HÅVARD LIDAL NO9505210
NEI -NO--562
CARBON DIOXIDE REMOVAL IN GAS TREATING PROCESSES
TH UNIVERSITETET I TRONDHEIM NORCES TEKNISKE HØGSKOLE
DOKTOR INGENIØR AVHANDLING 1992:26 INSTITUTT FOR KJEMITEKNIKK TRONDHEIM
DISTRIBUTION OF THIS DOCUMENT IS'UNUMITED'f
smamrnm
CARBON DIOXIDE REMOVAL
IN
GAS TREATING PROCESSES
NJH.IRVKH 199!
by
Håvard Lidal
A Thesis Submitted for the Degree of
Dr. Ing.
The University of Trondheim
The Norwegian Institute of Technology
Department of Chemical Engineering
Trondheim, June 1992
MASTER DISTRIBUTION OF THIS DOCUMENT IS UNLIMITED
ADDENDUM
The cooperation of the industrial participants in this SPUNG project, Norsk Hydro a.s and Kværner Engineering A/S, is greatly appreciated.
I
ACKNOWLEDGEMENTS
I am most obliged to my supervisor Olav Erga for all his
professional and personal support. His encouragement, inspiring
personality, and wholehearted interest in the field of gas
treating have given me the backing I needed during this work.
I wish to express my sincere appreciation to Dag Eimer of Norsk
Hydro a.s. I learned a lot from discussions we had, and I enjoyed
working with him on various projects.
Thanks are due to Olav Juliussen of SINTEF for his technical
assistance with the laboratory equipment. I also wish to
acknowledge the contributions of A.R. Fossen-Helle, J. Bjørvik,
W.E. Olsen, M. Schneider, and M. Tørnqvist for performing parts
of the experiments.
Thanks also to all those representatives of the gas industry,
professors and staff members from our university and universities
and research establishments around the world, and other people
I had the opportunity to meet and have inspiring discussions
with. In particular, I would like to thank Orville C Sandall
from UCSB who accepted to serve on my dissertation committee, and
travel all the way from California to do this.
Above all, I would like to give special thanks ;,o God and my
family, especially my late father, my mother, and my brother.
I gratefully acknowledge the financial support of the Royal
Norwegian Council for Scientific and Industrial Research (NTNF),
given as a part of the SPUNG Programme (State R&D Programme for
Utilization of Natural Gas). The support of the Foundation for
Scientific and Industrial Research at the Norwegian Institute of
Technology (SINTEF), as well as grants received from NTHs Fond,
M.H. Lungreens Enkes Fond, and Lise og Arnfinn Hejes Fond, are
greatly appreciated.
i n
ABSTRACT
A semiempirical thermodynamic model which represents the
equilibrium partial pressure of C02 over aqueous solutions of
tertiary and sterically hindered amines, is presented. The model
has been used on the tertiary amine methyldiethanolamine (MDEA),
and on the sterically hindered amine 2-amino-2-methyl-1-propanol
(AMP). Measurements of pH as a function of C02 concentration play
an important role in the modelling procedure. The model is based
on the pH data, together with solubility measurements performed
in this work and also solubility data collected from the
literature. Solubility and pH measurements were made over a
temperature range of 25 to 70°C, and for amine concentrations of
3M AMP and 4 and 4.28M MDEA.
The model relates the equilibrium partial pressure of CO2 as a
function of the amine concentration, the C02 loading, and the
temperature. For MDEA solutions, the model covers the temperature
interval of 25 to 140°C, and can be used for C02 loadings between
0.001 and 1 mol C02/mol MDEA, and at C02 partial pressures
between 0.00001 and 50 atm. The model is tested against
experimental data from several literature references with amine
concentrations ranging from 1.69 to 4.28M, and it is found to
predict the experimental data very well. While the presented
model covers both absorption and desorption conditions for MDEA
solutions, the application range is restricted to absorption
conditions for AMP solutions.
The technique of utilizing measured pH data in the modelling of
vapor-liquid equilibrium, distinguishes the present model from
equilibrium models found in the literature. Establishing accurate
relations for pH as a function of the C02 loading and the
IV
temperature, constitute the backbone on which the model is based.
The solubility of C02 has been measured over a temperature range
of 30 to 70°C in mixed nonaqueous solutions of glycols and
alkanolamines. The following systems have been studied:
Triethyleneglycol (TEG) with either monoethanolamine (MEA) or
diethanolamine (DEA), and diethyleneglycol (DEG) with MEA.
Measurements were made with amine contents of 5, 10, and
13.6mol%. The solubility in these mixed solvents is compared with
other mixed solvents and also with aqueous amine solutions. The
effect of temperature and amine concentration on solubility is
also discussed.
To be able to estimate the CO2 partial pressure at temperatures
above 70°C, a vapor-liquid equilibrium model is developed for the
TEG/MEA-system. The model, which is in many aspects similar to
the model developed for the aqueous system, shows satisfactory
agreement with the available experimental data.
The rates of C02 absorption into mixed solvents have been
measured using a string-of-discs experimental set-up. These
experiments were undertaken on five solvents with and without the
addition of 5mol% MEA. The following solvents were investigated:
N-methyl-pyrrolidone, ethanol, diethyleneglycol monomethylether,
TEG, and water. Due to problems with temperature rise, only
approximate data have been obtained.
To improve our laboratory facilities, two new experimental set
ups have been designed and built. These are an apparatus for
solubility measurements at temperatures above 70°C, and a one-
sphere apparatus for determinations of reaction kinetics. Both
sets of equipment are described in this thesis.
v
TABLE OF CONTENTS
Acknowledgements i i i Abs t rac t iv List of tables ix List of figures xi
Chapter One Introduction 1 1 .1 Acid Gas Removal Technologies 1 1 .2 Alkanolaraine Solutions 3 1 .3 Scope of the Work 7
Chapter Two Literature Review 9 2.1 VLE Data in Gas Treating Processes 9
2.1.1 VLE Measurements in Aqueous Alkanolamine Solutions 9
2.1.2 VLE Modelling in Aqueous Alkanolamine Solutions 10
2.1.3 VLE Measurements in Mixed Nonaqueous Solvents 14 2.1.4 VLE Measurements in Pure Physical Solvents... 15 2.1.5 VLE Modelling Techniques in Physical Solvents 16
2.2 Chemistry of C02 - Amine Systems 17 2.2.1 Introduction 17 2.2.2 Reactions between C02 and Amines in
Aqueous Solutions 18 2.2.3 Reaction Kinetics between C02 and
Aqueous MDEA 22 2.2.4 Reaction Kinetics between C02 and Aqueous AMP 24 2.2.5 Reaction Kinetics in Nonaqueous Solutions.... 25 2.2.6 Experimental Equipment for Kinetic
Determinations 27
Chapter Three Experimental 29 3 .1 Vapor-Liquid Equilibrium Measurements 29
3.1.1 Equilibrium Equipment for Temperatures up to 70°C 29
3 . 1 . 2 A New Equipment fo r T e m p e r a t u r e s up t o 120°C. 30 3 .2 pH Measurements 31 3 .3 K i n e t i c Measurements 32
3 . 3 . 1 S t r i n g - o f - d i s c s Column 32 3 . 3 . 2 O n e - s p h e r e A p p a r a t u s 33
3 . 4 Chemica l s and Gases 34 3 .5 L i q u i d A n a l y s i s 35
3 . 5 . 1 C02 C o n c e n t r a t i o n 35 3 . 5 . 2 Amine C o n c e n t r a t i o n 36
3 . 6 Gas A n a l y s i s 36
vi
Chapter Four Experimental Results 37 4.1 Vapor-Liquid Equilibrium Measurements 37
4.1.1 C02 Solubility in Aqueous MDEA Solutions 37 4.1.2 C02 Solubility in Aqueous AMP Solutions 40 4.1.3 C02 Solubility in Nonaqueous Amine Solutions. 41
4.2 pH Measurements 47 4.2.1 Aqueous MDEA Solutions 47 4.2.2 Aqueous AMP Solutions 48
4.3 Kinetic Measurements 49
Chapter Five A Model for Equilibrium Solubility of C02 in Aqueous Solutions of the Tertiary Amine MDEA 52
5 .1 Introduction 52 5.2 C02 Equilibrium Model for Aqueous 4M MDEA 54
5.2.1 Approximations 54 5.2.2 The Basic Model 55 5.2.3 A Correlation for pH 55 5.2.4 A Correlation for logK 58 5.2.5 A Preliminary Final Model 59 5.2.6 Comparison with Experimental Equilibrium Data 59
5.3 Extended Equilibrium Model, Valid for Aqueous Solutions with 1-4.5M MDEA at Temperatures between 25 and 1 40 °C 60 5.3.1 Introducing VLE Data from the Literature 60 5.3.2 A New Correlation for the Parameter K 60 5.3.3 The Final Model 61 5.3.4 Comparison with Experimental Equilibrium Data 62
