter 3
69
Figure 3. 1. Experimental setup for electrochemical corrosion tests.
In all the electrochemical measurements, aqueous amine or amine-RTIL solutions
incorporating 5 kmol/m3 amine were used.
Experimental procedure. Before each experiment, the working electrode surface was
polished by wet grinding with 600 grit SiC paper and alumina paste, respectively. The
specimen was then degreased with acetone followed by rinsing with deionized water. After
drying, the specimen was instantly immersed in the test solution in order to establish a
steady state open-circuit potential. To evaluate the effect of amine/RTIL type, solution
temperature, and water content, bubbling of a gaseous mixture comprising CO2 (15%), O2
(5%), and N2 (balance) was initiated one hour prior to the commencement of
electrochemical polarization run. In order to assess the effect of oxygen on corrosion of
steel; the O2 concentration was varied between 0%, 5%, and 10% in the simulated flue gas.
However, to estimate the influence of CO2 loading on the corrosion of steel, pure CO2 was
bubbled to attain the desired gas loading. After the accomplishment of the required
conditions, bubbling of the respective gaseous mixture of CO2+O2+N2 (bearing ≤3% CO2
to maintain gas loading during the execution of the electrochemical run) was sustained at a
flow rate of 100 cm3/min at ambient pressure. Calibrated gas flow meters were used for the
Chapter 3
70
purpose. Linear polarization resistance (LPR) curves were recorded at a constant scan rate
of 0.16 mVs-1
starting from a cathodic potential of -250 mV to an anodic potential of +250
mV (versus open circuit potential). During electrochemical experimentation, 500 rpm of
rotation speed was maintained for the working electrode. To ensure data reproducibility,
each experiment was replicated at least once and the stated corrosion rates are average
values with an uncertainty of ±5%.
The Tafel extrapolation method was applied to determine the corrosion current (icorr), which
was converted to the corrosion rate by the following equation:
Where, CR is the corrosion rate in milli-inch per year (mpy); icorr is the corrosion current in
Ampere; W is the equivalent weight of metal specimen in gram per equivalent; ρ is the
density of metal in g/cm3; and A is the area (in contact with experimental fluid) of the
rotating disk working electrode in cm2.
CO2 loading determination. For electrochemical corrosion tests, CO2 loading was
determined by a Chittick apparatus using the titration method [19]. Small samples were
withdrawn from the electrochemical cell and titrated against a standard solution of HCl
(1M) using methyl orange indicator to release the captured gas. CO2 loading was calculated
subsequently from the volume of CO2 evolved. In case of aqueous samples, the CO2 partial
pressure was corrected for the vapor pressure of water.
3.3. Results and Discussion
The main objective of current project is to find an amine-RTIL combination effective
enough in CO2 capture while averting the drawbacks (especially equilibrium limitations,
higher regeneration energy requirements and corrosion occurrence) of aqueous amine
systems.
5(1.29 10 ) corri WCR
A
Chapter 3
71
Contrary to aqueous alkanolamines, ionic liquids endow unique characteristic to the amine-
RTIL blends by making CO2-captured product (carbamate) to precipitate out as revealed in
Figure 3.2 (see also Figures B.1-B.4 in Appendix B) [16,17,20]. For this reason, under
ambient conditions, amine-RTIL mixtures enabled capture of CO2 up to stoichiometric
maximum avoiding any significant decelerating effect through equilibrium limitations
(Figure 3.3). The mere difference in capture kinetics seems only due to disparity in
viscosities of the RTILs (Table 3.2) as gas diffusivity decreases at higher viscosity.
Moreover, product precipitation, in case of primary/secondary alkanolamines, not only can
evade any active role of carbamate in electrochemical pathway regarding corrosion but
might also help facilitate easy removal of the product thus letting lesser volume to
regenerate.
Figure 3. 2. MEA-RTIL fluid showing solid carbamate, after CO2 bubbling: 1)
MEA+[bmim][BF4]; 2) MEA+[emim][BF4]; 3) MEA+[emim][Otf].
1 2 3
Chapter 3
72
Figure 3. 3. Thermogravimetric evolution of CO2 absorption for MEA-RTIL mixtures
(MEA: 5 kmol/m3) at 25 °C.
Table 3. 2. Viscosity values of the ionic liquids used.
Ionic liquid Viscosity (mPas)*
[bmim][BF4] 136.7
[emim][BF4] 34.0
[emim][Otf] 39.8 *at 25 °C
However, this particular chapter was specifically aimed at assessing thoroughly the role of
RTILs toward corrosion of steel in amine-RTIL mixtures. Replacing the aqueous phase
with RTIL seems a better practice to nullify corrosion. This was quite obvious from the
scanning electron microscopic study of working electrode surfaces (Figure 3.4) before and
after anodic polarization runs. In case of aqueous MEA (Figure 3.4a) the electrode surface
underwent deterioration due to corrosion, whereas in the case of RTIL based media the
surface morphology of the fresh and tested electrode (Figures 3.4b, 3.4c) appeared quite
similar revealing the protective function of RTIL.
Different process parameters were tested to evaluate the usefulness of RTILs against
corrosion of steel and the findings are being addressed in detail in the following sections.
Chapter 3
73
Figure 3. 4. SEM micrographs of steel electrode surface before and after electrochemical
polarization runs at 25 °C under CO2(15%)+O2(5%)+N2 atmosphere in: a) MEA (aqueous);
b) MEA+[bmim][BF4]; c) MEA+Water+[bmim][BF4].
3.3.1. Effect of amine type on corrosion of steel
The linear polarization curves (Figure 3.5a) and Tafel fit calculations (Table 3.3) for
aqueous amines demonstrated monoethanolamine to be the most corrosive among the tested
alkanolamines.
c)
b)
a) (before)
(before)
(before)
(after)
(after)
(after)
Chapter 3
74
Figure 3. 5. Linear polarization curves of carbon steel 1020 at 25 °C: a) in aqueous
alkanolamines; b) in alkanolamine+[bmim][BF4] mixtures.
-2
-1
0
1
2
3
-1100 -1000 -900 -800 -700 -600
log
I (
µA
)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
Aqueous amines
- AMP
- Blended amines
- DEA
- MEA
a)
-3
-2,5
-2
-1,5
-1
-0,5
0
0,5
1
-800 -700 -600 -500 -400 -300 -200 -100
log I
(µ
A)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
Amine+RTIL systems
- AMP
- Blended amines
- DEA
- MEA
b)
Chapter 3
75
Table 3. 3. Effect of amine type on corrosion parameters at 25 °C. Medium
(Amine: 5 kmol/m3)
CO2 loading
(mol CO2/mol amine)
Corrosion potential
(mV)
Corrosion current
(µA)
Corrosion rate
(mpy)
AMPaq 0.17 -800.0 4.2 9.8
Blended aminesaq* 0.09 -779.0 2.8 6.5
DEAaq 0.15 -726.8 2.3 5.4
MEAaq 0.23 -750.7 6.6 15.5
AMP+[bmim][BF4] 0.14 -363.1 0.15 0.35
Blended amines*#+[bmim][BF4] 0.06 -407.6 0.4 0.93
DEA+[bmim][BF4] 0.14 -422.3 0.16 0.37
MEA+[bmim][BF4] 0.24 -367.7 0.09 0.22
*MEA (0.5 kmol/m3) + MDEA (4.5 kmol/m
3);
#water (5 kmol/m
3)
As MEA is comparatively more reactive, during CO2 bubbling, it sharply converts to
carbamate/bicarbonate (RNHCOO-/HCO3
-) thus promptly increasing the concentration of
oxidants to take part in electrochemical corrosion reactions (equations 1-3) and favoring the
iron dissolution (reaction 4) [3,4].
RNHCOO- + H2O → RNH2 + HCO3
- (1)
2HCO3- + 2e
- → 2CO3
- + H2 (2)
2H2O+ 2e- → 2OH
- + H2 (3)
Fe → Fe2+
+ 2e- (4)
The cationic part (RR′NH2+) of carbamate might be another probable oxidizing agent
involved in the corrosion process as shown by equation 5 [4].
2RR′NH2+ + 2e
- → 2RR′NH + H2 (5)
Nevertheless, AMP and DEA caused corrosion problem to a lesser extent owing to their
low reactivity. Though, blended amines (in spite of low gas loading) induced relatively
higher corrosion, probably due to the formation of enough bicarbonate species. On the
other hand, in case of amine-RTIL systems (Figure 3.5b; Table 3.3) there was not much
effect of amine type as RTIL was the key component responsible for controlling the
corrosion phenomenon by forming a protective coating on the working electrode surface.
This mode was quite explicit from the positive shift of corrosion potential [21,22].
Furthermore, substituting RTIL for water diminished the active role of oxidizing species
involved in aqueous amine related corrosion reactions. Also, precipitation of gas-captured
product in case of amine-RTIL media involving MEA, AMP, and DEA made solid
Chapter 3
76
carbamate to move out of the reaction phase, thus avoiding its active participation in the
electrochemical corrosion process.
3.3.2. Effect of RTIL type on corrosion behaviour
[emim][BF4] and [bmim][BF4] alone or in combination with MEA caused minimal
corrosion at ambient conditions of temperature and pressure as can be seen from Figure 3.6;
Table 3.4.
Figure 3. 6. Effect of RTIL type on polarization behavior of carbon steel 1020 at 25 °C.
Table 3. 4. Effect of RTIL type on corrosion rate of carbon steel 1020 at 25 °C.
Medium
(Amine: 5 kmol/m3)
Corrosion potential
(mV)
Corrosion current
(µA)
Corrosion rate
(mpy)
Pure [emim][Otf] -159.3 5.5 12.8
Pure [emim][BF4] -305.2 0.27 0.62
Pure [bmim][BF4] -251.7 0.05 0.13
MEA+[emim][Otf] -525.9 0.27 0.63
MEA+[emim][BF4] -423.1 0.26 0.61
MEA+[bmim][BF4] -367.7 0.09 0.22
This consequence was due to the formation of a protective layer by RTIL on the metal
surface that blocked the active sites against approaching oxidizing species. EDX analysis of
the working electrode surface as well as noble shift of the corrosion potential confirmed
such developments. EDX results demonstrated the presence of the carbon peak as well as
the overlapping strong peak of fluorine (with that of iron) as shown in Figures 3.7b and
-3
-2
-1
0
1
2
3
-800 -700 -600 -500 -400 -300 -200 -100 0 100
log
I (
µA
)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
RTIL type
Chapter 3
77
3.7c indicating the existence of surface protective films of [bmim][BF4] and [emim][BF4]
respectively. This behavior is also evident from a decrease in iron peak intensities
compared to those in the EDX spectrum of the fresh surface (Figure 3.7a).
On the contrary, [emim][Otf] in pure form proved most corrosive among the three ionic
liquids tested. The presence of acidic impurities in [emim][Otf] appeared to be the main
cause of this behavior [23]. The pH (2.63) of the aqueous solution of [emim][Otf] (10%
w/w) also validated its acidic aspect, whereas [emim][BF4] and [bmim][ BF4], being of
neutral nature, showed excellent corrosion control potential. However, in case of amine-
RTIL systems, [emim][Otf] demonstrated similar behavior as that by [emim][BF4],
significantly diminishing corrosion. This is because the acidic impurities were
counterbalanced by the dominating effect of MEA present in excess, thus enabling the
RTIL to safeguard the metal surface effectively (Figure 3.7d).
Energy (keV)
Co
un
ts
(a)
Chapter 3
78
Energy (keV)
Co
un
ts
(b)
Energy (keV)
Co
un
ts
(c)
Chapter 3
79
Figure 3. 7. EDX analysis of steel electrode surface: a) freshly polished surface; b,c,d)
after electrochemical corrosion tests in MEA+[bmim][BF4], MEA+[emim][BF4],
MEA+[emim][Otf] blends respectively.
3.3.3. Effect of process temperature
To know the comparative effect of temperature on corrosion of steel in aqueous and RTIL
based media, LPR experiments were conducted at 25 °C and 60 °C. In the aqueous amine
system, the rise in temperature resulted in nearly doubling of the corrosion rate from 15.5
mpy (at 25 °C) to 26.4 mpy (at 60 °C), see Figure 3.8, Table 3.5. On the other hand in
amine-RTIL mixtures, at 25 °C, the corrosion phenomenon was suppressed to almost
negligible which seems to be not only because of the formation of RTIL protective layer on
the working electrode surface but also due to higher viscosity of the fluid that makes the
diffusion of electrochemical species (involved in corrosion phenomenon) between the
electrodes strenuous. However, at 60 °C, the likely depletion of the protective film as well
as decrease in viscosity led to a higher corrosion rate (7.4 mpy). The EDX surface
examination revealed weakening of the protective layer at 60 °C as is evident from lower
carbon/fluorine peak intensities (Figure 3.9) compared to the EDX spectrum (Figure 3.7b)
of the electrode surface used during a corrosion test at 25 °C in MEA-[bmim][BF4] media
having equivalent composition. Even so, the outcome confirms the positive role of
Energy (keV)
Co
un
ts(d)
Chapter 3
80
[bmim][BF4] in decreasing the corrosion rate up to 72% compared to that in aqueous
monoethanolamine solvent under the same process conditions.
Figure 3. 8. Comparison of temperature effect on steel corrosion in aqueous as well as
RTIL based media.
Figure 3. 9. EDX scan of steel electrode surface after electrochemical corrosion test in
MEA+[bmim][BF4] at 60 °C.
-3
-2
-1
0
1
2
3
-1100 -1000 -900 -800 -700 -600 -500 -400 -300 -200 -100
log I
(µ
A)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
Temperature Effect
- MEAaq (25 C)
- MEAaq (60 C)
- MEA+RTIL (25 C)
- MEA+RTIL (60 C)
Co
un
ts
Energy (keV)
Chapter 3
81
Table 3. 5. Effect of process temperature on corrosion rate of carbon steel 1020.
Medium
(Amine: 5 kmol/m3)
Process temperature
(°C)
Corrosion potential
(mV)
Corrosion current
(µA)
Corrosion rate
(mpy)
MEAaq 25 -750.7 6.6 15.5
MEAaq 60 -796.2 11.3 26.4
MEA+[bmim][BF4] 25 -367.7 0.09 0.22
MEA+[bmim][BF4] 60 -837.2 3.2 7.4
3.3.4. Effect of gas loading
In aqueous MEA, the corrosion rate intensified as the loading was increased from 0.23 to
0.53 mole of CO2 per mole of amine. This effect can be attributed to a higher concentration
of RNHCOO-/HCO3
- species, the key oxidizing agents involved in electrochemical
corrosion reactions. A decrease in pH of aqueous media with increasing CO2 loading was
another cause that also helped accelerate iron dissolution.