5 .4 Accuracy of the Model 71 5 .5 Conclusions 71
Chapter Six A Model for Equilibrium Solubility of C02
in an Aqueous Solution of the Sterically Hindered Amine AMP 72
6 .1 Introduction 72 6 .2 The Equilibrium Model for C02 75
6.2.1 Approximations 75 6.2.2 The Basic Model 75 6.2.3 A Correlation for pH 76 6.2.4 A Correlation for logK 78 6.2.5 The Final Model 79 6.2.6 Comparison with Experimental Equilibrium Data 79 6.2.7 Limitations 80
6 .3 Conclusions 80
vii
Chapter Seven Vapor-Liquid Equilibria of Mixed Nonaqueous Solvents 81
7.1 Equlibrium Solubility Model for C02 in TEG/MEA Solutions 81 7.1.1 Background 81 7.1.2 Modelling Procedure 82 7.1.3 Comparison with Experimental Equilibrium Data 85
7.2 Comparison with Aqueous Amine Solutions 86 7.3 Comparison with other Mixed Solvents 91 7.4 Comparison with Pure Physical Solvents 92
Chapter Eight Conclusions and Recommendations 93 8.1 Conclusions 93 8. 2 Recommendations 94
Nomenclature 96
References 98
Appendix A Tabulated Data of C02 Solubility in Aqueous Systems 108
Appendix B Tabulated Data of C0 2 Solubility in
Nonaqueous Systems 110
Appendix C Tabulated pH Data for Aqueous Systems 116
Appendix D Tabulated Results of Kinetic Measurements.... 118
Appendix E HP-42S Program for Calculation of Equilibrium Partial Pressure of C02 over Aqueous MDEA.... 123
viii
LIST OF TABLES
Table 1 Solubility of C02 in aqueous solutions of 4.00M MDEA at 30, 45, and 60°C 108
Table 2 Solubility of C02 in aqueous solutions of 4.28M MDEA at 25, 40, and 70°C 109
Table 3 Solubility of C02 in aqueous solutions of 3.00M AMP at 40 and 50°C 109
Table 4 Solubility of C02 in solutions of TEG and 5mol% MEA at 30, 50, and 70°C 110
Table 5 Solubility of C02 in solutions of TEG and 10mol% MEA at 30, 50, and 70°C 111
Table 6 Solubility of C02 in solutions of TEG and 5mol% DEA at 30, 50, and 70°C 112
Table 7 Solubility of C02 in solutions of TEG and 10mol% DEA at 30 and 50°C 113
Table 8 Solubility of C02 in solutions of TEG and 13.6mol% DEA at 30, 50, and 70°C 114
Table 9 Solubility of C02 in solutions of DEG and 5mol% MEA at 40°C 115
Table 10 Solubility of C02 in solutions of DEG and 10mol% MEA at 40°C 115
Table 11 pH values as a function of C02 loading in aqueous solutions of 4.00M MDEA at 30, 40, 50, and 60°C. 116
Table 12 pH values as a function of C0 2 loading in aqueous
solutions of 3.00M AMP at 20, 30, 40, and 50°C... 117
Table 13 Rate of absorption of C02 in water at 20°C 118
Table 14 Rate of absorption of C02 in a solution of water and 5mol% MEA at 20°C 118
Table 15 Rate of absorption of C02 in n-methyl-pyrrolidone at 20°C 119
Table 16 Rate of absorption of C02 in a solution of n-methyl-pyrrolidone and 5mol% MEA at 20°C 119
IX
Table 17 Rate of absorption of C02 in ethanol at 20°C 120
Table 18 Rate of absorption of C02 in a solution of ethanol and 5mol% MEA at 20°C 120
Table 19 Rate of absorption of C02 in triethyleneglycol at 20CC 121
Table 20 Rate of absorption of C02 in a solution of triethyleneglycol and 5mol% MEA 121
Table 21 Rate of absorption of C02 in diethyleneglycol monomethylether at 20°C 122
Table 22 Rate of absorption of C02 in a solution of diethyleneglycol monomethylether and 5mol% MEA at 20°C 122
x
LIST OF FIGURES
Figure 2.1
Figure 3.1
Figure 3.2
Figure 3.3
Figure 3.4
Figure 4.1
Figure 4.2
Figure 4.3
Figure 4.4
Figure 4.5
Figure 4.6
Figure 4.7
Figure 4.8
Figure 4.9
Figure 4.10
Figure 4.11
Molecular structure of amines used in acid gas removal processes 19
Gas-liquid equilibrium equipment 30
New gas-liquid equilibrium equipment, capable of measuring solubilities at temperatures encountered in desorption units.... 31
String-of-discs absorber 32
Schematic of operation of string-of-discs and one-sphere apparatus 34
Solubility of C02 in aqueous 4.00M MDEA solutions at 30, 45, and 60°C 38
Solubility of C02 in aqueous 4.28M MDEA solutions at 25, 40, and 70°C, compared with literature data 39
Solubility of C02 in aqueous 3.00M AMP solutions, compared with literature data.... 40
Solubility of C02 in TEG solutions containing 5mol% MEA at 30, 50, and 70°C.... 41
Solubility of C02 in TEG solutions containing 10mol% MEA at 30, 50, and 70°C... 42
Solubility of C02 in TEG solutions containing 5mol% DEA at 30, 50, and 70°C... 43
Solubility of C02 in TEG solutions containing 10mol% DEA at 30 and 50°C 44
Solubility of C02 in TEG solutions containing 13.6mol% DEA at 30, 50, and 70°C. 45
Solubility of C02 in DEG solutions containing 5mol% MEA and 1 Omol% MEA at 40°C 46
Experimental pH data for aqueous 4.00M MDEA solutions at 30, 40, 50, and 60°C 47
Experimental pH data for aqueous 3.00M AMP solutions at 20, 30, 40, and 50°C 48
XI
Figure 4.12
Figure 4.13
Figure 5.1
Figure 5.2
Figure 5.3
Figure 5.4
Figure 5.5
Figure 5.6
Figure 5.7
Figure 5.8
Figure 5.9
Rate of absorption of C02 in physical solvents as a function of wetting rate at 20°C. 50
Rate of absorption of C02 in physical solvents containing 5mol% MEA as a function of wetting rate at 20 °C 51
pKp' as a function of the temperature for aqueous 4.00M MDEA solution 57
logK as a function of the temperature for aqueous 4.00M MDEA solution 58
Comparison of the present model with experimental data from the literature on the system of 4.28M MDEA aqueous solution at 25, 40, 70, 100, and 120°C 63
Comparison of the present model with experimental data from the literature on the system of 4.28M MDEA aqueous solution at 140°C 64
Comparison of the present model with present experimental data and data taken from the literature on the system of 4.28M MDEA aqueous solution at 40°C 65
Comparison of the present model with experimental data from the literature on the system of 2.00M MDEA aqueous solution at 25, 40, 70, 100, and 120°C 66
Comparison of the present model with experimental data from the literature on the system of 2.00M MDEA aqueous solution at 40°C 67
Comparison of the present model with experimental data from the literature on the system of 3.04M MDEA aqueous solution at 40 and 100°C... 68
Comparison of the present model with experimental data from the literature on the system of 1.69M MDEA aqueous solution at 100°C 69
xii
Figure 5.10
Figure 6.1
Figure 6.2
Figure 7.1
Figure 7.2
Figure 7.3
Figure 7.4
Figure 7.5
Figure 7.6
Figure 7.7
Figure 7.8
Comparison of the present model with present experimental data for aqueous solutions of 4.00M MDEA at 30°C 70
pKp* as a function of 1/T for aqueous 3.00M AMP solution 77
logK as a function of C02 loading for aqueous 3. 0OM AMP solution 78
Equilibrium partial pressure of CO2 presented as a function of 1000/T for eight different C02 loadings in aqueous solutions Of 4.28M MDEA 83
Equilibrium partial pressure of CO2 presented as a function of 1000/T for five different C02 loadings in a solution of TEG and 10mol% MEA 84
Comparison of the present model with present experimental data for a solution of TEG and 10mol% MEA at 30, 50, and 70°C, and predicted equilibrium curves for 100 and 150°C. 86
Comparison of equilibrium curves at 40°C for three different solvents, all containing 5mol% MEA 88
Comparison of equilibrium curves for the TEG/DEA system at different amine concentrations at 30°C 89
Comparison of equilibrium curves at 50°C for TEG solutions containing 10mol% MEA and 10mol% DEA 90
Present C02 solubility data in a mixed TEG/MEA solution compared with the solubility in NMP/MEA solutions at 50°C 91
C02 solubility data for 5mol% and 10mol% MEA in TEG, compared with the solubility in pure TEG 92
xiii
Chapter One
Introduction
1 .1 ACID GAS REMOVAL TECHNOLOGIES
Acid gases such as carbon dioxide (C02), hydrogen sulfide (H2S),
and sulfur dioxide (S02) are removed from a variety of gas
streams, including natural gas, flue gas, synthesis gas, and
refinery gases. Acid gas treating generally refers to removal of
C02 and H2S, while the removal of S02 is often denoted r"lue gas
desulfurization, although the technology used is often very
similar. Removal of organic sulfur compounds such as carbonyl
sulfide (COS), carbon disulfide (CS2), mercaptans (RSH),
thiophene, and other impurities present at low concentration
levels (HCN, NH3, S03), are often required as well.
Kohl and Riesenfeld (1985) divides all gas purification processes
into three categories: absorption into liquid, adsorption on a
solid, and chemical conversion to another compound. In addition
both cryogenic and membrane technology can be applied favorably
in certain cases. Absorption into a liquid is the most used
method (Astarita et al. (1983)), and is the method studied in
this thesis. The liquid solution can consist of either a physical
solvent, a chemical solvent (or a blend of chemical solvents) in
water, or a mixed solvent containing both a chemical active and
a nonaqueous physical solvent.
The best method for a certain application is decided by
parameters such as feed gas composition, pressure and quantity
of gas treated, as well as the cleanup target. Since the process
1
to be chosen, will be the one that shows the best economics and
the most reliable operation, it is important to have at hand
design data for the processes. In the case of absorption
processes, models describing the gas-liquid equilibria are
important tools in the process design.
The acid gas content in feed gases to treating units can range
from less than 1% to well above 50%. The specification of acid
gas in treated gas varies markedly from application to
application. For example, according to Astarita et al. (1983),
the pipeline specification for natural gas is maximum 4 ppm H2S
and 1% C02. For natural gas to LNG plants the C02 content is
usually limited to 50 ppm, and in ammonia manufacturing, the C02
impurity of the feed gas must be reduced to 10 ppm.
As one can see, the range over which the feed gas compositions
and the desired treated gas specifications varies, is quite
large. The capability to remove acid gases at these different
levels, depends highly on the process chosen. For example, the
pure physical solvents are well suited for bulk C02 removal when
the inlet partial pressure of C02 is relatively high, above
approximately 7 atm according to Astarita et al. (1983), while
at the same time the C02 specification in the treated gas is
quite loose. If deep acid gas removal is required, the addition
of an alkanolamine may help. Aqueous alkanolamine solutions are
often used when the partial pressure of C02 in feed gas is
relatively low and C02 removal down to ppm levels is required.
The use of alkanolamines is discussed in some detail in the next
section. Other useful chemical solvents for certain applications
are aqueous solutions of salts of amino acids as well as promoted
hot carbonate solutions.
Membrane technology will have its potential for treating of high
2
pressure gases with high levels of acid gas (Funk and Li (1989)).
For small acid gas removal units, savings in capital and
operating costs might be expected. However, to minimize
hydrocarbon losses, membranes with high C02/CH4 selectivity must
be developed or complex recirculation schemes must be used. Work
has also been done on a laboratory scale to use facilitated gel
membranes containing amines to separate hydrocarbons and acid
gases (Pellegrino et al. (1989) and Chakma (1992)).
Pressure swing adsorption processes can be competitive with
absorption in small plants (Astarita et al. (1983)). Adsorption
are suited for trace removal of acid gases.
1 .2 ALKANOLAMINE SOLUTIONS
Alkanolamines are the most used chemical active agent in acid gas
removal processes (Astarita et al. (1983)). Since Bottoms (1930)
introduced the amines to "separate acidic gases", and recommended
the tertiary triethanolamine (TEA) because of its higher boiling
point, several new and more suited amines have become
commercially available. Among the most important ones are
monoethanolamine (MEA), diethanolamine (DEA), diisopropanolamine
(DIPA), B, B'-hydroxy-aminoethylether (DGA, also known as
diglycolamine), and methyldiethanolamine (MDEA).
The tertiary amine MDEA has come to extensive use quite recently
for a number of gas treating applications. MDEA is the major
constituent in solvent processes offered by Dow Chemical Co.
(Gas/Spec solvents), Union Carbide (Ucarsol solvents), Texaco
Chemical Co. (Textreat solvents), and BASF (Activated MDEA).
These are proprietary formulated solvents containing inhibitors,
activators, and other additives. MDEA solutions exhibit large
3
acid gas capacity as well as easy regenerability.
The use of corrosion inhibitors have made it possible to increase
the amine concentration in the solutions markedly, and thereby
reduce the solvent circulation rate, giving lower operating and
capital costs. For example, aqueous MEA solutions can now be used
in concentrations up to 5M, compared to typically 3M previously
(Astarita et al. (1983)). According to Niswander et al. (1992),
the last generation of MDEA based solutions have diminished
their corrosiveness with a factor of 10, compared to the first
generation of MDEA solvents.
In recent years a new class of amines has been introduced: the
sterically hindered amines. Sartori and Savage (1983) define a
sterically hindered amine to be a primary amine where the amino
group is attached to a tertiary carbon atom, or a secondary amine
in which the amino group is attached to a secondary or a tertiary
carbon atom. Examples are 2-amino-2-methyl-1-propanol (AMP), 2-
(tert-butylamino) ethanol (TBE), and 2-piperidine ethanol (PE).
Due to the bulky substituent attached to the amino group, a
strong bonding of CO2 to the nitrogen atom is prevented, and the
result is a low tendency to form carbamates. As we shall see in
Chapter 6, an improved thermodynamic capacity exceeding 0.5 mol
C02/n\ol amine can be expected, at favorable absorption rates.
Besides, the sterically hindered amines are well suited as
promoters for the hot carbonate process (Say et al. (1984)). They
are also suited for selective removal of H2S when CO2 is present,
as an alternative to tertiary amines.
By using blends of amines, one can make use of each amine' s
attractive properties. For example, the large capacity and easy
stripping of an MDEA solution can be combined with an MEA
solution's ability to produce high purity sweet gas (Chakravarty
4
(1985)). This opens for an interesting possibility of
tailormaking blended amine solutions to meet specific acid gas
removal requirements.
Evidently, there is a great need for more fundamental research
into these "new" chemical solvents. Rochelle (1991) suggests that
further studies should be undertaken to obtain solubility and
kinetic data for both sterically hindered amines and MDEA based
solutions mixed with primary and secondary amines.
Alkanolamines are also used in nonaqueous solutions. Savings due
to easier regeneration can be obtained. An example of this is the
Sulfinol process using a mixture of DIPA, sulfolane
(tetrahydrothiophene dioxide), and water. This process has shown
capability of removing carbonyl sulfide (COS) and mercaptans
(RSH) together with H2S and C02 (Kohl and Riesenfeld (1985)).
Information given in Gas Process Handbook (1990) indicates that
the Sulfinol process can deliver treated gas specified to 50 ppm
C02, and thus be used prior to liquefaction in an LNG plant.
Another example using a combined chemical and physical solvent
is the Amisol process where methanol is mixed with MEA or DEA.
Quite recently, Institut Francais du Petrole (IFP) has introduced
a 2-stage process called IFPEXOL. Based on methanol as the major
constituent, this process is capable of removing acid gases and
water, giving hydrate protection and controlling the dew point
(Minkkinen and Levier (1992)).
The use of di- and triethyleneglycol together with alkanolamines,
as studied in some detail in this thesis, was first described by
Hutchinson (1939). Kohl and Riesenfeld (1985) discuss the
advantages and the problems arising when such mixtures are used.
Among the problems, the most important one was that the glycol-
amine system requires a high reboiler temperature, causing a
5
corrosive environment in the stripper and the heat exchanger.
This eventually led to a decrease in the use of amine-glycol
solutions for gas treating purposes. In recent years more
resistant metals have been developed, and additives such as
corrosion inhibitors have become available. This can lead to a
renaissance for such processes, when it is important to reduce
the number of process units to save either space or weight. Such
solvents is capable of removing water and CO2 in one step.
Savings should be obtained in cases where removal of both these
components is necessary.
One advantage using the glycol-amine process is that the steam
consumption can be lowered compared to aqueous systems. In
addition, these solutions will have the capability to reduce the
C02 content in the gas down to extremely low levels, because the
CO2 is more readily stripped from the solution.
Vaporization losses and degradation problems may occur because
of the high temperatures. According to McCartney (1948), this can
be reduced by introducing a glycol wash after the glycol-amine
absorber. To minimize degradation, amines other than MEA could
be used. Secondary amines such as DEA look promising, while
tertiary and sterically hindered amines will face problems in
nonaqueous solvents due to their resistance towards the formation
of carbamates. Versteeg (1986) has shown that tertiary amines
(MDEA) do not show significant effect on C02 solubility in
nonaqueous solutions.