However in MEA-RTIL media, higher CO2 loading appears to be better to nullify the
corrosion phenomenon as is evident from Figure 3.10 and Table 3.6. The stoichiometric
maximum CO2 loading avoids any participation of amine species in corrosion phenomenon
by eliminating the entire amine out of the liquid phase as solid carbamate, and this behavior
is fully explicit from diminution of the corrosion rate to 0.08 mpy.
-1.5
-0.5
0.5
1.5
2.5
-1100 -1000 -900 -800 -700 -600 -500 -400
log
I (
µA
)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
MEAaq: CO2 loading
- 0.23- 0.42- 0.53
a)
Chapter 3
82
Figure 3. 10. CO2 loading effect on steel corrosion at 25 °C in a) aqueous MEA b)
MEA+[bmim][BF4] mixture.
Table 3. 6. Corrosion rate of steel in aqueous MEA and MEA+[bmim][BF4] blends at
different CO2 loadings and 25 °C.
Medium (Amine: 5 kmol/m3)
CO2 loading (mol CO2/mol amine)
pH Corrosion potential (mV)
Corrosion current (µA)
Corrosion rate (mpy)
MEAaq 0.23 10.55 -750.8 6.6 15.5
MEAaq 0.42 9.40 -790.0 19.0 37.2
MEAaq 0.53 8.29 -678.5 40.7 95.1
MEA+[bmim][BF4] 0.24 - -367.7 0.09 0.22
MEA+[bmim][BF4] 0.35 - -418.3 0.12 0.29
MEA+[bmim][BF4] 0.50 - -636.5 0.034 0.08
3.3.5. Presence of oxygen
The effect of oxygen on steel corrosion was studied by varying the concentration of oxygen
from 0 to 10% in the simulated flue gas. Oxygen had a significant effect in case of aqueous
MEA as the corrosion rate augmented from 3.04 mpy (in the absence of oxygen) to 15.5
mpy (in the presence of oxygen) as shown in Figure 3.11, Table 3.7. The presence of
oxygen increases the number of oxidizing species that along with water acts as a sink for
electrons oxidized from iron.
O2 + 2H2O + 4e- → 4OH
-
-3
-2,5
-2
-1,5
-1
-0,5
0
0,5
1
-900 -800 -700 -600 -500 -400 -300 -200 -100
log I
(µ
A)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
MEA+RTIL: CO2 loading
- 0.24- 0.35- 0.50
b)
Chapter 3
83
Conversely in amine-RTIL solutions, the protective coating of ionic liquid as well as
absence of water did not permit oxygen to play any active role in the corrosion
phenomenon.
Figure 3. 11. Effect of O2 concentration in flue gas on corrosion of steel a) in aqueous
MEA; b) in MEA+[bmim][BF4] mixture.
-3
-2
-1
0
1
2
3
-1100 -1000 -900 -800 -700 -600
log
I (
µA
)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
MEAaq: O2 ratio in flue gas
- O2 0%- O2 5%- O2 10%
a)
-3
-2,5
-2
-1,5
-1
-0,5
0
0,5
1
1,5
2
-1100 -1000 -900 -800 -700 -600 -500 -400 -300 -200 -100
log I
(µ
A)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
MEA+RTIL: O2 ratio in flue gas
b)
- O2 0%- O2 5%- O2 10%
Chapter 3
84
Table 3. 7. Effect of oxygen presence/absence on corrosion rate of carbon steel 1020 at 25
°C.
Medium
(Amine: 5 kmol/m3)
O2 content in flue gas
(Vol. %)
Corrosion potential
(mV)
Corrosion current
(µA)
Corrosion rate
(mpy)
MEAaq 0 -795.5 1.30 3.04
MEAaq 5 -750.7 6.62 15.5
MEAaq 10 -717.7 6.52 15.2
MEA+[bmim][BF4] 0 -819.7 0.20 0.46
MEA+[bmim][BF4] 5 -367.7 0.09 0.22
MEA+[bmim][BF4] 10 -376.8 0.15 0.36
3.3.6. Influence of water
To examine the effect of water content, 5 kmol/m3 of water was mixed in [bmim][BF4] and
MEA-[bmim][BF4] systems, respectively. The water-RTIL mixture wreaked corrosion of
steel analogous to that by aqueous amine solvent (Figure 3.12; Table 3.8). This behavior
reveals that the presence of water prompted formation of HCO3- species and also weakened
the RTIL protective coating thus making the electrode surface vulnerable to corrosion.
However in the water-MEA-RTIL system, the corrosion rate fell to 0.37 mpy. In the
presence of MEA, carbamate formation and subsequent precipitation seem to remove water
hygroscopically out of the RTIL phase, thus crafting the RTIL protective layer strong
enough to hinder corrosive action on the steel electrode surface.
Figure 3. 12. Effect of water content in CO2 capture medium on corrosion of steel.
-3,5
-2,5
-1,5
-0,5
0,5
1,5
2,5
3,5
4,5
-600 -500 -400 -300 -200 -100 0
log I
(µ
A)
Ew (mV) vs. Ag/AgCl/KCl (sat'd)
Water presence
- RTIL
- Water+RTIL
- MEA+RTIL
- Water+MEA+RTIL
Chapter 3
85
Table 3. 8. Influence of water content in the gas capture fluid on corrosion of steel at 25 °C.
Medium
(Amine: 5 kmol/m3)
Corrosion
potential
(mV)
Corrosion
current
(µA)
Corrosion
rate
(mpy)
Pure [bmim][BF4] -251.7 0.05 0.13
Water+[bmim][BF4] -215.4 1.55 3.62
MEA+[bmim][BF4] -367.7 0.09 0.22
Water+MEA+[bmim][BF4] -353.1 0.16 0.37
3.4. Conclusion
To cope with anthropogenic greenhouse gas emissions, an extremely efficient system is
required that can help detain the flow of CO2 into the atmosphere at reasonable costs. This
can be done by inducing stability to the current chemical capture solvents, and amine-RTIL
blends may provide a feasible opportunity. To test their suitability, the corrosion rate of
carbon steel 1020 in alkanolamine-RTIL solutions was studied and the results were
compared with aqueous alkanolamine systems. From the experimental data it can be
concluded:
[emim][BF4] and [bmim][BF4] ionic liquids not only exhibited excellent corrosion
control in the pure state but also demonstrated a similar trend in combination with
alkanolamines. Whereas pure [emim][Otf] was quite corrosive toward the electrode
surface primarily due to the presence of acidic impurities. However, the presence of
amine nullified the effect of acid content and consequently the MEA-[emim][Otf]
blend demonstrated good efficacy.
EDX analysis revealed the formation of a protective film that helped shield the
metal surface against the detrimental effect of the fluid contents.
In contrast to aqueous amine systems, a higher gas loading further improved the
corrosion prevention ability of amine-RTIL blends. Moreover, the presence/absence
of water and oxygen did not insert any negative impact on the shielding skill of
amine-RTIL mixtures against corrosion.
Chapter 3
86
At 60 °C, MEA-[bmim][BF4] does not seem as effective as at room temperature.
Even so, there is around 72% decrease in the corrosion rate than that in aqueous
amine system at the same temperature.
These results are quite encouraging, revealing that these new schemes may prove beneficial
in pre-/post-combustion CO2 capture. However, to fully exploit the advantages of amine-
RTIL blends, further screening is needed specifically to test their gas capturing and
regeneration ability.
3.5. References
[1] B. Metz, O. Davidson, H. de Coninck, M. Loos, L. Meyer, Eds., IPCC Special Report
on Carbon Dioxide Capture and Storage, Prepared by Working Group III of the
Intergovernmental Panel on Climate Change, Cambridge University Press: New York,
2005.
[2] A.L. Kohl, R.B. Nielsen, Gas Purification, 5th ed.; Gulf Publishing Company: Houston,
Texas, 1997.
[3] I.R. Soosaiprakasam, A. Veawab, Corrosion and polarization behavior of carbon steel in
MEA-based CO2 capture process, Int. J. Greenh. Gas Control 2 (2008) 553-562.
[4] N. Kladkaew, R. Idem, P. Tontiwachwuthikul, C. Saiwan, Corrosion Behavior of
Carbon Steel in the Monoethanolamine-H2O-CO2-O2-SO2 System: Products, Reaction
Pathways, and Kinetics, Ind. Eng. Chem. Res. 48 (2009) 10169-10179.
[5] A. Veawab, P. Tontiwachwuthikul, S.D. Bhole, Studies of Corrosion and Corrosion
Control in a CO2-2-Amino-2-methyl-1-propanol (AMP) Environment, Ind. Eng. Chem.
Res. 36 (1997) 264-269.
[6] M. Nainar, A. Veawab, Corrosion in CO2 capture unit using MEA-piperazine blends,
Energy Procedia 1 (2009) 231-235.
[7] A. Veawab, P. Tontiwachwuthikul, A. Chakma, Investigation of Low-Toxic Organic
Corrosion Inhibitors for CO2 Separation Process Using Aqueous MEA Solvent, Ind. Eng.
Chem. Res. 40 (2001) 4771-4777.
Chapter 3
87
[8] J.F. Brennecke, E.J. Maginn, Ionic Liquids: Innovative Fluids for Chemical Processing,
AIChE J. 47 (2001) 2384-2389.
[9] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Ionic Liquids for CO2 Capture -
Development and Progress, Chem. Eng. Process. 49 (2010) 313-322.
[10] D. Camper, P. Scovazzo, C. Koval, R. Noble, Gas Solubilities in Room-Temperature
Ionic Liquids, Ind. Eng. Chem. Res. 43 (2004) 3049-3054.
[11] A. Yokozeki, M.B. Shiflett, Separation of Carbon Dioxide and Sulfur Dioxide Gases
Using Room-Temperature Ionic Liquid [hmim][Tf2N], Energy Fuels 23 (2009) 4701-4708.
[12] A.P-S. Kamps, D. Tuma, J. Xia, G. Maurer, Solubility of CO2 in the Ionic Liquid
[bmim][PF6], J. Chem. Eng. Data 48 (2003) 746-749.
[13] C. Cadena, J.L. Anthony, J.K. Shah, T.I. Morrow, J.F. Brennecke, E.J. Maginn, Why
Is CO2 So Soluble in Imidazolium-Based Ionic Liquids? J. Am. Chem. Soc. 126 (2004)
5300-5308.
[14] R.E. Baltus, B.H. Culbertson, S. Dai, H. Luo, D.W. DePaoli, Low-Pressure Solubility
of Carbon Dioxide in Room-Temperature Ionic Liquids Measured with a Quartz Crystal
Microbalance, J. Phys. Chem. B 108 (2004) 721-727.
[15] A.M. Schilderman, S. Raeissi, C.J. Peters, Solubility of carbon dioxide in the ionic
liquid 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide, Fluid Phase
Equilibr. 260 (2007) 19-22.
[16] D. Camper, J.E. Bara, D.L. Gin, R.D. Noble, Room-Temperature Ionic Liquid-Amine
Solutions: Tunable Solvents for Efficient and Reversible Capture of CO2, Ind. Eng. Chem.
Res. 47 (2008) 8496-8498.
[17] Q. Huang, Y. Li, X. Jin, D. Zhao, G.Z. Chen, Chloride ion enhanced thermal stability
of carbon dioxide captured by monoethanolamine in hydroxyl imidazolium based ionic
liquids, Energy Environ. Sci. 4 (2011) 2125-2133.
[18] M.F. Arenas, R.G. Reddy, Corrosion of Steel in Ionic Liquids, J. Min. Metall. Sect. B-
Metall. 39 (2003) 81-91.
Chapter 3
88
[19] A. Dreimanis, Quantitative Gasometric Determination of Calcite and Dolomite by
Using Chittick Apparatus, J. Sediment. Petrol. 32 (1962) 520-529.
[20] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 Capture in Alkanolamine/Room-
Temperature Ionic Liquid Emulsions: A Viable Approach with Carbamate Crystallization
and Curbed Corrosion Behavior, Int. J. Greenhouse Gas Control 6 (2012) 246-252.
[21] D. Guzmán-Lucero, O. Olivares-Xometl, R. Martínez-Palou, N.V. Likhanova, M.A.
Domínguez-Aguilar, V. Garibay-Febles, Synthesis of Selected Vinylimidazolium Ionic
Liquids and Their Effectiveness as Corrosion Inhibitors for Carbon Steel in Aqueous
Sulfuric Acid, Ind. Eng. Chem. Res. 50 (2011) 7129-7140.
[22] A.K. Satapathy, G. Gunasekaran, S.C. Sahoo, K. Amit, P.V. Rodrigues, Corrosion
inhibition by Justicia gendarussa plant extract in hydrochloric acid solution, Corrosion Sci.
51 (2009) 2848-2856.
[23] Green Solvents - Impurities & Corrosion, 2006, IoLiTec Inc.,
(http://www.iolitec.de/en/Poster/).
Chapter 4
89
CO2 capture in alkanolamine-RTIL blends via carbamate crystallization:
route to efficient regeneration*
Abstract/Résumé
One of the major drawbacks of aqueous alkanolamine based CO2 capture processes is the
requirement of significantly higher energy of regeneration. This weakness can be overcome
by separating the CO2-captured product to regenerate the corresponding amine, thus
avoiding the consumption of redundant energy. Replacing aqueous phase with more stable
and practically nonvolatile imidazolium based room-temperature ionic liquid (RTIL)
provided a viable approach for carbamate to crystallize out as a supernatant solid. In the
present study, regeneration capabilities of solid carbamates have been investigated.
Diethanolamine (DEA) carbamate as well as 2-amino-2-methyl-1-propanol (AMP)
carbamate was obtained in crystalline form by bubbling CO2 in alkanolamine-RTIL
mixtures. Hydrophobic RTIL, 1-hexyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide ([hmim][Tf2N]), was used as aqueous phase substituent.
Thermal behavior of the carbamates was observed by differential scanning calorimetry and
thermogravimetric analysis, while the possible regeneration mechanism has been proposed
through 13
C NMR and FTIR analyses. The results showed that decomposition of DEA-
carbamate commenced at lower temperature (~55 °C), compared to that of AMP-carbamate
(~75 °C); thus promising easy regeneration. The separation of carbamate as solid phase can
offer two-way advantage by letting less volume to regenerate as well as by narrowing the
gap between CO2 capture and amine regeneration temperatures.
L'un des inconvénients majeurs des procédés de capture du CO2 basé sur l‟utilisation
d‟alcanolamines aqueuses est la grande quantité d‟énergie requise lors de la régénération.
Cette faiblesse peut être surmontée en séparant le produit issu de la capture du CO2 afin de
régénérer l'amine correspondante en évitant ainsi une consommation d'énergie
supplémentaire. La substitution de la phase aqueuse par une phase plus stable, pratiquement
non-volatile, comme un liquide ionique à température ambiante (RTIL) à base
* M. Hasib-ur-Rahman, F. Larachi, Environ. Sci. Technol. 46 (2012) 11443-11450.