Hydrocarbons are in general more soluble in nonaqueous physical
solvents than in aqueous solvents.
6
1.3 SCOPE OF THE WORK
This thesis deals with the problems related to acid gas treating
in general, and specifically to CO2 removal using alkanolamines.
Most emphasis has been put on developing simple and reliable
modelling procedures for vapor-liquid equilibria of aqueous amine
solutions. The modelling technique presented here has been tested
with the tertiary amine MDEA and the sterically hindered amine
AMP. The model predicts equilibrium partial pressures of C02 in
good agreement with experimental values.
Experimental equilibrium and pH data are presented, and the model
is based on these data and solubility data from the literature.
The objective was to develop a model which could be used at both
absorption and desorption conditions. In the case of MDEA we have
succeeded in covering the temperature interval from 25 to 140°C.
C02 loadings between 0.001 and 1 mol C02/mol amine, and C02
partial pressures between 0.00001 and 50 atm, are correlated. The
model is tested against present experimental data and data
published previously by several investigators, and found to be
accurate for the following amine molarities: 1.69, 2.00, 3.04,
4.00, and 4.28M. In the case of AMP the application range of the
model is restricted to absorption conditions.
This work also includes equilibrium measurements for nonaqueous
alkanolamine solutions at absorption temperatures. The following
systems are investigated: TEG/MEA, TEG/DEA, and DEG/MEA. The
measurements are undertaken to compare the C02 solubility in
these mixed nonaqueous solvents with the solubility in aqueous
amine solutions and other solvents containing amines, reported
in the literature. The influence of temperature and amine
concentration on the C02 solubility is also investigated.
7
The TEG/MEA-system with 0.79M MEA (10mol%) is modelled to enable
estimation of C02 partial pressures at elevated temperatures,
well outside the range where the measurements were undertaken.
Some screening measurements to determine absorption rates of C02
into five different solvents including water, are also reported.
The same solvents, with addition of MEA, are also investigated
with respect to kinetics. A string-of-discs column was used for
these experiments.
The work described in this thesis is in many ways the first
comprehensive treatment of acid gas removal processes done in our
laboratory. Previous studies have mostly been related to S02
absorption. As a result of preliminary experiments undertaken in
the start-up phase of this study, we realized that both a new
experimental set-up for high temperature solubility measurements,
as well as an improved apparatus for kinetic determinations, were
desirable. These two new experimental set-ups were built in 1991
and are now in use. Having these apparatuses at our hands, we can
conduct experimental research in most areas related to gas
treating technology. The new experimental facilities are
described in some detail in Chapter 3.
8
Chapter Two
Literature Review
2.1 VAPOR-I.IOUID EQUILIBRIUM DATA IN GAS TREATING PROCESSES
2.1.1 VXE MEASUREMENTS IN AQUEOUS ALKANOLAMINE SOLUTIONS
A number of investigators have presented vapor-liquid equilibrium
data on aqueous C02-alkanolamine systems. Some examples are:
- For MEA systems, contributions are made by Mason and Dodge
(1936), Leibush and Shneerson (1950), Muhlbauer and Monaghan
(1957), Jones et al. (1959), and Lee et al. (1975, 1976).
- DEA systems are investigated by Bottoms (1931), Mason and Dodge
(1936), Reed and Wood (1941), Murzin and Leites (1971), Lee et
al. (1972, 1974), Lawson and Garst (1976), Kennard and Meisen
(1984), and Lai et al. (1985).
- For TEA systems measurements are presented by Bottoms (1931),
Mason and Dodge (1936), Byudkovskaya and Leibush (1949), and Jou
et al. (1985).
- VLE data for aqueous DIPA are given by Isaacs et al. (1977).
More recently, equilibrium data for the tertiary amine MDEA and
also for sterically hindered amines have become available.
Equilibrium data for the MDEA system is given by Jou et al.
(1982, 1986), Bhairi (1984), Chakma and Meisen (1987), Austgen
(1989), and Lidal and Erga (1991). Sharma (1965) observed that
9
sterical hindrance has a pronounced effect on the stability of
the carbamates, see section 2.2 and 6.1. The sterically hindered
amines were later introduced to acid gas treating by Exxon (Chem.
Eng. News (1981)). A few investigators have reported equilibrium
data in some hindered amines. Measurements on one of the best
known hindered amines, AMP, have been undertaken by Sartori and
Savage (1983), Komorowicz and Erga (1987), Roberts and Mather
(1988a), Teng and Mather (1989, 1990), Erga and Lidal (1990), and
Tontiwachwuthikul et al. (1991).
In a research report from the Gas Processors Association,
equilibrium solubility of CO2 in aqueous solutions of MEA, DGA,
DEA, and MDEA are given (Maddox et al. (1987)). Equilibrium data
for DGA are also presented by Martin et al. (1978) and Dingman
et al. (1983) .
2.1.2 VLE MODELLING FOR C02 IN AQUEOUS ALKANOLAMINE SOLUTIONS
Mason and Dodge (1936) made the first attempt to correlate the
equilibrium solubility data for C02 in alkanolamines. Since the
reactions between amines and CO2 had not been properly
investigated at that time, they were limited to use curv* . itting
methods.
A method for predicting C02/amine equilibria in aqueous solutions
based on the use of apparent equilibrium constants, without
activity coefficients, was described by Danckwerts and McNeil
(1967). They used the same approach as Van Krevelen (1949) had
developed for aqueous solutions of ammonia and C02. The apparent
equilibrium constants, and their dependence on the ionic strength
of the solutions, are used to describe the chemical equilibria.
A similar approach was used by Kent and Eisenberg (1976) for MEA
10
and DEA solutions/ the main difference being that the apparent
equilibrium constants were regarded as constant, irrespective of
the Ionic strength. In the Kent-Eisenberg method, the approach
made by Danckwerts and McNeil (1967) was modified by forcing the
apparent equilibrium constants to fit published equilibrium data
as a function of the temperature.
An early attempt to include activity coefficients into a
predictive model for C02/amine equilibria was made by Klyamer and
Kolesnikova (1972), and was further developed to describe the
C02/H2S/amine equilibria by Klyamer et al. (1973). They used a
method proposed for the H2S/amine system by Atwood et al. (1957),
where the activity coefficients of all ionic species are assumed
to be equal. According to Deshmukh and Mather (1981), the
generalized model given by Klyamer et al. is algebraically
equivalent to the Kent-Eisenberg model if the activity
coefficients are set equal to unity.
These earlier models exhibit a useful description of the chemical
equilibria for many compositions of the aqueous amine solutions.
However, they often fail at compositions outside the range where
the apparent equilibrium constants are fitted, or the activity
coefficients are determined. To be able to broaden the range
where such models could be applied, one has to use equilibrium
constants expressed as functions of C02 concentration, amine
molarity, and temperature.
Realizing that the Kent-Eisenberg model has certain limitations,
improvements of the method have been achieved by several
investigators over the years, such as Chakma and Meisen (1990)
for the C02/DEA/water system. The most important improvement is
that the apparent equilibrium constant of the main DEA-C02
reaction is recorrelated using a more comprehensive set of
11
experimental data. In the new correlation the apparent
equilibrium constant is expressed as a function not only of the
temperature» but also of the CO2 concentration and the amine
molarity- Jou et al. (1982) used a similar procedure to correlate
their VI.E data for the MDEA system.
A "new generation" of equilibrium models has been developed in
recent y&ars. A. thermodynamic framework was established by
Edwards et al. (1975, 1978) to calculate gas-liquid equilibria
in aqueous solutions containing one or more volatile weak
electrolytes, such as C02. The framework was so constructed that
the equilibrium compositions of multisolute systems could be
predicted using only binary interaction parameters. Beutier and
Renon (1978) used a similar approach, in which two ternary
interaction parameters were fitted to the actual ternary
experimental data. In this way, a better agreement between
calculated and expsrimental data for a two-solute system, was
obtained.
Deshmukh and Mather (1981) proposed a mathematical model based
on the extended Debye-Hiickel theory of electrolyte solutions,
using the Guggenheim (1935) equation, which represents the
activity coefficients by the use of two terms. The first of these
terms is the standard Debye-Huckel term representing the
electrostatic forces. The second term includes binary interaction
parameters accounting for short-range Van der Waals forces.
Because many of these interaction parameters were unavailable,
Deshmukh and Mather adjusted some of them to ternary VLE data
(CO2/MEA/H2C) and H2S/MEA/H20) . In this, they made use of the
assumption that the interaction parameters for the species which
were present in very small concentrations could be neglected. The
fugacity coefficients were calculated using the Peng-Robinson
(1976) equation of state. The model exhibits a good fit to the
12
experimental data for the MEA system except for high C02
loadings. This especially applies to the quaternary system;
C02/H2S/MEA/H20, where the assumptions made seem to lead to an
underprediction of the equilibrium partial pressures. Chakravarty
(1985) in his work used a similar approach. By extensive use of
literature data an equilibrium model was developed, applicable
to four single amine systems (MEA, DEA, DIPA, and MDEA) as well
as blends of amines (MEA/MDEA and DEA/MDEA).