Chapter 4
90
d‟imidazolium offre une approche viable pour la cristallisation du carbamate en un
surnageant solide. Dans cette étude, les possibilités de régénération des solides carbamates
ont été étudiées. Le carbamate de diéthanolamine (DEA) ainsi que le carbamate du 2-
amino-2-methyl-1-propanol (AMP) ont été cristallisés en barbotant du CO2 dans un
mélange alcanolamine-(RTIL). Le RTIL hydrophobe 1-hexyl-3-methylimidazolium bis-
(trifluoromethylsulfonyl) imide ([hmim][Tf2N]) a été utilisé en tant que substituant de la
phase aqueuse. Le comportement thermique des carbamates a été observé par calorimétrie
différentielle à balayage et par analyse thermogravimétrique, tandis qu‟un mécanisme
possible de régénération a été proposé grâce aux analyses au carbone 13 et par
spectroscopie infrarouge à transformée de Fourier. Les résultats ont montré que la
décomposition du carbamate de DEA débute à plus basse température (~ 55 °C)
comparativement à celle de du carbamate d‟AMP (~ 75 °C); ceci promettant une
régénération aisée. La séparation de carbamate comme phase solide peut offrir un avantage
double en laissant d‟une part moins de volume à régénérer et d‟autre part, la réduction de
l'écart entre les températures de capture du CO2 et de régénération des amines.
4.1. Introduction
Anthropogenic industrial activities are causing serious increase in atmospheric
concentration of greenhouse gases; and carbon dioxide, being the most important of these
in perspective of its contributions toward global warming, is considered as the main cause
of environmental problems in this regard [1-4]. Major CO2 emission sources that offer
potential capture convenience comprise fossil-fuel based power generation installations [5].
Various measures are being explored to check CO2 emissions from large point sources into
the atmosphere. These include physical/chemical sorption, membrane separation, and
cryogenic distillation techniques. In industry, the most preferred gas absorption processes
comprise alkanolamine based aqueous solvents executing absorber-stripper arrangements,
and can principally be used for postcombustion CO2 capture [5-7]. At temperatures around
40 °C aqueous solutions of primary and secondary amines, such as monoethanolamine
(MEA), diethanolamine (DEA) respectively, are subjected to absorb CO2 through
carbamate formation whereas tertiary amines, such as N-methyldiethanolamine (MDEA),
along with water react with the sour gas to form ammonium bicarbonate. In case of
Chapter 4
91
primary/secondary amines, predominantly one mole of CO2 reacts with two moles of amine
obeying the following mechanism (eqs 1 and 2) [8,9]:
2
2
CO RR NH RR NH COO
RR NH COO RR NH RR NH RR NCOO
However, in presence of water, tertiary amines react with CO2 in 1:1 molar ratio, as shown
below (eqs 3 and 4):
2 2 3CO H O H HCO
RR R N H RR R NH
Then the regeneration of these solvents is carried out by heat stripping at temperatures in
the range of 100 °C to 140 °C [5]. In case of primary/secondary aqueous alkanolamines, the
following regeneration mechanism (eqs 5 and 6) has been proposed [8,10]:
2 2
2 2
RR NCOO H O CO RR NH OH
RR NH OH RR NH H O
While regeneration of tertiary amines occurs as follows (eqs 7 and 8):
3 2
2
HCO CO OH
RR R NH OH RR R N H O
Nevertheless, there are many downsides of these CO2 capture systems like low gas loading,
degradation/evaporation of amines, and corrosion of equipment [11-13]. Higher
regeneration energy requirement is one of the major drawbacks of aqueous alkanolamine
based state-of-the-art technologies. In a power generation plant, up to 40% additional
energy is required for carbon dioxide capture and storage (CCS). Out of this extra bite,
roughly 50% is consumed in regeneration step alone [5].
Recently, unique room-temperature ionic liquids (RTILs), owing to their tunable
physicochemical characteristics, have been emerging as potential contenders for CO2
capture [6,14]. In this context, thermally stable imidazolium based RTILs are being
(1)
(2)
(3)
(4)
(5)
(6)
(7)
(8)
Chapter 4
92
investigated extensively as prospective alternates [15-19]. Pressure swing technique can be
used to regenerate such solvents. However, like other physical solvents such as methanol,
dimethyl ethers of polyethylene glycol (currently being used industrially as rectisol/selexol
processes), these alone cannot be employed effectively for separating CO2 from flue gases
with low CO2 partial pressures [20,21]. Neither aqueous alkanolamines nor RTILs solely
are proficient enough for economical CO2 separation.
In search of an efficient CO2 separation process, various methodologies are being
scrutinized. These include amino functionalized solid adsorbants, task specific ionic
liquids, as well as supported ionic liquid membranes [14,22]. Work has also been initiated
to combine the advantages of RTILs with those of primary/secondary alkanolamines, and in
this regard Camper et al. were the first to report MEA-carbamate precipitation in amine-
RTIL solution [23-26]. In case of alkanolamine solvents, replacing aqueous phase with
more stable room-temperature ionic liquid (RTIL) can avoid the corrosion and equilibrium
limitation problems particularly arising due to the presence of water. More significantly,
the presence of RTIL provides the favorable environment for CO2-captured product to
crystallize out, thus making it possible to easily separate the solid carbamate from the liquid
counterpart in addition to completing the reaction to its full stoichiometric potential. As
CO2 is about 3 times more soluble (in terms of moles of CO2 per volume of the solvent) in
imidazolium based RTILs than in water [17,27,28], this new approach of CO2 absorption in
alkanolamine-RTIL mixtures can ensure greater mass transfer capacity thus compensating
to a certain extent the downside posed by higher viscosity of the ionic liquids.
The objectives of this study were to look for an apposite alkanolamine-hydrophobic RTIL
combination that can (a) guarantee stoichiometric maximum CO2 loading by evading
equilibrium constraints; (b) minimize stripping temperatures; (c) manage less volumes to
regenerate through separation of CO2-captured product thus letting ensue probable cut
down of the gratuitously high regeneration energy to affordable limit. The overall concept
has been envisaged in Figure 4.1.
Chapter 4
93
Figure 4. 1. The simplified process flow diagram of alkanolamine-RTIL based CO2 capture
process.
The current activity was focused on looking into the regeneration scenario of CO2
absorption process comprising AMP/DEA-RTIL blends. Single crystal X-ray diffraction
technique and 13
C NMR/FTIR analyses were employed to infer the nature of CO2-captured
products and the regenerated amines. Whereas decomposition behavior of solid carbamates,
obtained by bubbling CO2 through amine-RTIL blends containing either 2-amino-2-methyl-
1-propanol (AMP) or diethanolamine (DEA), has been investigated in detail using
differential scanning calorimetry (DSC), thermogravimetry (TG), 13
C NMR, and FTIR
techniques.
4.2. Experimental
4.2.1. Materials
2-Amino-2-methyl-1-propanol (AMP: purum, ≥97.0%) and Diethanolamine (DEA: ACS
reagent, ≥99.0%) were purchased from Sigma-Aldrich, and Triton® X-100 (t-
Octylphenoxypolyethoxyethanol, a nonionic surfactant) was obtained from EMD
Chapter 4
94
Chemicals. IoLiTec Inc. supplied RTIL, 1-hexyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide ([hmim][Tf2N]: 99% purity). While carbon dioxide and
nitrogen gases (≥99% purity) were obtained from Praxair Canada Incorporation. All the
materials were used as received.
4.2.2. Procedures and techniques
4.2.2.1.CO2 capture studies
Gas absorption studies were carried out by thermogravimetric analyzer (Perkin-Elmer
Diamond TG/DTA) under carbon dioxide atmosphere isothermally at 35 °C. For this
purpose, 18 (±1) mg sample (amine-RTIL mixture) was loaded in an aluminum pan and
placed in the analyzer under N2 atmosphere. Then the sample was exposed to pure CO2 to
obtain CO2 uptake profile. Mass flow meters were used to adjust gas flow rates at 100
mL/min.
Prior to gas absorption capacity measurements by thermogravimetric analyzer,
alkanolamine-RTIL samples were prepared using Omni homogenizer (Omni International,
Kennesaw, GA) fitted with rotor-stator generator. Fifteen wt% amine (AMP/DEA) was
mixed in [hmim][Tf2N]. Though, in case of DEA/[hmim][Tf2N] blend, Triton X-100
surfactant was added to stabilize the homogeneity of the mixture.
In order to get solid carbamates, CO2 was bubbled through 15 wt% amine-RTIL blends
(without surfactant) at 35 °C along with continuous stirring for two hours. The suspension
obtained as a result of carbamate crystallization was allowed to stand for 48 hours to help
the two phases settle apart. This enabled easy separation of supernatant crystals that were
washed thoroughly with acetone, dried and stored at room temperature.
4.2.2.2.Carbamate characterization
To know the nature of the CO2-captured products (AMP-carbamate, DEA-carbamate), 13
C
NMR spectra were recorded on a Varian Inova Spectrometer (Palo Alto, CA) at a
frequency of 100 MHz with proton decoupling, after dissolving the crystals in DMSO-d6
solvent (CND Isotopes, QC, Canada). Whereas a Nicolet Magna 850 spectrometer (Thermo
Scientific, Madison, WI) equipped with high temperature Golden Gate ATR accessory was
Chapter 4
95
used to perform FTIR analysis, and Single crystal X-ray diffraction technique provided the
detailed information about crystalline structures.
4.2.2.3.Regeneration behavior
Amine regeneration studies were carried out using thermogravimetric (TG) analyzer and
differential scanning calorimetry (DSC). In case of TG analysis, 9 (±1) mg of ground
carbamate sample was taken in an aluminum pan and the analysis was conducted using a
heating rate of 5 °C per minute. The regeneration behavior of carbamates was studied under
two different environments, that is, pure N2, and pure CO2. The onset temperature for
carbamate decomposition under N2 atmosphere, at which gas evolution started, was
detected by quadrupole mass spectrometer (Thermostar Prisma QMS200, Pfeiffer Vacuum
GmbH, Asslar, Germany) coupled with thermogravimetric analyzer. The gas flow rate was
maintained at 100 mL/min. To ensure the reproducibility, each experiment was repeated at
least once. Differential scanning calorimetric analyses were performed using a Mettler-
Toledo DSC1 (Columbus, OH) instrument. DSC scans were also managed at a temperature
scan rate of 5 °C per minute. 13
C NMR and ATR-FTIR techniques were employed to
confirm the likely regeneration mechanism.
4.3. Results and Discussion
4.3.1. Maximum gas capture capacity
CO2 absorption in AMP-RTIL and DEA-RTIL blends resulted in crystallization of the
product. This development enabled the product (carbamate) to move out of the reaction
phase and hence helped overcome the equilibrium limitation barrier thus not only allowing
maximum CO2 loading but also enabling easy separation of the solid product [25].
However, due to higher volatility of AMP [29,30], regarding AMP-RTIL combination, it
was not possible to maintain the initial concentration of amine in AMP-RTIL blends. And
so the CO2 capture capacity apparently appeared inferior to what the theoretical maximum
would have been with respect to initial AMP concentration (Figure 4.2). The evaporation
phenomenon was quite evident from the mass loss profile of AMP-RTIL blend acquired
under N2 atmosphere at 35 °C (Figure 4.3).
In order to verify the CO2 capture capacity in case of AMP-RTIL blend, the resulting AMP
carbamate was titrated against 1M HCl to release captured gas, using Chittick apparatus.
Chapter 4
96
This practice substantiated the 50 mol % absorption limit of CO2 (w.r.t. AMP) in AMP-
RTIL blend. The procedure has been described in the previous work [26].
Figure 4. 2. CO2 absorption isotherm for alkanolamine-[hmim][Tf2N] systems obtained at
atmospheric pressure and 35 °C temperature.
Figure 4. 3. Evaporation profiles of amines (in amine-RTIL blends) at 35 °C under N2.
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150 200 250 300
Mo
le r
ati
o C
O2/A
min
e
Time (min)
— DEA (15 wt%)/RTIL
— AMP (15 wt%)/RTIL
0
20
40
60
80
100
0 50 100 150 200 250 300
Am
ine
(w
eig
ht
%)
Time (min)
— DEA (15 wt%)/RTIL
— AMP (15 wt%)/RTIL
Chapter 4
97
However, no detectable evaporation loss was observed in case of emulsified DEA-RTIL
mixture under the specified conditions, and CO2 capture resulted in theoretical maximum
mass uptake (0.5 mole of CO2 per mole of DEA, in accordance with the mechanism
proposed by Caplow [8]).
CO2 capture studies at ambient conditions using DEA/[hmim][Tf2N] emulsion has been
discussed in our previous study [25].
4.3.2. Nature of CO2-captured products
Single crystal structure determination confirmed the formation of carbamate product,
originating from chemical interaction of CO2 with amine; both (AMP-carbamate and DEA-
carbamate) possessing monoclinic crystal system with P21/n and Pn space groups
respectively (Figure 4.4; see also supporting information, Appendix C). Appearance of
additional 13
C NMR signals at 162.59 ppm and 162.57 ppm, regarding corresponding CO2-
captured products (AMP-carbamate and DEA-carbamate respectively), also validated the
CO2 absorption exclusively through carbamate formation. These outcomes were further
complemented by FTIR analysis (Figures 4.5, 4.6).
Figure 4. 4. Packing diagrams: a) AMP-carbamate; b) DEA-carbamate [25].
(a) (b)
Chapter 4
98
As is observed in case of aqueous AMP based CO2 separation processes, AMP being a
sterically hindered amine favors CO2 absorption via bicarbonate formation owing to water
involvement that can guarantee higher sorption capacity. On the other hand, in present
work, absence of water prohibited the formation of bicarbonate species, limiting the gas
capture capacity to 50 mol % of CO2. Thermogravimetric isotherms as well as Chittick
apparatus measurements also confirmed the same outcome as CO2 capture capacity never
exceeded 0.5 CO2/amine molar ratio. DEA interacts with CO2 preferably through zwitterion
mechanism yielding carbamate product in either case, regarding aqueous DEA or DEA-
RTIL blends.
The detailed description of crystal structure determination of AMP-carbamate is provided
in Appendix C, whereas single crystal X-ray diffraction study of DEA-carbamate has been
discussed in the previous work [25].
4.3.3. Regeneration ability
Regeneration was brought about by thermal decomposition of carbamates at 110 °C that
resulted in quick release of CO2 and corresponding alkanolamine (AMP/DEA). 13
C NMR
as well as ATR-FTIR analyses of fresh and regenerated amines demonstrated the excellent
regeneration ability of both AMP and DEA. Theoretically, the probable mechanism might
comprise the following reactions (eqs 9 and 10) responsible for CO2 liberation during heat
treatment.