Based on a generalized excess Gibbs energy model that treats both
long-range electrostatic interactions between ions, and short-
range interactions between all liquid phase species, Austgen
(1989) has developed a thermodynamically consistent model
describing the vapor-liquid equilibria in acid gas-amine-water
systems. The vapor phase fugacity coefficients were calculated
by the use of the Redlich-Kwong-Soave equation of state (Soave
(1972)). The Electrolyte-NRTL equation (Chen and Evans (1986))
was used to represent the liquid phase activity coefficients. The
Electrolyte-NRTL equation requires binary interaction parameters
to be estimated from experimental data. In addition the carbamate
stability constant was treated as an adjustable parameter within
the VLE model. The model was extended to describe C02
solubilities in blends of amines (MEA/MDEA and DEA/MDEA).
An attempt to correlate the CC^/AMP/water system was made by
Chakraborty et al. (1986) based on equilibrium constants at
vanishingly small ionic strength. As would be expected, the model
could not describe the equilibrium curve very well at high C02
loadings. Tontiwachwuthikul et al. (1991) proposed a modified
Kent-Eisenberg model for the same system, and obtained a better
agreement between calculated and experimental data.
VLE models for aqueous DGA solutions have been developed by
13
Dingman et al. (1983) and Hu and Chakma (1990). While Hu and
Chakma based their method on similar principles as Kent and
Eisenberg (1976), Dingman et al. took a more fundamental
approach, by using the framework introduced by Edwards et al.
(1975), and thereby included activity coefficients in the
description of the vapor-liquid equilibria.
In this review of the literature, no reports were found regarding
the use of measured pH data in the modelling of VLE in
alkanolamine systems.
2.1.3 VLE MEASUREMENTS IN MIXED NONAQUEOUS SOLVENTS
Parts of this thesis concern the absorption of C02 into a mixed
solution, containing an alkanolamine and a glycol solvent, with
virtually no water present. Since Hutchinson (1939) proposed the
glycol-amine process for simultaneous acid gas removal and
dehydration, few investigations on the C02 solubility in these
systems have been reported. Most of the literature published
deals with the technical specifications of the process, or with
the limitations and problems related to the process. Examples are
Chapin (1947), Kohl and Blohm (1950), Polderman et al. (1955),
and Holder (1966).
Literature data are scarce for all systems combining amines and
nonaqueous physical solvents. Murrieta-Guevara and Tre jo
Rodriguez (1984) presented solubility data for C02, H2S, and
methane in nonaqueous mixtures of alkanolamines (MEA, DEA) and
physical solvents (n-methyl-pyrrolidone (NMP), propylene
carbonate (PC)). Murrieta-Guevara et al. (1992) introduced some
additional data for the solubility of C02 in NMP solutions
containing either MEA or DEA. Solubilities of C02, H2S, and
14
ethane in PC, NMP, and sulfolane (tetrahydrothiophene dioxide)
in mixtures with alkanolamines, were measured by Rivas and
Prausnitz (1979). Dimov et al. (1976) measured low pressure VLE
data for MEA solutions of ethyleneglycol, NMP and
tetrahydrofurfuryl alcohol at different water levels (also
without water). Leites et al. (1972) compared the C02 solubility
between several nonaqueous solvents containing MEA. Takeshita and
Kitamoto (1988) measured the C02 solubility in complete water
free solutions of methanol/ octane and triethylamine with
different primary and secondary amines.
Woertz (1972) investigated a number of mixtures containing an
amine, a physical solvent, and a small amount of water (3 or
10vol%). In the literature one can find VLE data for aqueous
systems containing both an amine and a physical solvent, an
example being the data of Roberts and Mather (1988b), where the
solubility of acid gases in a mixed solvent of 16.5wt% AMP,
32.2wt% sulfolane, and 51.3wt% water was reported. Oyevaar et al.
(1989) measured the C02 solubility in aqueous ethyleneglycol
solutions containing DEA.
2.1.4 VLE MEASUREMENTS IN PORE PHYSICAL SOLVENTS
TEG-CO2 equilibria without amine present were measured by
Takahashi et al. (1984) and Jou et al. (1987). Takahashi et al.
also presented solubility data for the DEG-C02 system. C02
solubilities in other useful physical solvents are reported by
a series of investigators. Some examples are: Isaacs et al.
(1977), Laddha et al. (1981), Sweeney (1984, 1988), Roberts and
Mather (1988c), Murrieta-Guevara et al. (1988), Jou et al.
(1990a, 1990b), and Yogish (1991). Fogg and Gerrard (1990) have
collected published C02 solubility data for more than 100
15
different solvents.
2.1.5 VLE MODELLING TECHNIQUES IN PHYSICAL SOLVENTS
The solubility of acid gas in a pure physical solvent can be
described by Henry's law (Eqn. (5.7)). However, at higher
concentrations and partial pressures most systems show a
considerable deviation from the linearity assumed in the simple
form of Henry's law (Fogg and Gerrard (1990) and Carroll (1991)).
Thus, in order to successfully correlate experimental results up
to high concentrations, one needs a method based on an equation
of state valid for the solvent and dilute solutions of the solute
in the solvent. Such an approach was used by Jou et al. (1987,
1990a) to correlate the solubility of C02 and H2S and the lower
alkanes in solutions of TEG and sulfolane. They used the Peng-
Robinson (1976) equation of state, and obtained interaction
parameters for these systems. These interaction parameters were
further used to determine the three parameters in the equation
developed by Krichevsky and Iliinskaya (1945). The Krichevsky-
Iliinskaya equation has been shown to be applicable also for
mixtures of components with strong intermolecular interactions.
For such systems simple equations of state are insufficient for
description of the phase behaviour (Jou et al. (1987)).
For the mixed solvents described in section 2.1.3, correlations
for the C02 solubility are given by Rivas and Prausnitz (1979)
and Roberts and Mather (1988b). Rivas and Prausnitz determined
equilibrium constants to describe the chemical equilibria for the
absorbed gas and the chemical solvent. Roberts and Mather used
the solubility model of Deshmukh and Mather (1981) to predict the
equilibrium partial pressures of C02 in a mixture of a chemical
active agent (AMP), a physical solvent (sulfolane), and water.
16
2.2 CHEMISTRY OF COo - AMINE SYSTEMS
2.2.1 INTRODUCTION
In this work, emphasis has been put into the developing of simple
and reliable methods for representing the equilibria of C02-amine
systems. This cannot be done without having an understanding of
the chemistry encountered in these systems.
Several comprehensive investigations have been undertaken to
study the kinetics of the reactions in alkanolamine processes.
As a result, rate-based process models for acid gas removal are
developed, see for example Glasscock (1990) and Carey (1990).
Carey gives an overview of rate-based models available. Tomcej
(1987) developed a nonequilibrium stage model to simulate acid
gas absorption into alkanolamine solutions. This model has become
commercially available as a simulation program under the name of
AMSIM. Other models for commercial use are the TSWEET program.
According to informations given at the last GPA convention in
1992, TSWEET will soon offer the capability of simulating systems
using blends of amines (Bullin et al. (1992)). A simulation
program described by Sardar and Weiland (1985) is also
commercially available. Several of the larger companies such as
DOW Chemical Co. are known to have in-house amine process
simulators (Katti and Langfitt (1985)). Based on a mass transfer
model described in the literature (Versteeg et al. (1989, 1990)),
researchers at Twente University have developed a simulation
program called SIMULTER for the calculation of the absorption
rates of Co2 and H2S into aqueous solutions of tertiary amines.
Versteeg (1986) in his work studied the reaction between C02 and
different alkanolamines both in aqueous and nonaqueous solutions.
Khalil (1984), Yu (1985) and Al-Ghawas (1988) studied the
17
kinetics of absorption of H2S and C02 in aqueous MDEA solutions.
Al-Ghawas (1988) and Glasscock and Rochelle (1989) recapitulate
the different mass transfer models presented in the literature.
Such models are the film theory model, still surface models,
surface renewal models, penetration models, and combinations of
these models. This theory will not be taken any further in this
thesis.
Complete understanding of the mechanism of the reaction of C02
with alkanolamines is still ahead of us. For well investigated
systems, however, kinetic expressions which are in good agreement
with experimental data, are established. In the following
sections, the basic C02 - amine chemistry, and the kinetics
proposed in the literature for MDEA and AMP systems, are
presented. The last section in this chapter presents different
experimental techniques for determinations of reaction kinetics.
2.2.2 REACTIONS BETWEEN C02 AND AMINES IN AQUEOUS SOLUTIONS
Compared with the instantaneous proton transfer reaction when H2S
reacts with an alkanolamine, the reaction between C02 and
alkanolamines is more complex, and the reaction rate depend
highly on the structure of the alkanolamine molecule. Primary and
secondary amines, have the capability to react with C02, forming
carbamate ions. These are amines with one or two carbon-
containing groups attached to the nitrogen atom. Tertiary amines,
like TEA and MDEA, with three carbon-containing groups attached
to the nitrogen atom, cannot form carbamates, and bicarbonate
formation becomes the only main reaction.
18
H C2H4OH
I I H - N - C2H4OH H - N - C2H4OH
Monoethanolamine (MEA) Diethanolamine (DEA)
C2H4OH C2H4OH
I C2H4OH - N - C2H4OH CH3 - N - C2H4OH
TriethanolaminetTEA) Methyldiethanolamine(MDEA)
CH 3
HO - C H 2 - C - NH2 I
CH 3
2 - amino - 2 - methyl - I - propanol (AMP)
Figure 2.1 Molecular structure of amines used in acid gas removal processes
19
Figure 2.1 shows the molecular structure of MEA, DEA, and TEA,
as well as the two amines especially studied in this
investigation, MDEA and AMP. AMP is denoted a sterically hindered
amine (Sartori and Savage (1983)), since the amino group is
attached to a tertiary carbon atom. For definition of a
sterically hindered amine, see Chapter 1.
An important reaction in aqueous solutions containing C02 is the
"OH"-reaction":
C02 + OH" = HC03" (2.1)
At pH values above 8, the most important reaction mechanism of
this reaction is the direct one, where Eqn. (2.1) is the actual
kinetic step (Astarita et al. (1983)). At lower pH values/ a
competing mechanism occurs. In this C02 is first hydrated:
C02 + H20 = H2C03 (2.2)
R e a c t x o n ( 2 . 2 ) i s t h e n fo l lowed by t h e d i s s o c i a t i o n o f t h e
c a r b o n i c a c i d :
H2C03 = HC03" + H+ ( 2 . 3 )
For reactions involving amines at sufficiently high pH-vaJues
Astarita et al. (1983) suggest that in general three main
reactions should be considered. Taking a primary amine (RNH2) as
an example:
Carbamate Formation,CF: C02 + 2RNH2 = RNH3+ + RNHCOCT (2.4)
Bicarbonate Formation,BF: C02 + RNH2 + H20 = RNH3+ + HC03" (2.5)
Carbamate Reversion,CR: RNHCOO" + H20 = RNH2 + HC03" (2.6)
20
We now introduce the C02 loading, y, expressed as mol C02/mol
amine. For primary and secondary amines, CF will take place at
y<0.5, CR at y>0.5, and BF at all values of y. For tertiary
amines, CF does not take place, and BF is the only reaction. For
hindered amines, CF may be very small or negligible.
Al-Ghawas (1988) in his work also includes the direct formation
of carbonic acid by the reaction of C02 and H20 (Eqns. (2.2-
2.3)), and also an alkylcarbonate formation reaction. However,
according to Astarita et al. (1983) and Yu et al. (1985), both
of these reactions will proceed to a negligible extent at the pH-
values and temperatures usually encountered in gas treating
processes.
According to Danckwerts (1979) and Astarita et al. (1983), with
later support also by other investigators, the CF mechanism is
believed to proceed by the steps:
C02 + RNH2 = RN+H2COO" (2.7)
RN+H2COO" + RNH2 = RNH3+ + RNHCOO" (2.8)
This zwitterion mechanism was first proposed by Caplow (1968) for
amines without alcoholic groups. The rate-determining step in
this mechanism is believed to be the zwitterion formation (Egn.
(2.7)). This is verified for the MEA system, where a rate
equation as follows has been verified:
r = kCF • Cc02 • CRNH2 (2.9)
For some of the other amines, such as DEA, there are data
supporting Eqn. (2.9), while other data suggest the reaction to
be second-order with respect to the amine (Hikita et al. (1977)).
21
Versteeg and van Swaaij (1988a) explain how the same reaction,
for different amines, can assume different reaction orders.
Among the first to investigate the reactions between C02 and
amines were Danish researchers. They studied the carbamate
formation from a number of amines, such as dimethylamine
(Faurholt (1925)) and glycine (Jensen et al. (1954)). The
reactions between C02 and alkanolamines were also studied. The
rate of reaction of C02 with both MEA and DEA (Ballund Jensen et
al. (1954)), as well as TEA (Jørgensen and Faurholt (1954)), was
measured. Their work is also commented on in the next section.
A possible mechanism of the CF reaction is described in the next
section.
2.2.3 REACTION KINETICS BETWEEN C02 AND AQUEOUS MDEA
MDEA is today the most used tertiary amine for acid gas removal.
MDEA has outdone for example TEA, which was the first amine to
be used for gas sweetening purposes (Bottoms (1930)). When the
correct additives are used, MDEA offers several advantages over
other amines also for bulk C02 removal (Bullin et al. (1990,
1992)). This is discussed elsewhere in this thesis (Chapter 5).
A number of investigations have been conducted on the kinetics
of the MDEA-C02 system in recent years. Examples are Barth et al.
(1981 1984), Haimour and Sandall (1984), Yu et al. (1985),
Versteeg (1986), Haimour et al. (1987), Tomcej and Otto (1989),
Crooks and Donnellan (1990), and Al-Ghawas and Sandall (1991).
There are some controversies in the literature about the reaction
rate of C02 with MDEA. Glasscock (1990) suggests that the
discrepancies found in the literature is due to the fact that the
22
reactions involved is more complex than assumed by the authors.
It should also be emphasized that, according to Glasscock and
Rochelle (1989) and Littel et al. (1990), a serious depletion of
0H~ toward the gas-liquid interface usually occurs. The
contribution of the C02 reaction with OH" to the observed
reaction rate may therefore have been overestimated by previous
investigators.
The most accepted mechanism for the BF reaction involving MDEA,
is that the amine acts as a base catalyst for the C02 hydration
reaction. This is the same theory as presented by Donaldson and
Nguyen (1980) for TEA solutions. Haimour et al. (1987) explain
their observations using this theory, and have found the
hydrolysis rate of C02 in MDEA to be second order, i.e. first
order with respect to both the amine and the C02 concentration.
Haimour et al. (1987) have found the rate constant to be 2.47
1/mol s, which are in acceptable agreement with Barth et al.
(1984), but about half the value determined by Blauwhoff et al.
(1984). Yu et al. (1985) also concludes that MDEA acts as a
homogenous catalyst for C02 hydrolysis, and they speculate that
a zwitterion could be formed and constitute the intermediate in
the catalytic path. Versteeg and van Swaaij (1988b) also
concludes that the base catalysis of the C02 hydration describes
the reaction between C02 and tertiary amines.
A possible reaction between C02 and the alcoholic group(s) on the
alkanolamine molecule, is not favored by the pH levels at which
acid gas treating processes usually occur (7-10). Such a
reaction, forming an alkylcarbonate, require a very high pH to
be able to proceed to a significant extent. Jørgensen and
Faurholt (1954) made experiments with C02 and TEA at a pH value
of 13, and found that monoalkyl carbonate was indeed formed.
Blauwhoff et al. (1984) concluded from their measurements, that
23
no alkylcarbonate formation occured in the case of TEA and MDEA
at pH values lower than 10.7, which is close to the pH range of
industrial interest. The effect of basicity on the kinetics of
C02 absorption in tertiary amines is discussed further by
Benitez-Garcia et al. (1991).
2.2.4 REACTION KINETICS BETWEEN C02 AND AQUEOUS AMP
Rochelle (1991) states that hindered amines appear to have much
of the same kinetic behavior as tertiary amines. Toman and
Rochelle (1990) have investigated the C02 absorption rates into
aqueous solutions of the severely hindered amine 2-(tert-
butylamino) ethanol (TBE). According to Sartori et al. (1987),
a severely hindered amine is characterized by a very low rate of
C02 absorption.
Kinetic studies of the sterically hindered amine 2-piperidine
ethanol (PE), were done by Shen et al. (1991).
The kinetics of the reaction between C02 and 2-amino-2-methyl-1-
propanol (AMP) are studied by several researchers, such as
Sartori and Savage (1983), Yih and Shen (1988), and Bosch et al.
(1990).
Chakraborty et al. (1986) and Zioudas and Dadach (1986) measured
the absorption rates of C02 into AMP solutions. Bosch et al.
(1989) used these measurements to demonstrate that no new
reaction paths were necessary for explaining the observed
absorption behavior of AMP solutions.
Yih and Shen (1988) concludes that the reaction is first order
with respect to both C02 and amine, and the rate constant was
24
found to have a value of 1270 1/mol s. The authors assume that
the reaction proceeds via a zwitterion mechanism, as explained
in section 2.2.2.