The FTIR as well as 13
C NMR spectra of fresh/regenerated amines and relevant carbamates
are shown in Figures 4.5 and 4.6. The emergence of respective carbon signals in 13
C NMR
spectra at 162.59 ppm and 162.57 ppm (Figures 4.5b, 4.6b) confirmed the CO2 absorption
via AMP-carbamate and DEA-carbamate formation. Two series of carbon signals
(compared to one series for corresponding fresh amine) in the range of 20-80 ppm, one
originating from protonated amine and the other from carbamate moiety, also
complemented the findings. Besides, the identical nature of NMR spectra of fresh and
2
2 2
RR NCOO RR N CO
RR N RR NH RR NH
(9)
(10)
Chapter 4
99
regenerated amines ruled out any probability of degradation occurrence at least after single
absorption/desorption cycle. FTIR analysis (Figures 4.5a, 4.6a) too revealed the same
outcome.
Chapter 4
100
Figure 4. 5. a) FTIR spectra, and b) 13
C NMR spectra: AMP (fresh amine), AMPC (AMP-
carbamate) and RAMP (regenerated AMP).
C
CH2CH3
NHCOO-
CH3
OHC
CH2 CH3
NH3
+CH3
OH
AMP Carbamate
C
CH2CH3
NH2CH3
OH
AMP
C
CH2CH3
NH2CH3
OH
Regenerated AMP
(a) (b)
AMP
C
CH2CH3
NH2CH3
OH
AMP
AMPC
C
CH2CH3
NHCOO-
CH3
OHC
CH2 CH3
NH3
+CH3
OH
AMP Carbamate
RAMP
C
CH2CH3
NH2CH3
OH
Regenerated AMP
Chapter 4
101
Figure 4. 6. a) FTIR spectra, and b) 13
C NMR spectra of DEA (fresh amine), DEAC
(DEA-carbamate) and RDEA (regenerated DEA).
NOH OH
OO-
NH2
+OH OH
DEA Carbamate
NHOH OH
DEA
NHOH OH
Regenerated DEA
(a) (b)
DEA
NHOH OH
DEA
DEAC
NOH OH
OO-
NH2
+OH OH
DEA Carbamate
RDEA
NHOH OH
Regenerated DEA
Chapter 4
102
4.3.4. Amine (AMP/DEA) regeneration behavior
Under N2 atmosphere, decomposition of AMP-carbamate commenced around 75 °C with
CO2 liberation, accompanied by simultaneous evaporation of amine (Figure 4.7). Whereas,
DEA-carbamate started decomposing at much lower temperature (~55 °C) and the
transition was completed at about 70 °C, as is evident from TG/DSC plots in Figure 4.8. In
case of TG profile of AMP-carbamate, the weight loss can be seen originating much before
the decomposition onset temperature. AMP-carbamate, owing to its unstable nature in
humid air [31], most probably underwent hydrolysis to some extent generating free amine
during sample grinding/mounting process; the evaporation of which resulted in mass loss as
appeared in TG plot prior to the commencement of carbamate decomposition. The
hydrolytic transformation of AMP-carbamate might occur as follows (eq 11):
Figure 4. 7. DSC/TG profiles of AMP-carbamate: Thermal behavior observed under N2
atmosphere at heating rate of 5 °C/min.
C
CH2CH3
NHCOO-
CH3
OHC
CH2 CH3
NH3
+CH3
OH
AMP-Carbamate
C
CH2CH3
NH2CH3
OH
AMP
C
CH2 CH3
NH3
+CH3
OHHCO3
-H2O
+
bicarbonate protonated-AMP
-60
-50
-40
-30
-20
-10
0
0
20
40
60
80
100
30 40 50 60 70 80 90 100 110 120 130
He
at fl
ow
(m
W)
We
igh
t (%
)
Temperature ( C)
Weight
Heat flow
(11)
Chapter 4
103
Figure 4. 8. DSC/TG curves of DEA-carbamate: Thermal behavior under N2 atmosphere,
using heating rate of 5 °C/min.
To detect CO2 release, QMS was coupled with TG. The QMS signals showed the evolution
of CO2 above 70 °C in case of AMP-carbamate, and around 55 °C in case of DEA-
carbamate (Figure 4.9); thus complementing the TG/DSC analyses outcomes. The
temperature was increased at the rate of 5 °C/min under N2 (100 mL/min flow rate) and
continued until the positive molecular ion current intensity, originating from CO2+ (m/z =
44), reached the initial levels.
-20
-15
-10
-5
0
85
90
95
100
30 40 50 60 70 80 90 100 110 120 130
He
at fl
ow
(m
W)
We
igh
t (%
)
Temperature ( C)
Weight
Heat flow
Chapter 4
104
Figure 4. 9. QMS monitoring of carbamates‟ decomposition by measuring positive ion
current m/z = 44 (CO2) under N2 atmosphere (100 mL/min. flow rate) at 5 °C/min heating
rate.
Quite prolonged release of CO2, as appears in ion current versus time plots (obtained via
QMS), might be due to the foaming buildup as well as slow heat transfer at lower
temperatures (above decomposition point). Variations in ion current intensity possibly were
35 55 75 95 115 135
0.0E+00
5.0E-11
1.0E-10
1.5E-10
2.0E-10
2.5E-10
3.0E-10
0 5 10 15 20
Temperature (deg. C)i (C
O2,
44
) (A
)
Time (min.)
AMP-carbamate
35 75 115 155 195 235
0.0E+00
3.0E-11
6.0E-11
9.0E-11
0 10 20 30 40
Temperature (deg. C)
i (C
O2,
44
) (A
)
Time (min.)
DEA-carbamate
Chapter 4
105
fallout of change in foaming make-up with temperature. The foaming phenomenon was
also observed during ATR-FTIR analysis while studying regeneration behavior.
Thermal decomposition temperatures of both AMP-carbamate and DEA-carbamate were
also verified through temperature-programmed FTIR analysis, revealing the disappearance
of carbamate absorption peaks above 70 °C and 50 °C respectively (Figures C.4 and C.5 in
Appendix C).
However under 100% CO2 atmosphere, the beginning of decomposition was delayed
significantly (now starting at ~65 °C) regarding DEA-carbamate (Figure 4.10). While
apparent mass loss, observed under N2 atmosphere in case of AMP-carbamate below 75 °C
(decomposition onset temperature), appears to have been suppressed under CO2 cover. This
trend probably emerged due to the presence of one of the reactants (CO2) in excess.
Concerning AMP-carbamate, the CO2 atmosphere would also have helped revert some
proportion of free amine (stemmed from hydrolytic activity during sample preparation) to
carbamate thus curtailing the evaporation occurrence.
0
20
40
60
80
100
35 45 55 65 75 85 95 105 115 125
Wei
gh
t (%
)
Temperature ( C)
AMP-carbamate
Chapter 4
106
Figure 4. 10. TG profiles of carbamates: Thermal behavior under CO2 atmosphere, using
heating rate of 5 °C/min.
The observations stated above indicate that using RTIL, in place of water, can act as a
suitable medium for carbamate crystal growth thus allowing easy recovery of lower density
CO2-captured product. This not only can provide feasible opportunity to regenerate solely
active species but also can promise milder regeneration conditions. From regeneration
capabilities of AMP-/DEA-carbamates, it is quite obvious that DEA-RTIL blends can help
improve the process efficiency more successfully, regarding regeneration energy penalty in
particular. From perspective of amine evaporation loss, DEA-RTIL recipe is undoubtedly
better option compared to AMP-RTIL combination.
4.4. Implications
In case of alkanolamine based gas capture systems, better efficiency can be attained by
avoiding energy wastage during regeneration by targeting the active species (responsible
for CO2 capture) alone; and for this purpose incorporation of thermally stable RTIL can
provide with the prospect of CO2-captured product (carbamate) precipitation and thereby
easy separation. When compared to aqueous alkanolamine based processes, carbamate
crystallization in alkanolamine-RTIL systems is not only meant to lessen the quantity
required to regenerate but also can help narrow the gap between capture and regeneration
temperatures. Besides, with this strategy we may well overcome the difficulties being faced
0
20
40
60
80
100
35 45 55 65 75 85 95 105 115 125
Wei
gh
t (%
)
Temperature ( C)
DEA-carbamate
Chapter 4
107
regarding gas loading restraints (due to corrosion/degradation detriments) in current
alkanolamine based industrial processes [13,32].
In general, a secondary alkanolamine blended with pertinent RTIL can be a better pick for
CO2 capture as is evident from lower thermal stability of DEA-carbamate compared to that
of AMP-carbamate.
Since bringing about regeneration at lower temperature can help decrease the magnitude of
solvent degradation, future work will be focused on amine degradation studies using
alkanolamine-RTIL based CO2 capture processes. Moreover, measures/conditions will be
optimized to minimize foaming as well as evaporation phenomena.
4.5. References
[1] C. Cooney, Nations Seek "Fair" Greenhouse Gas Treaty in Kyoto, Environ. Sci.
Technol. 31 (1997) 516A-518A.
[2] H. Herzog, What future for carbon capture and sequestration? Environ. Sci. Technol. 35
(2001) 148A-153A.
[3] J. Figueroa, T. Fout, S. Plasynski, H. McIlvried, R. Srivastava, Advances in CO2
capture technology - The U.S. Department of Energy‟s Carbon Sequestration Program, Int.
J. Greenh. Gas Control 2 (2008) 9-20.
[4] J.C.M. Pires, F.G. Martins, M.C.M. Alvim-Ferraz, M. Simões, Recent developments on
carbon capture and storage: An overview, Chem. Eng. Res. Des. 89 (2011) 1446-1460.
[5] B. Metz, O. Davidson, H. de Coninck, M. Loos, L. Meyer, Eds., IPCC Special Report
on Carbon Dioxide Capture and Storage, Prepared by Working Group III of the
Intergovernmental Panel on Climate Change, Cambridge University Press: New York,
2005.
[6] J. Brennecke, B. Gurkan, Ionic Liquids for CO2 Capture and Emission Reduction, J.
Phys. Chem. Lett. 1 (2010) 3459-3464.
Chapter 4
108
[7] J. Kittel, E. Fleury, B. Vuillemin, S. Gonzalez, F. Ropital, R. Oltra, Corrosion in
alkanolamine used for acid gas removal: From natural gas processing to CO2 capture,
Mater. Corros. 63 (2012) 223-230.
[8] M. Caplow, Kinetics of carbamate formation and breakdown, J. Am. Chem. Soc. 90
(1968) 6795-6803.
[9] P.V. Danckwerts, The reaction of CO2 with ethanolamines, Chem. Eng. Sci. 34 (1979)
443-446.
[10] Z. Pei, S. Yao, W. Jianwen, Z. Wei, Y. Qing, Regeneration of 2-amino-2-methyl-1-
propanol used for carbon dioxide absorption, J. Environ. Sci. 20 (2008) 39-44.
[11] J.N. Knudsen, J.N. Jensen, P.-J. Vilhelmsen, O. Biede, Experience with CO2 capture
from coal flue gas in pilot-scale: testing of different amine solvents, Energy Procedia 1
(2009) 783-790.
[12] S. Chi, G. Rochelle, Oxidative Degradation of Monoethanolamine, Ind. Eng. Chem.
Res. 41 (2002) 4178-4186.
[13] I. Soosaiprakasam, A. Veawab, Corrosion and polarization behavior of carbon steel in
MEA-based CO2 capture process, Int. J. Greenh. Gas Control 2 (2008) 553-562.
[14] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Ionic liquids for CO2 capture – development
and progress, Chem. Eng. Process. 49 (2010) 313-322.
[15] J. Anthony, J. Anderson, E. Maginn, J. Brennecke, Anion Effects on Gas Solubility in
Ionic Liquids, J. Phys. Chem. B 109 (2005) 6366-6374.
[16] J. Anderson, J. Dixon, J. Brennecke, Solubility of CO2, CH4, C2H6, C2H4, O2, and N2
in 1-Hexyl-3-methylpyridinium Bis(trifluoromethylsulfonyl)imide: Comparison to Other
Ionic Liquids, Acc. Chem. Res. 40 (2007) 1208-1216.
[17] J. Bara, T. Carlisle, C. Gabriel, D. Camper, A. Finotello, D. Gin, R. Noble, Guide to
CO2 Separations in Imidazolium-Based Room-Temperature Ionic Liquids, Ind. Eng. Chem.
Res. 48 (2009) 2739-2751.
Chapter 4
109
[18] J. Bara, D. Camper, D. Gin, R. Noble, Room-Temperature Ionic Liquids and
Composite Materials: Platform Technologies for CO2 Capture, Accounts Chem. Res. 43
(2010) 152-159.
[19] A. Yokozeki, M. Shiflett, Separation of Carbon Dioxide and Sulfur Dioxide Gases
Using Room-Temperature Ionic Liquid [hmim][Tf2N], Energy Fuels 23 (2009) 4701-4708.
[20] A.L. Kohl, R.B. Nielsen, Gas Purification, 5th ed. Gulf Publishing Company:
Houston, Texas, 1997.
[21] X. Gui, Z. Tang, W. Fei, CO2 Capture with Physical Solvent Dimethyl Carbonate at
High Pressures, J. Chem. Eng. Data 55 (2010) 3736-3741.
[22] E.G. Langeroudi, F. Kleitz, M.C. Iliuta, F. Larachi, Grafted Amine/CO2 Interactions in
(Gas−)Liquid−Solid Adsorption/Absorption Equilibria, J. Phys. Chem. C 113 (2009)
21866-21876.
[23] D. Camper, J.E. Bara, D.L. Gin, R.D. Noble, Room-temperature ionic liquid-amine
solutions: tunable solvents for efficient and reversible capture of CO2, Ind. Eng. Chem. Res.
47 (2008) 8496–8498.
[24] Q. Huang, Y. Li, X. Jin, D. Zhao, G.Z. Chen, Chloride ion enhanced thermal stability
of carbon dioxide captured by monoethanolamine in hydroxyl imidazolium based ionic
liquids, Energy Environ. Sci. 4 (2011) 2125-2133.
[25] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 capture in alkanolamine/room-
temperature ionic liquid emulsions: A viable approach with carbamate crystallization and
curbed corrosion behavior, Int. J. Greenh. Gas Control 6 (2012) 246-252.
[26] M. Hasib-ur-Rahman, H. Bouteldja, P. Fongarland, M. Siaj, F. Larachi, Corrosion
Behavior of Carbon Steel in Alkanolamine/Room-Temperature Ionic Liquid Based CO2
Capture Systems, Ind. Eng. Chem. Res. 51 (2012) 8711-8718.
[27] H.A. Al-Ghawas, D.P. Hagewlesche, G. Rulz-Ibanez, O.C. Sandall, Physicochemical
Properties Important for Carbon Dioxide Absorption in Aqueous Methyldiethanolamine, J.
Chem. Eng. Data 34 (1989) 385-391.
Chapter 4
110
[28] R. Crovetto, Evaluation of Solubility Data of the System CO2-H2O from 273 K to the
Critical Point of Water, J. Phys. Chem. Ref. Data 20 (1991) 575-589.