Based on new absorption rate experiments, Bosch et al. (1990)
have difficulties in explaining the absorption behaviour solely
by the zwitterion mechanism. They suggest that the kinetics of
the CO2-AMP system might be more complex. For example an
alkylcarbonate formation could have some influence. The direct
reaction with OH" (Eqn. (2.1)) should also be considered.
In addition to the base catalysis reaction mechanism described
in the previous section, the Bp reaction (Egn. (2.5)) could
proceed by the reactions given in Eqns. (2.1)-(2.3), followed by
the instantaneous proton reaction with the amine molecule.
Sartori and Savage (1983) describe a possible parallel path,
where the zwitterion formed in the first step of the CF reaction
(Eqn. (2.7)), undergoes a direct reaction with water forming
bicarbonate and ammonium ions.
2.2.5 REACTION KINETICS IN NONAQUEOUS SOLUTIONS
Less literature is found on the kinetics of C02 reactions in
nonaqueous solutions. Here we shall restrict ourselves to mention
some of the most important publications on this subject in recent
years.
Alvarez-Fuster et al. (1980, 1981) present rate data for C02
absorption in mixed solvents. Solutions with MEA, DEA and
cyclohexylamine (CHA) in ethanol, toluene, and ethyleneglycol
(ETG) were investigated. The reaction rate data was interpreted
using the zwitterion reaction mechanism. It was found that most
25
of these systems exhibit a third order kinetics, first order with
respect to C0 2 and second order with respect to amine. However,
the C02-CHA-ETG system was found to be first order to both C0 2
and amine.
Since this thesis deals partly with mixed solvents using glycols,
the work done by Alvarez-Fuster et al. is of interest, since it
includes solutions containing ETG. Other investigators who have
studied the kinetics of similar systems, are Såda et al. (1985a),
and Oyevaar et al. (1990). Oyevaar et al. used the absorption
kinetics of the aqueous C02-DEA-ETG system for determining
interfacial areas in gas-liquid reactors.
Såda et al. (1985a, 1985b, 1986a, 1986b, 1989) have undertaken
a series of investigations on C02 absorption in nonaqueous amine
solutions. In most of their experiments different alcohols are
used as the nonaqueous solvent. They conclude that the zwitterion
mechanism, here explained earlier, can describe the reaction
between C02 and primary and secondary amines. As for the tertiary
amine TEA, they expect that in alcoholic solutions, dissolved C02
will react with solvated TEA forming an ion pair.
This last statement is in contrast with the finding of Versteeg
and van Swaaij (1988b) for the MDEA-ethanol system. They conclude
that in nonaqueous solutions no reaction, not even alkylcarbonate
formation, occurs between C02 and the tertiary amine. This may
be an important finding, since it would exclude MDEA as a useful
amine for the purpose of simultaneous removal of water and C02,
using a glycol-amine solution.
Versteeg and van Swaaij (1988a) have also studied the kinetics
in nonaqueous solutions of primary and secondary amines, and they
conclude that the solvent used has a pronounced effect on both
26
renction order and reaction rate.
2.2.6 EXPERIMENTAL EQUIPMENT FOR KINETIC DETERMINATIONS
Descriptions of various laboratory experimental set-ups for
kinetic determinations are given by Danckwerts (1970) and
Astarita et al. (1983). Astarita et al. classify mass transfer
experiments into three types. Type 1 is recognizable by that the
physical mass transfer coefficient can be estimated from the
solution of the appropriate hydrodynamic equations. The
interfacial area is known at the outset. Examples of type 1 set
ups are the laminar jet, the short wetted-wall column, and the
one-sphere apparatus. In type 2 units the interfacial area is
still known, but the physical mass transfer coefficient cannot
be estimated from first principles. Examples are the stirred
cell, the long wetted-wall column, the string-of-spheres, and the
string-of-discs columns. In type 3 units neither the interfacial
area nor the mass transfer coefficient are known. Examples are
the sparged cell, the single sieve tray, and the single bubble
cap plate. Type 1 is in general prefered to type 2, which again
is prefered to type 3.
As we have seen, a large number of investigations are reported
in the literature on the kinetics of the reactions involved in
acid gas removal processes. Several of the above mentioned set
ups have been used, and here some examples will be given.
Blauwhoff et al. (1984), Khalil (1984), Yu efc al. (1985),
Versteeg (1986), Haimour et al. (1987), Sada et al. (1989),
Littel et al. (1990), Oyevaar et al. (1990), and Glasscock
(1990), used a stirred vessel. Sartori and Savage (1983), Tomcej
(1987), Al-Ghawas (1988), and Benitez-Garcia (1991), used a one-
27
sphere set-up. Haimour and Sandall (1984) and Al-Ghawas et al.
(1989), used a laminar jet apparatus. A. wetted-wall technique was
used by Alvarez-Fuster et al. (1980, 1981), Yih and Shen (1988),
Såda et al. (1989), and Toman and Rochelle (1990).
Barth et al. (1981, 1984) used a stopped flow method with optical
detection of the proceeding of the reaction, while Crooks and
Donnellan (1990) used a stopped flow method with conductimetric
detection. Donaldson and Nguyen (1980) used a tracer 14C02
membrane transport technique.
28
Chapter Three
Experimental
3.1 VAPOR-LIQUID EQUILIBRIUM MEASUREMENTS
3.1.1 EQUILIBRIUM EQUIPMENT FOR TEMPERATURES UP TO 70°C
The thermostatic equilibrium equipment, which is essentially the
same as the one used by Erga (1988) in S02 absorption studies,
is shown in Fig. 3.1 . Here, an IR C02-analyzer (URAS 3G, Hartmann
& Braun) measures C02 partial pressures between 0.0050 and 0.30
atm. The instrument was calibrated at atmospheric pressure with
4 standardized gases, containing 1, 5, 10, and 30vol% of CO2,
respectively. The CO2 equilibrium partial pressure was computed
from the formula:
PC02 = 10"6 ' PP m v • (P - 6p) (3.1)
where ppmv = instrument reading, P = total pressure in the gas
leaving the last gas wash bottle (the pressure drop downstream
from this point to the manometer was negligible) and 6p =
difference in water vapor partial pressure between the gas
leaving the buffer solution and the condenser, corrected for the
reduction in vapor pressure due to the amine concentration. In
the case of nonaqueous solvents there is no 6p-correction.
29
Figure 3.1 Gas-liquid equilibrium equipment
3.1.2 A NEW EQUIPMENT FOR TEMPERATURES UP TO 120°C
The equipment described in the previous section has been used in
a number of investigations prior to this work. It constitutes a
very simple and rapid way of obtaining VLE data for aqueous
systems at temperatures below 70°C. To be able to undertake
experiments at temperatures above 70°C, it was recognized that
our laboratory was in need of a new equipment. Equilibrium data
at higher temperatures up to 120-140°C are important in the
design of the regeneration units in acid gas treating processes.
An equipment for this purpose was built, and it is now in use.
The new equipment consists of a 300ml autoclave made in the inert
material Hastelloy C. The autoclave acts as an equilibrium cell
with a magnetically driven stirrer. The autoclave was delivered
by PARR Instrument Co., and it was furnished with two seeglasses.
As can be seen from Fig. 3.2, the new set-up consists basically
of the same components as the low temperature equipment described
above. The main differences are: i) in the new apparatus all
parts are made in heat and corrosion resistant material, and
ii) the C02-analyzer (URAS 3GH, Hartmann & Braun) works at a
30
temperature of 160°C. Therefore, there is no need to condense the
gas stream going into the C02-analyzer. This makes the gas
analyzing much simpler and also more accurate. The water vapor
partial pressure can now be the same in the autoclave and in the
analyzer. This makes the correction in 6p in Eqn. (3.1),
unnecessary.
The new equipment was built so that it could be extended to
measure VLE data at pressures above atmospheric. By introducing
a back pressure valve downstream from the autoclave, and
replacing the gas compressor with a more powerful one, C02
partial pressures up to 7-8 atm could be measured.
I Tl V
< '• ? -
lii
Figure 3.2
• _ , /
AU10CUWE
in
—-vr.:f[)-e-<l- i • EX)
C02 ANALYZER
J
FLOWMETER
- X —
TT¥
-x^
<5>
*
New gas-liquid equilibrium equipment, capable of measuring solubilities at temperatures encountered in desorption units
3.2 pH MEASUREMENTS
pH values were measured as a function of the C02 loading using
an Orion Ross Sure-Flow combination electrode, and an Orion SA720
pH-meter. The measurements were performed at different
temperatures, using a thermostatic water bath.
31
3.3 KINETIC MEASUREMENTS
3.3.1 STRING-OF-DISCS-COLUMN
The kinetic measurements reported in this study were all done on
a string-of-discs-column. This apparatus was first introduced by
Stephens and Morris (1951) and also described by Morris and
Jackson (1953). Astarita et al. (1983) classifies the disc-column
as a type 2 laboratory mass transfer unit, confer section 2.2.6.