[29] K. Klepacova, P.J.G. Huttenhuis, P.W.J. Derks, G.F. Versteeg, Vapor Pressures of
Several Commercially Used Alkanolamines, J. Chem. Eng. Data 56 (2011) 2242-2248.
[30] T. Nguyen, M. Hilliard, G.T. Rochelle, Amine volatility in CO2 capture, Int. J.
Greenh. Gas Control 4 (2010) 707-715.
[31] E. Jo, Y.H. Jhon, S.B. Choi, J.-G. Shim, J.-H. Kim, J.H. Lee, I.-Y. Lee, K.-R. Jang, J.
Kim, Crystal structure and electronic properties of 2-amino-2-methyl-1-propanol (AMP)
carbamate, Chem. Commun. 46 (2010) 9158-9160.
[32] G.T. Rochelle, Thermal degradation of amines for CO2 capture, Current Opinion in
Chemical Engineering 1 (2012) 183-190.
Chapter 5
111
Kinetic behavior of carbon dioxide absorption in diethanolamine/ionic-
liquid emulsions*
Abstract/Résumé
Room-temperature ionic liquids (RTILs) have been found to induce precipitation of CO2-
captured carbamate product in case of amine-RTIL systems that may lead to an efficient
carbon dioxide capture process. We have evaluated the kinetic behaviour of CO2 absorption
in DEA-[hmim][Tf2N] blends in a laboratory scale stirred-cell reactor at ambient pressure
(~1 atm) to assess the effects of amine concentration (≤ 2M DEA), CO2 partial pressure,
agitation speed (1500-4500 rpm), and temperature variation (25°C to 41°C). A CO2 probe
was used to monitor the change in gaseous CO2 volume ratio during the absorption
experiments. It was evident from the outcome that with the increase in CO2 percentage in
the simulated gaseous mixture (CO2+N2), the gas absorption rate was correspondingly
improved. Since the constituents (DEA and [hmim][Tf2N]) were immiscible, agitation
speed appeared to have a significant influence on CO2 absorption behaviour resulting most
probably from the better dispersion of amine droplets at higher homogenising speeds.
Les liquides ioniques à température ambiantes (RTILs) conduisent à la précipitation des
carbamates produits par la capture de CO2 dans des mélanges amine-RTIL qui peuvent
conduire à un processus efficace de capture du dioxyde de carbone. Nous avons évalué le
comportement cinétique d'absorption du CO2 dans des mélanges DEA-[hmim][Tf2N] dans
un réacteur agité, à l‟échelle laboratoire, à pression ambiante (~1 atm) afin d'évaluer les
effets de la concentration en amine (≤ 2M DEA), la pression partielle de CO2, la vitesse
d‟agitation (1500-4500 rpm), et la température (25°C to 41°C). Une sonde de CO2 a été
utilisée pour suivre la variation de la composition gazeuse en CO2 au cours des expériences
d'absorption. Il ressort clairement des résultats que l'augmentation de la proportion de CO2
dans le mélange gazeux accélère le taux d'absorption de gaz. Étant donné que les
constituants (DEA and [hmim][Tf2N]) sont immiscibles, la vitesse d'agitation semble
influencer significativement le comportement d‟absorption du CO2 comme résultat direct
* M. Hasib-ur-Rahman, F. Larachi, Sep. Purif. Technol. Submitted February 2013.
Chapter 5
112
d'une meilleure dispersion de gouttelettes d'amines à des vitesses plus élevées
d'homogénéisation.
5.1. Introduction
To quench the anthropogenic CO2 emissions into the atmosphere and hence controlling the
global warming phenomenon resulting from greenhouse gases buildup, efficient gas
separation systems are needed [1-3]. In this regard, large point sources such as fossil-fueled
power plants are the most convenient sites for CO2 capture. State-of-the-art aqueous
alkanolamines are the most developed schemes being employed widely in natural gas
purification installations. The major hindrance in large scale application of aqueous
alkanolamine based CO2 capture processes is the unaffordably high regeneration energy
requirement [4]. Equilibrium limitations, equipment corrosion, and amine degradation are
some other drawbacks of the process, mainly inherited by the aqueous moiety [5-10].
Park et al. tried an alternate methodology by dispersing aqueous amine droplets in benzene
in order to enhance the gas capture efficiency of alkanolamine based processes by taking
advantage of the concept that in a gas-liquid system the water-in-oil (emulsion) type
arrangement can enhance mass transfer of the dissolved gas [11,12]. But such an approach
only appeared to add to environmental concerns arising from the use of toxic and volatile
organic solvents. Besides, the use of surfactant further complicated the process. Work has
also been underway to merge the advantages of both aqueous amines with those of ionic
liquids evading some of the drawbacks posed by higher viscosity of pure ionic liquids [13-
15]. Still it was hard to fully elude the drawbacks solely related to the presence of water as
shown by Hamah-Ali et al. regarding corrosion occurrence [16]. Nonetheless, all these
strategies seem unlikely to impart any significant improvement to alkanolamine based CO2
capture techniques [11-16].
Consequently, it may be a viable approach to replace aqueous phase wholly with more
stable and secure solvent such as a room-temperature ionic liquid (RTIL). Being thermally
stable, virtually non-volatile, as well as possessing lower heat capacities [17-19], RTILs
may lead to an energy efficient pathway to CO2 capture and amine regeneration. Moreover,
Chapter 5
113
availability of numerous combinations of constituent ionic counterparts makes it quite
feasible to tailor an ionic liquid in accordance with the required specifications.
Typically imidazolium based ionic liquids either solely [20-22] or in combination with
alkanolamines [23-27] are being investigated as potential alternates for the current
physical/chemical absorption processes. Among these, the most striking aspect of
alkanolamine-RTIL combinations is the emergence of carbamate (CO2-captured product)
precipitation that not only helps reach stoichiometric maximum gas loading capacity but
also provides the opportunity to separate CO2-captured product, thereby offering likely
reduction in regeneration energy. Also the suppression of corrosion occurrence particularly
in case of gas absorption system comprising alkanolamine and hydrophobic ionic liquid
adds further value to the process [25,26].
However, there have not been enough methodical efforts to assess the practicability of
amine-RTIL based CO2 separation schemes. Accordingly, the objective of the current study
was to scrutinize the kinetic aspects of such systems. To achieve this goal, CO2 absorption
behaviour was monitored using different amine concentrations and varying gas partial
pressures. Moreover, the influence of agitation speed and temperature was also
investigated. The exercise was conducted in a continuously stirred-cell reactor to probe the
role of above stated experimental variables regarding CO2 capture in (immiscible) DEA-
[hmim][Tf2N] blends.
5.2. Reaction mechanism in non-aqueous amines
For chemical absorption of CO2 in alkanolamine based systems, the major reaction
comprises the carbamate formation involving CO2 and amine interaction in 1:2 molar ratio
respectively. Considering primary/secondary alkanolamines, zwitterion mechanism is the
most widely accepted model first proposed by Caplow in 1968 [28] and later reiterated by
Danckwerts [29]. This mechanism involves the formation of an intermediate (zwitterion) in
the first step that follows the abstraction of proton by a base:
1 2 2 1 2
1 2 1 2
R R NH CO R R NH COO
R R NH COO B R R NCOO BH
Chapter 5
114
In aqueous amines the deprotonation species (B) include water, OH-, and amine itself [30]
but, contrary to aqueous amine systems, in non-aqueous media primary/secondary amine
can be the only base available to deprotonate the zwitterion [31], and hence the gas loading
capacity becomes limited to 0.5 mole CO2 per mole of amine (stoichiometric maximum).
Thus the reaction can be presented as follows:
1 2 2 1 2
1 2 1 2 1 2 1 2 2
R R NH CO R R NH COO
R R NH COO R R NH R R NCOO R R NH
The same is pertinent to the amine-RTIL blends as the room-temperature ionic liquid does
not involve in any kind of chemical interaction either with CO2 or with amine [23-27].
5.3. Experimental
5.3.1. Materials
A secondary alkanolamine, diethanolamine (DEA: ACS reagent, ≥99.0%), was purchased
from Sigma-Aldrich while the room-temperature ionic liquid, 1-hexyl-3-
methylimidazoilium bis(trifluoromethylsulfonyl)imide ([hmim][Tf2N]: 99%), was provided
by IoLiTec Inc. Carbon dioxide and nitrogen (≥99% purity) gases were obtained from
Praxair Canada Inc.
5.3.2. Setup
Gas absorption experiments were carried out in a double jacketed stirred-cell reactor as
shown schematically in figure 5.1. An Omni homogenizer, fitted with rotor-stator
generator, was immersed in the cell to agitate the liquid during absorption experiments
whereas a CO2 probe (GMP221, Vaisala) was positioned in the headspace to monitor
volumetrically the CO2 consumption rate. A K type thermocouple was used to measure the
temperature of the liquid sample during the course of experiments. The reactor volume was
100 ml. The gaseous mixture was continuously circulated between the absorption cell and
the reservoir (18.7 L volume) with the help of a peristaltic pump at a flow rate of 1±0.01
L/min. The gas inlet comprised two partitions, one exiting the gas in the headspace area
while the second continued till underneath the homogeniser in the liquid phase to help
Chapter 5
115
enhance gas sparging. The temperature of the stirred-cell reactor as well as of the
headspace area was controlled by a thermostatic bath.
5.3.3. Procedure
Each time, prior to the absorption experiments, the setup was purged with nitrogen gas to
remove any gaseous contaminant. Then the gas reservoir was filled with desired
proportions of CO2 and nitrogen using Bronkhorst mass-flow controllers. After the
introduction of a specified volume of pure RTIL into the cell through an inlet needle, the
gaseous mixture was continuously recirculated for 120 min with the help of a peristaltic
pump so that, under the given circumstances, the RTIL became saturated with CO2, till the
probe showed stable reading. Subsequently a known quantity of DEA (being immiscible
with the ionic liquid) was injected into the RTIL containing cell reactor and the process was
continued for 3 h. For each experiment, 12 ml of DEA-[hmim][Tf2N] fluid was used.
During the experiment, the liquid was constantly stirred using an Omni homogeniser fitted
with rotor-stator generator. CO2 probe was linked to a computerized acquisition system,
delivering data in terms of CO2 available as volume % in the gaseous mixture. This allowed
the calculation of CO2 absorption per unit time.
Figure 5. 1. Experimental set-up scheme: A) Gas inlet; B) Gas outlet (A & B connect to a
gas reservoir via closed loop system); C) CO2 probe; D) Injection port; E) Thermocouple;
F) Rotor-stator homogeniser; G) Absorption cell; Hi) Heating bath inlet; Ho) Heating bath
outlet.
Chapter 5
116
5.4. Results and Discussion
As has been observed during the earlier work [23-27], primary/secondary alkanolamines
when blended with room-temperature ionic liquids offer some exceptional advantages
regarding CO2 capture. Absorption of CO2 in amine-RTIL blends results in precipitation of
the CO2-captured product (carbamate) as shown in Figure 5.2.
Figure 5. 2. CO2-captured product (carbamate) precipitation in DEA-[hmim][Tf2N]: a)
immediately after CO2 bubbling; b) 24 hours later.
Since there is no supplementary deprotonating species except amine in DEA-RTIL system,
the maximum loading capacity does not exceed 50 mol% of CO2 as primary/secondary
amine (DEA in this case) reacts with CO2 in 2:1 ratio obeying the following reaction:
2DEA CO DEAH COO
DEAH COO DEA DEACOO DEAH
During the course of the current experiments, effects of various parameters, i.e., amine
concentration, CO2 gaseous ratio, agitation speed, and temperature were considered to
study the CO2 uptake behaviour in DEA-[hmim][Tf2N] mixtures using stirred-cell reactor.
(a) (b)
Chapter 5
117
5.4.1. Impact of variation in amine concentration
The CO2 absorption mode shows two distinct regions in the curve, as shown in Figure 5.3.
The initial steeper part depicts an abrupt gas absorption phenomenon that seems to have
occurred through mutual contribution of physically confined CO2 in the RTIL (solubilized
prior to the injection of amine into the stirred-cell) and the additional CO2 approaching
directly via continuous gas bubbling. While the other somewhat horizontal portion could
have evolved after the unreacted amine started accumulating over RTIL surface. As diluted
gaseous mixture containing CO2 ≤ 10 vol% was used to observe the gas absorption trends
of DEA-[hmim][Tf2N] blends, higher amine content (DEA: 2M) did not appear to be
compatible with the experimental conditions and consequently the absorption rate was the
lowest among the four concentrations of DEA tested. However, decrease in amine content
corresponded well to the low CO2 gaseous ratio. The results thus depict that the gas-liquid
contact zone (between CO2 and dispersed amine), in case of higher amine concentration,
was not large enough to accommodate most of the active sites (the dispersed amine
droplets) inside the bulk room-temperature ionic liquid phase. Neither the immiscibility as
well as the difference in the respective densities (1.09 g/cm3 and 1.37 g/cm
3) of both the
components, DEA and [hmim][Tf2N], let the amine droplets to stay dispersed long enough.
This situation caused a significant amount of unreacted amine to agglomerate to the surface
of the RTIL (the coalescence of amine droplets accelerated as the amine content was
increased) and continued to capture CO2 but at a much slower pace, as is obvious from the
CO2 uptake (mole of CO2 per mole of amine vs time) curves for DEA-RTIL blends with
higher amine ratios (Figure 5.3).
Chapter 5
118
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
CO2 Concentration
10.0%
5.0%
2.5%
(a)
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
CO2 Concentration
10.0%
5.0%
2.5%
(b)
Chapter 5
119
Figure 5. 3. Influence of [DEA] molar concentration on absorption rate with respect to
initial CO2 vol% in the gaseous mixture, at 33 °C and 3000 rpm agitation speed: a) 2M
DEA in [hmim][Tf2N]; b) 1M DEA in [hmim][Tf2N]; c) 0.5M DEA in [hmim][Tf2N].
Smoothed lines show trends.
5.4.2. CO2 volume ratio in the gaseous mixture
Influence of the variation of CO2 vol% (in the gaseous mixture) on absorption also
corroborates the discussion in the previous section. As is evident from the results presented
in figure 5.4, the 0.5M DEA appears well-suited to the gaseous mixture containing 10 vol%
CO2 for quick gas absorption. However, as the gaseous CO2 concentration was lowered the
capture rate decreased accordingly. In case of DEA-RTIL blends with higher amine ratio
(1-2M DEA), even 10 vol% CO2 was not sufficient to drive the process quickly to
maximum gas loading.
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
CO2 Concentration 10.0%
5.0%
2.5%
(c)
Chapter 5
120
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
[DEA] in [hmim][Tf2N]
(a)
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
[DEA] in [hmim][Tf2N]
(b)
Chapter 5
121
Figure 5. 4. Influence of initial CO2 volume ratio (in gaseous mixture) on absorption rate
w.r.t. [DEA], at 33 °C and 3000 rpm agitation speed: a) 10 vol% CO2; b) 5 vol% CO2; c)
2.5 vol% CO2. Smoothed lines show trends.
5.4.3. Influence of agitation speed
Since diethanolamine and [hmim][Tf2N] are immiscible and there is significant density
difference between the two (DEA: 1.09 g/cm3; [hmim][Tf2N]: 1.37 g/cm
3), it is hard to
keep DEA dispersed in [hmim][Tf2N] without the addition of a surfactant. Yet, proper
agitation can not only help induce DEA dispersion for extended duration, but also can
provide with increased surface area of DEA to interact with CO2. As has been shown in
figure 5.5, increase in agitation speed from 1500 rpm to 4500 rpm (keeping other variables
constant: 2 M DEA; 10 vol% CO2; 33 °C) caused faster CO2 absorption. This seems to be
the outcome of greater residence time of dispersed amine inside the RTIL phase and/or
smaller amine droplet size. Thus agitation speed can be optimized in accordance with the
flue gas composition and the other experimental parameters (such as amine ratio, gas flow
rate, process temperature, etc.) to acquire a sufficient absorption rate.
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
[DEA] in [hmim][Tf2N]
(c)
Chapter 5
122
Figure 5. 5. Influence of agitation on CO2 absorption rate (2M DEA in [hmim][Tf2N]; 10
vol% CO2; 33 °C). Smoothed lines show trends.
5.4.4. Effect of temperature variation
Three different temperatures (25, 33 and 41 °C) were chosen to evaluate the role of
temperature on CO2 absorption behaviour. In spite of the fact that increase in temperature
resulted in decreased liquid viscosity (Table 5.1) and hence gas transfer rate could have
improved, the experimental outcome did not depict any systematic change in capture rate as
shown in figure 5.6. Decrease in physical solubility of CO2 in RTIL at higher temperature
might have undone the lower viscosity advantage if there was any. This behaviour suggests
that CO2-captured product (carbamate) precipitation might be the dictating factor that
possibly has overshadowed the influence of temperature on CO2 absorption rate.
Table 5. 1. Viscosities* of the capture fluid components at three temperatures.
Component 25 °C 33 °C 41 °C
DEA 470 cP 241 cP 139 cP
[hmim][Tf2N] 61 cP 38 cP 26 cP
*measured by AR-G2 rheometer (TA Instruments) with parallel plate geometry
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mo
lar
rati
o)
Time (min)
Agitation Speed
Chapter 5
123
Figure 5. 6. Effect of temperature on CO2 capture rate (1M DEA in [hmim][Tf2N]; 10
vol% CO2; 3000 rpm). Smoothed lines show trends.
As in the simulated gaseous mixture the CO2 ratio (opposed to the pure gas stream) was
maintained within the concentration range of post-combustion flue gases (˂ 15%), the CO2
solubility in the RTIL phase must had undergone a negative impact [32,33]. The CO2
absorption trends (Figures 5.3, 5.4) reveal the fact that amine fraction in the blend should
be defined in accordance with the CO2 percentage in the flue gas. Besides, the agitation
speed which dictates the mass transfer phenomena within the gas capturing fluid should
also be taken into consideration. The immiscible nature of DEA as well as a significant
difference between the densities of both constituents (DEA and [hmim][Tf2N]) renders it
compulsory to agitate the fluid during gas capture.
5.5. Conclusion
The results of this study reveal that though amine-RTIL blends are blessed with a unique
advantage, i.e., CO2-captured product (carbamate) precipitation, there are some other
factors that should be taken into consideration for profiting from maximal absorption
capabilities of the immiscible amine-RTIL systems:
0
0.1
0.2
0.3
0.4
0.5
0.6
0 50 100 150
CO
2u
pta
ke
(mola
r ra
tio)
Time (min)
Temperature Variation
Chapter 5
124
Amine ratio in the blends should be specified according to the CO2 proportion in
flue gas.
As both components (DEA and [hmim][Tf2N] in the present case) are immiscible,
continuous agitation is an obligatory requirement for the process to achieve better
and prolonged dispersion of amine inside RTIL phase.
Even though temperature did not appear to have a momentous effect on CO2
absorption rate, maximum possible temperature during absorption can be quite
advantageous; firstly, by easing the processing through decrease in fluid viscosity
and, secondly, by lowering the gap between absorption and regeneration
temperatures.
The experimental findings may as well help carve the way out towards designing a
pertinent absorption column.
5.6. References
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[2] C. Stewart, M.-A. Hessami, A study of methods of carbon dioxide capture and
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[3] B. Metz, O. Davidson, H. de Coninck, M. Loos, L. Meyer, (Eds.), IPCC Special Report
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[4] R. Idem, M. Wilson, P. Tontiwachwuthikul, A. Chakma, A. Veawab, A. Aroonwilas, D.
Gelowitz, Pilot plant studies of the CO2 capture performance of aqueous MEA and mixed
MEA/MDEA solvents at the University of Regina CO2 Capture Technology Development
Plant and the Boundary Dam CO2 Capture Demonstration Plant, Ind. Eng. Chem. Res. 45
(2006) 2414-2420.
Chapter 5
125
[5] F. Barzagli, F. Mani, M. Peruzzini, Continuous cycles of CO2 absorption and amine
regeneration with aqueous alkanolamines: a comparison of the efficiency between pure and
blended DEA, MDEA and AMP solutions by 13
C NMR spectroscopy, Energy Environ. Sci.
3 (2010) 772-779.
[6] S. Bishnoi, G.T. Rochelle, Absorption of carbon dioxide into aqueous piperazine:
reaction kinetics, mass transfer and solubility, Chem. Eng. Sci. 55 (2000) 5531-5543.
[7] I.R. Soosaiprakasam, A. Veawab, Corrosion and polarization behavior of carbon steel in
MEA-based CO2 capture process, Int. J. Greenh. Gas Control 2 (2008) 553-562.
[8] M. Hasib-ur-Rahman, F. Larachi, Corrosion in amine systems - a review, Carbon
Capture Journal, Sept - Oct 2012, 22-24.
[9] B.R. Strazisar, R.R. Anderson, C.M. White, Degradation pathways for
monoethanolamine in a CO2 capture facility, Energy Fuels 17 (2003) 1034-1039.
[10] S. Chi, G. Rochelle, Oxidative Degradation of Monoethanolamine, Ind. Eng. Chem.
Res. 41 (2002) 4178−4186.
[11] S.W. Park, H.B. Cho, I.J. Sohn, H. Kumazawa, CO2 absorption into w/o emulsion with
aqueous amine liquid droplets, Sep. Sci. Technol. 37 (2002) 639-661.
[12] V. Linek, P. Beneš, A study of the mechanism of gas absorption into oil-water
emulsions, Chem. Eng. Sci. 31 (1976) 1037-1046.
[13] A Ahmady, M.A. Hashim, M.K. Aroua, Experimental Investigation on the Solubility
and Initial Rate of Absorption of CO2 in Aqueous Mixtures of Methyldiethanolamine with
the Ionic Liquid 1-Butyl-3-methylimidazolium Tetrafluoroborate, J. Chem. Eng. Data 55
(2010) 5733-5738.
[14] N.A. Sairi, R. Yusoff, Y. Alias, M.K. Aroua, Solubilities of CO2 in aqueous N-
methyldiethanolamine and guanidinium trifluoromethanesulfonate ionic liquid systems at
elevated pressures, Fluid Phase Equilibr. 300 (2011) 89-94.
Chapter 5
126
[15] Z. Feng, F. Cheng-Gang, W. You-Ting, W. Yuan-Tao, L. Ai-Min, Z. Zhi-Bing,
Absorption of CO2 in the aqueous solutions of functionalized ILs and MDEA, Chem. Eng.
J. 160 (2010) 691-697.
[16] B. Hamah-Ali, B.S. Ali, R. Yusoff, M.K. Aroua, Corrosion of Carbon Steel in
Aqueous Carbonated Solution of MEA/[bmim][DCA], Int. J. Electrochem. Sci. 6 (2011)
181-198.
[17] J.F. Brennecke, E.J. Maginn, Ionic Liquids: Innovative Fluids for Chemical
Processing, AIChE J. 47 (2001) 2384-2389.
[18] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, Ionic Liquids for CO2 Capture -
Development and Progress, Chem. Eng. Process 49 (2010) 313-322.
[19] D. Waliszewski, I. Stepniak, H. Piekarski, A. Lewandowski, Heat capacities of ionic
liquids and their heats of solution in molecular liquids, Thermochim. Acta 433 (2005) 149-
152.
[20] D. Camper, P. Scovazzo, C. Koval, R. Noble, Gas Solubilities in Room-Temperature
Ionic Liquids, Ind. Eng. Chem. Res. 43 (2004) 3049-3054.
[21] A.M. Schilderman, S. Raeissi, C.J. Peters, Solubility of carbon dioxide in the ionic
liquid 1-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide, Fluid Phase
Equilib. 260 (2007) 19-22.
[22] A. Yokozeki, M.B. Shiflett, Separation of Carbon Dioxide and Sulfur Dioxide Gases
Using Room-Temperature Ionic Liquid [hmim][Tf2N], Energy Fuels 23 (2009) 4701-4708.
[23] D. Camper, J.E. Bara, D.L. Gin, R.D. Noble, Room-Temperature Ionic Liquid-Amine
Solutions: Tunable Solvents for Efficient and Reversible Capture of CO2, Ind. Eng. Chem.
Res. 47 (2008) 8496-8498.
[24] Q. Huang, Y. Li, X. Jin, D. Zhao, G.Z. Chen, Chloride ion enhanced thermal stability
of carbon dioxide captured by monoethanolamine in hydroxyl imidazolium based ionic
liquids, Energy Environ. Sci. 4 (2011) 2125-2133.
Chapter 5
127
[25] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 Capture in Alkanolamine/Room-
Temperature Ionic Liquid Emulsions: A Viable Approach with Carbamate Crystallization
and Curbed Corrosion Behavior, Int. J. Greenhouse Gas Control 6 (2012) 246-252.
[26] M. Hasib-ur-Rahman, H. Bouteldja, P. Fongarland, M. Siaj, F. Larachi, Corrosion
behavior of carbon steel in alkanolamine/room-temperature ionic liquid based CO2 capture
systems, Ind. Eng. Chem. Res. 51 (2012) 8711-8718.
[27] M. Hasib-ur-Rahman, F. Larachi, CO2 Capture in Alkanolamine-RTIL Blends via
Carbamate Crystallization: Route to Efficient Regeneration, Environ. Sci. Technol. 46
(2012) 11443-11450.
[28] M. Caplow, Kinetics of carbamate formation and breakdown, J. Am. Chem. Soc. 90
(1968) 6795-6803.
[29] P.V. Danckwerts, The reaction of CO2 with ethanolamines, Chem. Eng. Sci. 34 (1979)
443-446.
[30] P.M.M. Blauwhoff, G.F. Versteeg, W.P.M. van Swaaij, A study on the reaction
between CO2 and alkanolamines in aqueous solutions, Chem. Eng. Sci. 38 (1983) 1411-
1429.
[31] G.F. Versteeg, L.A.J. van Dijck, W.P.M. van Swaaij, On the kinetics between CO2 and
alkanolamines both in aqueous and non-aqueous solutions. An overview, Chem. Eng.
Commun. 144 (1996) 113-158.
[32] C.W. Jones, CO2 Capture from Dilute Gases as a Component of Modern Global
Carbon Management, Annu. Rev. Chem. Biomol. Eng. 2 (2011) 31-52.
[33] R.E. Baltus, B.H. Culbertson, S. Dai, H. Luo, D.W. DePaoli, Low-Pressure Solubility
of Carbon Dioxide in Room-Temperature Ionic Liquids Measured with a Quartz Crystal
Microbalance, J. Phys. Chem. B 108 (2004) 721-727.
Chapter 6
129
Conclusions and recommendations
6.1. General conclusions
Anthropogenic emission of greenhouse gases, predominantly carbon dioxide, is a matter of
great concern with reference to the consequences of global warming phenomenon. To
mitigate emissions, carbon dioxide capture from large point sources especially involving
power generation can be a viable practice and is being investigated extensively at present.
Yet there is no quick and easy way to decrease the emissions to acceptable level, as current
CO2 capture technologies would increase the cost of electricity production by 35-70%
(IPCC, 2005). State-of-the-art aqueous alkanolamine based chemical absorption processes
are in use industrially since 1930s. But the major impediments in this regard are high
energy consumption, equilibrium limitations, equipment corrosion, and solvent loss. Most
of these issues are directly related to the presence of water. One feasible route may be the
replacement of aqueous phase with some stable solvent. Room-temperature ionic liquids
(RTILs), with tunable physicochemical nature, higher thermal stability and practically no
volatility even at elevated temperatures, are emerging as promising candidates. More
importantly RTILs have shown quite significant affinity for CO2. In this perspective,
combining alkanolamines with RTILs can provide an opportunity to couple the chemical
and physical capabilities. Switching from aqueous to organic phase can also be productive
enough to alleviate some of the problems posed by aqueous amine systems.
Accordingly, the current experimental strategy was aimed at investigating the CO2 capture
potential of alkanolamine-RTIL combinations, and to ascertain if these new schemes can
counter the drawbacks of aqueous counterparts. It involved studying CO2 absorption
behaviour using primary/secondary alkanolamines (2-amino-2-methyl-1-propanol,
monoethanolamine, or diethanolamine) blended in imidazolium based room-temperature
ionic liquids (hydrophobic/hydrophilic). The CO2-captured product‟s separation and amine
regeneration trends were also probed. Besides, carbon steel 1020 was selected to look into
the corrosivity of these amine-RTIL blends.
Chapter 6
130
The experimental results can be reviewed in the following three sections, i.e. CO2
absorption behaviour, amine regeneration, and corrosion studies:
CO2 absorption behaviour
The most significant aspect of the amine-RTIL based gas capture fluids, involving
primary/secondary alkanolamine, is the phase changing characteristic as CO2-
captured product (carbamate) is insoluble in the studied RTILs. Opposed to what is
observed in aqueous amine processes, this feature helped avoid equilibrium
limitation thus enabling the absorption process to continue at a good rate in spite of
higher viscosity as has been experienced in case of DEA-[hmim][Tf2N] emulsions.
This factor also helped reach maximum stoichiometric capture capacity. The
characterization of solid products (AMP-carbamate and DEA-carbamate) confirmed
that there was no direct involvement (chemical interaction) of RTIL in the
formation of CO2 absorption product. Hence ionic liquid phase served as reservoir
to hold the product as precipitate.
However, to avoid the use of a surfactant in case of immiscible DEA-RTIL blends,
an apposite agitation technique is required to maintain better and prolonged
dispersion of amine droplets inside the RTIL phase in order to attain good CO2
absorption kinetics.
Amine regeneration
Carbamate products, resulting from CO2 absorption by primary/secondary amines,
crystallized out as supernatant solid. This behavior eased product removal and
offered smaller volume to regenerate, thus promising a probable decrease in
regeneration energy requirement. As secondary amine carbamates are less stable
compared to those of primary alkanolamines, combining diethanolamine (DEA)
with RTIL proved to be a better choice. Regeneration of DEA-carbamate
commenced at about 55 °C and this aspect can further help cut down regeneration
energy needs. In order to know the regeneration capabilities, AMP-carbamate as
well as DEA-carbamate was heated at 110 °C to enable quick release of CO2. After
Chapter 6
131
single absorption/desorption cycle, no amine degradation products were formed as
was confirmed by NMR/FTIR analyses of the regenerated amines.
Corrosion studies
Both hydrophobic and hydrophilic room-temperature ionic liquids were tested
regarding the corrosion phenomenon.
At room temperature (25 °C), alkanolamine/s and hydrophilic RTIL mixtures
showed excellent corrosion control (˂ 1 mpy) even in the presence of water and
oxygen, primarily owing to the coating of working electrode (carbon steel 1020)
surface by RTIL. However, higher temperature (60 °C) caused depletion of the
RTIL protective layer making the electrode surface vulnerable to the corrosive
action of oxidants (probably originating from water impurities). Still, compared to
the corresponding aqueous amine, amine-RTIL blends exhibited much better
performance by reducing corrosion rate up to 72% at 60 °C.
On the other hand, amine-RTIL blends involving hydrophobic ionic liquids proved
to be excellent in nullifying the corrosion occurrence even at higher temperature,
primarily by establishing a water repellent environment.
6.2. Future work recommendations
The experimental outcomes appear to be quite inspiring; however, to prove these novel
schemes to be practicable further exploration is required.
As has been observed that one absorption/desorption cycle did not induce any amine
degradation; detailed examination is needed to fully assess the maximum number of
cycles that can be run with one batch without any loss of absorption capability
through amine degradation. Moreover it will also be worthwhile to explore to what
extent it can be possible to control amine loss via evaporation specifically during
regeneration step and what measures can be taken in this regard.
Chapter 6
132
It needs to be researched how flue gas impurities (such as NOx, SOx, water vapours,
generally found in power plants‟ flue gases) can affect the absorption performance
of amine-RTIL blends.
Though immiscible amine-RTIL combinations (DEA-[Cnmim][Tf2N]) appeared to
be advantageous regarding product (solid carbamate) separation and regeneration, to
fully profit from these immiscible blends, a better mixing/processing technique is
needed to help keep amine phase dispersed in the bulk ionic liquid long enough
during gas absorption activity to reach maximum capture capacity at an acceptable
rate.
A traditional procedure would not apply when there are two such phases. So to take
advantage of the approach of carbamate (CO2-captured product) separation to
minimize regeneration volumes, a major diversion from conventional processing
will be required.
Appendix A
133
Supporting Information (Chapter 2)
Table A.1. Bond lengths and angles in CO2-captured product (DEA-carbamate).
Figure A.1. Packing diagram of the C9H22N2O6 compound (DEA-carbamate). Hydrogen atoms not participating in hydrogen bonding
were omitted for clarity.
Figure A.2. Solid-state 13
C NMR of solid DEA-carbamate (obtained from emulsion without Triton® X-100).
Figure A.3. 13
C NMR spectrum of DEA-carbamate crystals (obtained in the absence of Triton® X-100) taken in DMSO-d6, also
revealing the traces of trapped [hmim][Tf2N] in the crystal.
Figure A.4. 13
C NMR of pure DEA in DMSO-d6.
Figure A.5. 13
C NMR of pure [hmim][Tf2N] in DMSO-d6.
Figure A.6. FTIR spectrum of solid DEA-carbamate (obtained from emulsion without Triton® X-100).
Figure A.7. FTIR spectrum of the liquid separated from crystalline product (showing physical solubility of CO2).
Figure A.8. FTIR spectra of a) solid DEA-carbamate, b) pure DEA, and c) pure [hmim][Tf2N].
Appendix A
134
Table A.1. Bond lengths and angles in CO2-captured product (DEA-carbamate).
Bond lengths [Å]
O(1)-C(1) 1.2695(18) C(4)-C(5) 1.523(2)
O(2)-C(1) 1.2764(17) O(5)-C(6) 1.430(2)
O(3)-C(3) 1.4155(18) O(6)-C(9) 1.415(2)
O(4)-C(5) 1.4274(19) N(2)-C(7) 1.4996(19)
N(1)-C(1) 1.3730(19) N(2)-C(8) 1.502(2)
N(1)-C(2) 1.4596(18) C(6)-C(7) 1.515(2)
N(1)-C(4) 1.4633(18) C(8)-C(9) 1.505(2)
C(2)-C(3) 1.528(2)
Bond angles [°]
C(1)-N(1)-C(2) 122.81(12) N(1)-C(4)-C(5) 113.06(12)
C(1)-N(1)-C(4) 119.42(11) O(4)-C(5)-C(4) 110.60(12)
C(2)-N(1)-C(4) 117.44(11) C(7)-N(2)-C(8) 112.87(11)
O(1)-C(1)-O(2) 123.18(13) O(5)-C(6)-C(7) 111.69(13)
O(1)-C(1)-N(1) 119.33(13) N(2)-C(7)-C(6) 111.58(12)
O(2)-C(1)-N(1) 117.49(12) N(2)-C(8)-C(9) 112.06(12)
N(1)-C(2)-C(3) 112.71(12) O(6)-C(9)-C(8) 112.36(14)
O(3)-C(3)-C(2) 111.23(13)
Appendix A
135
Figure A.1. Packing diagram of the C9H22N2O6 compound (DEA-carbamate). Hydrogen atoms not participating in hydrogen bonding
were omitted for clarity.
Appendix A
136
Figure A.2. Solid-state 13
C NMR of solid DEA-carbamate (obtained from emulsion without Triton® X-100).
PPM 180.0 160.0 140.0 120.0 100.0 80.0 60.0 40.0 20.0
163.5
572
61.1
500
60.7
606
58.9
831
54.3
603
53.1
910
51.2
525
Appendix A
137
Figure A.3. 13
C NMR spectrum of DEA-carbamate crystals (obtained in the absence of Triton® X-100) taken in DMSO-d6, also
revealing the traces of trapped [hmim][Tf2N] in the crystal.
PPM 180.0 160.0 140.0 120.0 100.0 80.0 60.0 40.0 20.0
162.5
700
61.3
468
58.5
883
51.3
312
50.8
435
Appendix A
138
Figure A.4. 13
C NMR of pure DEA in DMSO-d6.
PPM 180.0 160.0 140.0 120.0 100.0 80.0 60.0 40.0 20.0
60.7
803
52.1
266
Appendix A
139
Figure A.5. 13
C NMR of pure [hmim][Tf2N] in DMSO-d6.
PPM 180.0 160.0 140.0 120.0 100.0 80.0 60.0 40.0 20.0
137.0
965
124.9
287
124.1
225
122.7
736
121.7
290
118.5
292
115.3
295
49.4
685
36.1
736
31.1
321
29.9
640
25.7
317
22.4
161
14.1
045
Appendix A
140
Figure A.6. FTIR spectrum of solid DEA-carbamate (obtained from emulsion without Triton® X-100).
ATR-FTIR
DEA-carbamate
Appendix A
141
Figure A.7. FTIR spectrum of the liquid separated from crystalline product (showing physical solubility of CO2).
ATR-FTIR
[hmim][Tf2N] + CO2
Appendix A
142
Figure A.8. FTIR spectra of a) solid DEA-carbamate, b) pure DEA, and c) pure [hmim][Tf2N].
a
b
c
Appendix B
143
Supporting Information (Chapter 3)
Figure B.1. 13
C NMR spectrum of MEA-carbamate precipitate (strong signal at 162.46 ppm) taken in DMSO-d6, also revealing traces
of [bmim][BF4].
Figure B.2. 13
C NMR spectrum of AMP-carbamate (weak signal at 162.57 ppm indicates its unstable nature) precipitate, along with
[bmim][BF4], taken in DMSO-d6.
Figure B.3. 13
C NMR spectrum of DEA-carbamate (strong signal at 163.03 ppm) precipitate taken in DMSO-d6, also showing
presence of [bmim][BF4].
Figure B.4. 13
C NMR of pure [bmim][BF4] in DMSO-d6.
Appendix B
144
CO2 capture by MEA (in amine-RTIL blend):
CO2 capture by AMP (in amine-RTIL blend):
CO2 capture by DEA (in amine-RTIL blend):
NH2OH + CO2
NHCOO-
OHNH3
+
OH+2
CH3OH
CH3
NH2
+ CO2 +CH3
OH
CH3
NHCOO-
CH3OH
CH3
NH3
+
2
NH
OH
OH
NCOO-
OH
OH
NH2
+
OH
OH
+ CO2 +2
Appendix B
145
Carbamate characterization by 13
C NMR
Liquid state 13
C NMR spectra were recorded on a Varian Inova Spectrometer at a frequency of 100 MHz with proton decoupling.
Samples were dissolved in DMSO-d6 (CND Isotopes, QC, Canada) and 600 scans were recorded for each spectrum.
Figure B.1.
13C NMR spectrum of MEA-carbamate precipitate (strong signal at 162.46 ppm) taken in DMSO-d6, also revealing traces of
[bmim][BF4].
MEA-Carbamate
162.4
6
NHCOO-
OH NH3
+ OH
Appendix B
146
Figure B.2. 13
C NMR spectrum of AMP-carbamate (weak signal1 at 162.57 ppm indicates its unstable nature) precipitate, along with
[bmim][BF4], taken in DMSO-d6.
1 Jo et al. Crystal structure and electronic properties of 2-amino-2-methyl-1-propanol (AMP) carbamate, Chem. Commun. 46 (2010) 9158–9160.
AMP-Carbamate
162.5
7
CH3OH
CH3
NHCOO-
CH3OH
CH3
NH3
+
Appendix B
147
Figure B.3. 13
C NMR spectrum of DEA-carbamate (strong signal at 163.03 ppm) precipitate taken in DMSO-d6, also showing
presence of [bmim][BF4].
DEA-Carbamate
163.0
3
NCOO-
OH
OH
NH2
+
OH
OH
Appendix B
148
Figure B.4. 13
C NMR of pure [bmim][BF4] in DMSO-d6.
[BMIM][BF4]
NN +
BF4
-
Appendix C
149
Supporting Information (Chapter 4)
Figure C.1. Carbamate crystals in RTIL medium, as seen under optical microscope: a) AMP-carbamate; b) DEA-carbamate.
Figure C.2. Structural units: a) AMP-carbamate; b) DEA-carbamate [C8].
Table C.1. Crystal data and structure refinement for AMP-carbamate
Table C.2. Bond lengths and angles for C9H22N2O4 (AMP-carbamate).
Figure C.3. Packing diagram of C9H22N2O4 compound (AMP-carbamate).
Figure C.4. FTIR spectrum of AMP-carbamate in the temperature range of 30-100 °C (regenerated AMP appeared in spectra taken at
80 °C and above).
Figure C.5. FTIR spectrum of DEA-carbamate in the temperature range of 30-80 °C (regenerated DEA appeared above 50 °C).
Appendix C
150
Single crystal X-ray diffraction studies:
To obtain carbamate crystals, CO2 was bubbled in respective amine-RTIL (AMP-[hmim][Tf2N] and DEA-[hmim][Tf2N]) blends.
Supernatant crystals were separated from the RTIL and washed thoroughly with acetone. The crystals were then dried and stored at
room temperature for characterization.
(a) (b)
Figure C.1. Carbamate crystals in RTIL medium, as seen under optical microscope: a) AMP-carbamate; b) DEA-carbamate.
Appendix C
151
Crystal structure determination:
Crystallographic data measurements were made at 100 K on a Bruker APEX II area detector diffractometer equipped with a source of
CuKα monochromatic radiation (λ = 1.54178 Å). APEX 2 and SAINT programs were used for retrieving cell parameters and data
collection [C1,C2]. Absorption corrections were performed using SADABS [C3]. The structure was solved and refined by full-matrix
least-squares against F2 using SHELXS-97 and SHELXL-97 programs [C4-C6]. Refinement of all non-hydrogen atoms was done
anisotropically. The hydrogen atoms were placed at geometrically idealized positions using a riding model. Further experimental
details are provided in Tables C.1-C.2. Crystal structure data for AMP-carbamate has also been reported previously by Jo et al., 2010
[C7]. Single crystal X-ray diffraction analysis of DEA-carbamate has been discussed in our previous work [C8].
Appendix C
152
Figure C.2. Structural units: a) AMP-carbamate; b) DEA-carbamate [C8].
(a) C9H22N2O4 (AMP-Carbamate) (b) C9H22N2O6 (DEA-Carbamate)
Appendix C
153
Table C.1. Crystal data and structure refinement for AMP-carbamate Empirical formula C9H22N2O4
Formula weight 222.29
Temperature 100K
Wavelength 1.54178 Å
Crystal system Monoclinic
Space group P21/n
Unit cell dimensions a = 6.02881(12) Å α = 90°
b = 9.88517(19) Å β = 97.8757(8)°
c = 20.4701(4) Å γ = 90°
Volume 1208.43(4)Å3
Z 4
Density (calculated) 1.222 g/cm3
Absorption coefficient 0.790 mm-1
F(000) 488
Crystal size 0.12 x 0.08 x 0.04 mm
Theta range for data collection 4.36 to 70.86°
Index ranges -6≤h≤7, -12≤k≤12, -25≤l≤24
Reflections collected 23054
Independent reflections 2270 [Rint = 0.025]
Refinement method Full-matrix least-squares on F2
Data / restraints / parameters 2270 / 0 / 164
Goodness-of-fit on F2 1.030
Final R indices [I>2sigma(I)] R1 = 0.0328, wR2 = 0.0864
R indices (all data) R1 = 0.0336, wR2 = 0.0872
Appendix C
154
Table C.2. Bond lengths and angles for C9H22N2O4 (AMP-carbamate).
Bond lengths [Å]
O(1)-C(1) 1.4322(12) C(2)-C(4) 1.5338(14)
O(2)-C(3) 1.2800(13) O(4)-C(6) 1.4129(13)
O(3)-C(3) 1.2688(13) N(2)-C(7) 1.5022(12)
N(1)-C(3) 1.3625(13) C(6)-C(7) 1.5306(14)
N(1)-C(2) 1.4785(13) C(7)-C(8) 1.5220(14)
C(1)-C(2) 1.5343(15) C(7)-C(9) 1.5227(14)
C(2)-C(5) 1.5288(14)
Bond angles [°]
C(3)-N(1)-C(2) 127.14(9) O(3)-C(3)-N(1) 117.45(9)
O(1)-C(1)-C(2) 115.03(8) O(2)-C(3)-N(1) 120.08(9)
N(1)-C(2)-C(5) 110.59(9) O(4)-C(6)-C(7) 113.41(8)
N(1)-C(2)-C(4) 105.89(8) N(2)-C(7)-C(8) 107.96(8)
C(5)-C(2)-C(4) 110.06(9) N(2)-C(7)-C(9) 107.96(8)
N(1)-C(2)-C(1) 112.45(8) C(8)-C(7)-C(9) 111.61(9)
C(5)-C(2)-C(1) 111.37(9) N(2)-C(7)-C(6) 108.25(8)
C(4)-C(2)-C(1) 106.24(9) C(8)-C(7)-C(6) 111.43(9)
O(3)-C(3)-O(2) 122.45(9) C(9)-C(7)-C(6) 109.51(9)
Appendix C
155
Figure C.3. Packing diagram of C9H22N2O4 compound (AMP-carbamate).
Appendix C
156
Figure C.4. FTIR spectrum of AMP-carbamate in the temperature range of 30-100 °C (regenerated AMP appeared in spectra taken at
80 °C and above).
Appendix C
157
Figure C.5. FTIR spectrum of DEA-carbamate in the temperature range of 30-80 °C (regenerated DEA appeared above 50 °C).
Appendix C
158
References:
[C1] APEX2, Bruker Molecular Analysis Research Tool. Bruker AXS Inc., Madison, WI, 2009.
[C2] SAINT, Release 7.34A; Integration Software for Single Crystal Data. Bruker AXS Inc., Madison, WI, 2006.
[C3] G.M. Sheldrick, SADABS; Bruker Area Detector Absorption Corrections. Bruker AXS Inc., Madison, WI, 2008.
[C4] G.M. Sheldrick, A short history of SHELX, Acta Cryst. A 64 (2008) 112-122.
[C5] SHELXTL, version 6.12; Bruker Analytical X-ray Systems Inc., Madison, WI, 2001.
[C6] XPREP, Version 2008/2; X-Ray Data Preparation and Reciprocal Space Exploration Program. Bruker AXS Inc., Madison, WI,
2008.
[C7] E. Jo, Y.H. Jhon, S.B. Choi, J.-G. Shim, J.-H. Kim, J.H. Lee, I.-Y. Lee, K.-R. Jang, J. Kim, Crystal structure and electronic
properties of 2-amino-2-methyl-1-propanol (AMP) carbamate, Chem. Commun. 46 (2010) 9158-9160.
[C8] M. Hasib-ur-Rahman, M. Siaj, F. Larachi, CO2 capture in alkanolamine/room-temperature ionic liquid emulsions: A viable
approach with carbamate crystallization and curbed corrosion behavior, Int. J. Greenh. Gas Control 6 (2012) 246-252.
Appendix D
159
Corrosion in amine systems – a review*
Irrespective of consistent and dominant usage of aqueous amine based processes in acid gas capture facilities since 1930s, there is
constant concern over a number of operational snags including, but not limited to, corrosion. Various physical/chemical factors like
process temperature, amine type/concentration, metallurgy, CO2 concentration, other gaseous impurities, gas loading, suspended
particles, and heat stable salts, play their respective role in intensifying the corrosion occurrence that also favours solvent degradation.
This obligates the use of additives that not only supplement the cost but also pose a risk to the environment, as typically heavy metals
such as arsenic, vanadium etc. constitute the more efficient corrosion inhibitors.
What else then?
Replacing water with more stable room-temperature ionic liquid (RTIL) in amine based systems might be an optimistically workable
option as it renders three benefits: excellent corrosion control, stoichiometric maximum gas loading by overcoming equilibrium
limitations through carbamate precipitation, and separation of CO2-captured product in the form of carbamates.
This strategy might also promise additive-free capture fluids. Moreover, easy separation of solid carbamate can offer cost-effective
regeneration. As comprehensive scrutiny is still needed in this regard, the limited work has shown some good prospects of
alkanolamine/RTIL mixtures as more optimal successors of aqueous amines for CO2 capture.
* M. Hasib-ur-Rahman, F. Larachi, Carbon Capture Journal, Sept - Oct 2012, 22-24.
Appendix D
160
Perspective
Amine-based chemical solvents have been in practice for over half a century in the oil and gas processing industry and are being
considered as one of the best potential candidates for CO2 capture from flue gases. However, this cannot be a trouble-free technology
regarding post-combustion capture in particular as flue gases contain particulate matter as well as acid gas impurities other than CO2
(like H2S/SO2) that ought to be removed separately [1,2]. Besides, corrosion of equipment and amine degradation further adds to
process downsides.
All aqueous amine-based CO2 capture installations are susceptible to corrosion that not only adds to process costs but also raises
concerns about the safety of personnel and environment. A number of factors like higher CO2 loading, increased amine concentration,
elevated process temperature, as well as presence of oxygen greatly intensifies the corrosion of metal (Figure 1). Moreover, presence
of suspended solid particles and amine degradation products/heat stable salts also causes to augment corrosion phenomenon. The
process equipment typically vulnerable to corrosion includes absorber, amine exchanger, regenerator, and pumps [3,4].
In aqueous alkanolamine-CO2 systems corrosion is the result of anodic (iron dissolution) and cathodic (reduction of oxidizers present
in the solution) electrochemical reactions on metal surfaces. In CO2-loaded aqueous amines, iron dissolution is induced by various
oxidizing species such as hydrogen/bicarbonate ions, protonated amine, carbamate ions, and undissociated water [4]. The most
significant redox reactions arising in this regard are:
2
2
2
3 3 2
2 2
2
2 2
2 2 2
2 2 2
Fe Fe e
H e H
HCO e CO H
H O e OH H
Appendix D
161
Figure D.1. Effect of various parameters on the corrosion rate of carbon steel C1020 in aqueous MEA (basal conditions: MEA conc. 5
kmol/m3; gas loading 0.4 mol CO2/mol MEA; 80 °C temperature) [4].
In case of either flue gas or raw natural gas, CO2 generally occurs in conjunction with some other acidic impurities like SO2, H2S that
also help accelerate wear and tear of the metallic tools. For instance, the SO2 amount in the flue gas causes a proportionate increase in
iron dissolution through the formation of hydrogen ions as shown below [5].
Appendix D
162
2 2 3
2
3 3
2122 2 2 42
SO H O H HSO
HSO H SO
SO O H O H SO
The H+ ions serve to abstract electrons from metallic iron resulting in oxidative decay of the equipment:
2
22Fe H Fe H
Amine degradation products have also been found to increase the corrosion rate and same is the behavior of corrosion products toward
amine degradation. Degradation occurrence not only depletes the active CO2 capturing material but the resulting species also speed up
the corrosion rate by introducing additional oxidants. In fact corrosion and degradation phenomena are closely interrelated. For optimal
functionality, perpetual removal of contaminants (degradation/corrosion products, particulate matter, etc.) from the chemical solvent is
required [2,6,7].
Corrosion control
Various approaches can be practiced to prevent corrosion to avoid severe operational problems in amine treating units. These may
include process-specific equipment metallurgy/design, removal of contaminants, and use of corrosion inhibitors. The last option has
been accomplished in industry quite effectively. A number of corrosion inhibitors, based on arsenic, antimony, vanadium, copper (like
NaVO3, CuCO3) are being used in order to control and prevent corrosion that not only adds to the capital cost but most of these are
toxic and hazardous to life as well. The more strict regulations in the case of toxic/hazardous substances in the very near future may
limit the use of such compounds due to high disposal costs [8,9].
Appendix D
163
Alternative workable route
Replacing the problematic aqueous phase (chiefly responsible for corrosion occurrence) with some apposite solvent such as non-
corrosive room-temperature ionic liquids under gas capture conditions might be a viable option at least as a near-term solution.
Alkanolamine/room-temperature ionic liquid blends
Imidazolium based room-temperature ionic liquids (RTILs) are thermally stable, virtually non-volatile, and generally non-corrosive.
RTILs being of tunable nature, because of the availability of manifold ion-pair combinations, can be tailored by choice in accordance
with the individual process requirements and hence can be used as a replacement of water in alkanolamine based CO2 capture
processes. These can significantly suppress corrosion phenomena when combined with primary/secondary alkanolamines [10,11].
Moreover, such novel schemes also offer some momentous benefits regarding CO2 separation methodology [10-13]:
Carbamate precipitation/crystallization
Stoichiometric maximum gas loading by avoiding equilibrium limitations contrary to what is experienced in aqueous amine
based systems
Enabling easy separation of solid carbamate thus promising cost-effective regeneration
Appendix D
164
Figure D.2. Schematic concept of CO2 scrubbing by amine-RTIL blends.
Larachi‟s research group at Laval University has studied the corrosion behaviour of carbon steel 1020 in alkanolamines blended with
hydrophobic or hydrophilic ionic liquids [10,11]. Linear polarization resistance (LPR) measurements followed by Tafel extrapolation
method was employed using a Bio-Logic VSP potentiostat. Diethanolamine (DEA)/hydrophobic 1-hexyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide ionic liquid ([hmim][Tf2N]) combination appeared to better control corrosion occurrence even at
Appendix D
165
higher temperature. The results showed that at 60 °C, even in the presence of oxygen and moisture along with CO2, the corrosion rate
was negligibly small (<1 mpy) as shown in Figure 3.
To know the effect of ionic liquid‟s hydrophobic/hydrophilic nature, three hydrophilic RTILs (1-butyl-3-methylimidazolium
tetrafluoroborate [bmim][BF4], 1-ethyl-3-methylimidazolium tetrafluoro borate [emim][BF4], 1-ethyl-3-methylimidazolium
trifluoromethanesulfonate [emim][Otf]) were studied in more detail. Effects of amine/RTIL type, water content, CO2 loading, O2
concentration in simulated flue gas, water content, as well as temperature were evaluated. At 25°C, the amine/RTIL blends showed
good corrosion control but at higher temperature (60°C) the carbon steel underwent a substantial amount of corrosion, however, it was
still lower up to about 70 % when compared to what was observed in aqueous monoethanolamine (Figure 4).
Appendix D
166
Figure D.3. Corrosion rate of steel 1020 under CO2+O2+H2O(vap.) atmosphere, a) Aqueous diethanolamine (15% w/w); b) Pure
[hmim][Tf2N]; c) Diethanolamine/RTIL emulsion (15% w/w).
Since the CO2-captured product (carbamate) moves away from the reaction phase as solid moieties, it is no longer involved in the
electrochemical corrosion reactions. Also, RTIL coating on metal surface barricades the access of any oxidants to the working
electrode facet. Furthermore, the absence of aqueous phase, that in combination with CO2/O2 provides the bulk share of oxidizing
species in the case of aqueous amine solvents, diminishes the chances of the occurrence of redox process. The results also
Appendix D
167
demonstrated that hydrophobic ionic liquids, compared to hydrophilic ones, could efficiently prevent metal deterioration at higher
temperatures and might offer more success in case if the whole bulk of gas capturing amine/RTIL fluid would be subjected to thermal
regeneration. This supremacy is probably due to its superior safeguarding through coating/adsorption on the metal surface and also
because of its repelling behaviour toward water species.
In spite of the above cited outcomes, prior to large scale applications, a significant amount of work is still required to exactly evaluate
the corrosion phenomenon under real regeneration conditions. Moreover, it is yet to be scrutinized if we can avoid amine degradation
by using this stratagem, and the impact of impurities/contaminants also needs appraisal.
Appendix D
168
Figure D.4. Various process conditions tested for amine/RTIL (hydrophilic) blends, a) Amine type; b) RTIL type; c) CO2 loading; d)
O2 conc. in flue gas; e) Water content in the fluid; f) Temperature effect.
Appendix D
169
References
[1] Perry et al. CO2 Capture Using Phase-Changing Sorbents, Energy Fuels 26 (2012) 2528-2538.
[2] http://www.mprservices.com/pdfs/corrosionenhancers.pdf
[3] I.R. Soosaiprakasam and A. Veawab, Corrosion and polarization behavior of carbon steel in MEA-based CO2 capture process, Int.
J. Greenh. Gas Control 2 (2008) 553-562.
[4] N. Kladkaew, R. Idem, P. Tontiwachwuthikul and C. Saiwan, Corrosion behaviour of carbon steel in the monoethanolamine-H2O-
CO2-O2-SO2 system: products, reaction pathways, and kinetics, Ind. Eng. Chem. Res. 48 (2009) 10169-10179.
[5] N. Kladkaew, R. Idem, P. Tontiwachwuthikul and C. Saiwan, Corrosion behaviour of carbon steel in the monoethanolamine-H2O-
CO2-O2-SO2 system, Ind. Eng. Chem. Res. 48 (2009) 8913-8919.
[6] W. Tanthapanichakoon and A. Veawab, Electrochemical Investigation on the Effect of Heat-stable Salts on Corrosion in CO2
Capture Plants Using Aqueous Solution of MEA, Ind. Eng. Chem. Res. 45 (2006) 2586-2593.
[7] S. Chi and G.T. Rochelle, Oxidative Degradation of Monoethanolamine, Ind. Eng. Chem. Res. 41 (2002) 4178-4186.
[8] A. Veawab, P. Tontiwachwuthikul and A. Chakma, Investigation of low-toxic organic corrosion inhibitors for CO2 separation
process using aqueous MEA solvent, Ind. Eng. Chem. Res. 40 (2001) 4771-4777.
[9] I.R. Soosaiprakasm and A. Veawab, Corrosion inhibition performance of copper carbonate in MEA-CO2 capture unit, Energy
Procedia 1 (2009) 225-229.
Appendix D
170
[10] M. Hasib-ur-Rahman, M. Siaj and F. Larachi, CO2 Capture in Alkanolamine/Room-Temperature Ionic Liquid Emulsions: A
Viable Approach with Carbamate Crystallization and Curbed Corrosion Behavior, Int. J. Greenhouse Gas Control 6 (2012) 246-252.
[11] M. Hasib-ur-Rahman, H. Bouteldja, P. Fongarland, M. Siaj and F. Larachi, Corrosion Behavior of Carbon Steel in
Alkanolamine/Room-Temperature Ionic Liquid Based CO2 Capture Systems, Ind. Eng. Chem. Res. 51 (2012) 8711-8718.
[12] D. Camper, J. E. Bara, D. L. Gin and R. D. Noble, Room-Temperature Ionic Liquid-Amine Solutions: Tunable Solvents for
Efficient and Reversible Capture of CO2. Ind. Eng. Chem. Res. 47 (2008) 8496-8498.
[13] Q. Huang, Y. Li, X. Jin, D. Zhao and G. Z. Chen, Chloride ion enhanced thermal stability of carbon dioxide captured by
monoethanolamine in hydroxyl imidazolium based ionic liquids. Energy Environ. Sci. 4 (2011) 2125-2133.