These are units for which the interfacial area is known, but the
mass transfer coefficient cannot be predicted from first
principles. For more details, reference is made to Morris and
Jackson (1953).
A sketch of the apparatus is given in Fig. 3.3. The measurements
were all undertaken at feed temperatures of 20°C (±1°C) for both
phases. However, significant temperature rise was noticed due to
the heat of reaction when C02-amine reactions were studied in
this equipment.
TO ri'HE HOOD
• Æ SOAP
HtTtR
jpai
LIQUID Tilt m
at
-Q+ r*cm
SOLUTION i i /mr
Figure 3.3 S t r i n g - o f - d i s c s absorber
32
3.3.2 ONE-SPHERE APPARATUS
ln the string-of-discs column, the physical mass transfer
coefficient cannot be estimated from first principles because the
hydrodynamic equations cannot be solved explicitly. This is the
reason why it was decided to build a type 1 apparatus (Astarita
et al.'s (1983) notation), where the physical mass transfer
coefficient can be obtained by solving the appropriate
hydrodynamic equations. We decided to use a "wetted one-sphere"
apparatus. This set-up is described in detail in the literature
by Tomcej (1987) and Al-Ghawas (1988), confer section 2.2.6.
A schematic description of the operation of the one-sphere
apparatus which would also cover the string-of-discs column, is
given in Fig. 3.4, Astarita et al. (1983).
The sphere is made of the highly resistant material Hastelloy C
in order to reduce the risk of corrosion. The sphere is polished
to give it a very smooth surface. The diameter of the sphere is
50.0 mm. A special construction is provided to ensure the rod on
which the sphere is mounted, to be accurately centered in the
opening of the liquid feed distributor.
The sphere is placed in a thermostatic environment, making it
possible to do experiments at controlled elevated temperatures,
pertaining to desorption conditions.
The equipment was constructed at the Norsk Hydro Research Centre
in Porsgrunn, where it is presently located. The apparatus has
now been tested, and the first experiments have been conducted
(Eimer (1992)). The results so far look promising as to the
applicability of this equipment.
33
Liquid in
Gas in
Absorber
Soap film meter
£x* Gas out
Liquid out
Release valve
Gas feed
Figure 3.4 Schematic of operation of string-of-discs and one-sphere apparatus
3.4 CHEMICALS AND GASES
The amines used in the experiments were supplied by Merck. MDEA
had a minimum purity of 98% and less than 0.2% water. AMP had a
minimum purity of 95% and less than 0.3% water. MEA had a minimum
purity of 99%, and DEA of 98% and less than 0.3% water.
The TEG was supplied by Merck, and had a minimum purity of 98%,
and contained less than 0.3% water. The DEG was supplied by
British Drug House (BDH), and had a minimum purity of 99.5%, and
contained less than 0.2% water.
As can be seen, even the amines and glycols used were not water
free. When the expression nonaqueous is used for some of the
34
experiments undertaken in this work, it means that the solutions
contained less than 0.3% water.
The C02 and N 2 gases were supplied by AGA and had purities of
99.9% and 99.99%, respectively. The standardized gases, used to
calibrate the IR gas analyzer, were supplied by Hydrogas and AGA,
and were delivered with analysis certificates to accurately
certify the CO2 content in the N2 gas.
3.5 LIQUID ANALYSIS
3.5.1 C0 2 CONCENTRATION
Two different methods were used to determine the C02 content of
the liquid samples: At the outset the measurements of the C02
concentration were determined by injection of a 5 ml sample into
a thermostatic closed vessel (1000 ml) containing 25 ml of 5M
HC1. The pressure increase caused by the liberated C02 gas, was
measured with an accurate dp-cell and recorded. The system was
calibrated using solutions of known bicarbonate concentrations.
Because of some difficulties in obtaining satisfactory
reproducibility with the above described method, the liquid
analysis was changed for the MDEA-studies. The C02 concentration
in the liquid phase was now determined by injecting a sample into
a 0.1M NaOH solution, then adding BaCl2 in excess to precipitate
the carbonate as BaC03. After at least 36 hours, the precipitated
BaC03 was filtered, then dissolved in distilled water, and
finally titrated with standard 0.1M HC1 (Jou et al. (1982)). The
endpoint was verified using pH measurements. The carbonate
content of the NaOH solution was corrected for. This procedure
was used in all experiments reported in this work, except for the
35
measurements on the AMP-system, where the pressure increase
detection method, described above, was used.
3.5.2 AMINE CONCENTRATION
Amine concentrations were measured using an acid-base titration
with standard 1.0M HCl, using methyl red as indicator.
3.6 GAS ANALYSIS
The gas from the equilibrium cell, see Fig. 3.1, was analyzed by
using a continuous infrared gas analyzer with a measuring limit
up to 30vol% C02. For measurements of extremely low C02 partial
pressures, the measuring cell was substituted with one adapted
for a lower measuring limit (10vol%).
As mentioned in section 3.1.1, the analyzer was calibrated using
4 standardized gases. The atmospheric pressure was monitored with
a high accuracy barometer.
36
Chapter Four
Experimental Results
4.1 VAPOR-LIQUID EQUILIBRIUM MEASUREMENTS
4.1.1 C02 SOLUBILITY IN AQUEOUS MDEA SOLUTIONS
Results of the C02 solubility measurements for aqueous solutions
of 4M MDEA at 30, 45, and 60°C are summarized in Appendix A, and
are presented graphically in Fig. 4.1.
Similarly, C02 solubility measurements for 4.28M aqueous MDEA at
25, 40, and 70CC, summarized in Appendix A, are presented
graphically in Fig. 4.2. In this figure, also literature data are
presented.
The solid lines in both figures represent the model developed in
section 5.2.
37
0 0.1 0.2 0.3 OA 0.5 0.6
y Imol C02/mol amiucl
Figure 4.1 Solubility of C02 in aqueous 4.00M MDEA solutions at 30, 45, and 60°C
38
y [rnol C02/mol amincl
Figure 4.2 Solubility of C02 in aqueous 4.28M MDEA solutions at 25, 40, and 70°c, compared with literature data D • Jou et al. (1982) O • This work
39
4.1.2 CO, SOLUBILITY IN AQUEOUS AMP SOLUTIONS
Results of the C02 solubility measurements in aqueous solutions
of 3M AMP at 40 and 50°C axe summarized in Appendix A, and are
presented graphically in Fig. 4.3 together with literature data.
The lines in the figure represent the model developed in Chapter
6.
0.2 0.4 0.6 0.B
y [mol COj/mol AMP J
1.0
Figure 4.3 Solubility of C02 in aqueous 3.00M AMP solutions, compared with literature data • Sartori and Savage (1983) O Komorowicz and Erga (1987) O This work, 40°C © This work, 50°C
40
4.1.3 C02 SOLUBILITY IN NONAQUEOUS AMINE SOLUTIONS
Measurements of C02 solubility have been undertaken for the
following nonaqueous systems; TEG/MEA (5 and 1Omol% MEA), TEG/DEA
(5, 10, and 13.6mol% DEA), and DEG/MEA (5 and 10 mol% MEA), at
30, 40, 50, and 70°C. The results are summarized in Appendix B,
and are presented graphically in Figs. 4.4-4.9. The lines in the
figures are for most systems pure curve-fitting lines obtained
using a power function. However, for the particular system with
10mol% MEA in TEG, a predictive model has been developed in
section 7.1 and is used in Fig. 4.5.
e m
o.
0.001
0.01
0.00 0.20 0.40 0.60 y [mol C02/mol amine]
0.80
Figure 4.4 Solubility of C02 in TEG solutions containing 5mol% MEA at 30, 50, and 70°C
41
100 3
10 =
e id
0.1 =
Q. 0.01 -r
0 . 0 0 1 =
0 . 0 0 0 1 =
0 . 0 0 0 0 1 I r i i i i i i i i i i i i i i i r i i i i i i i i i i i i i i i i r i i i i i i i i i r i 0.00 0.10 0.20 0.30 0.40
y [mol C02/mol amine]
F igure 4.5 S o l u b i l i t y of C02 i n TEG s o l u t i o n s con ta in ing 10mol% MEA a t 30, 50, and 70°C. The l i n e s follow Eqn. (7.4)
42
0.00 0.20 0.40 0.60 0.80 y [mol C02/mol amine]
Figure 4.6 Solubility of C02 in TEG solutions containing 5mol% DEA at 30, 50, and 70°C
43
0.00 0.20 0.40 0.60 0.80 y [mol C02/mol amine]
Figure 4.7 Solubility of C02 in TEG solutions containing 10mol% DEA at 30 and 50°C
44
0.00 0.20 0.40 0.60 0.80 y [mol C02/raol amine]
Figure 4.8 Solubility of C02 in TEG solutions containing 13.6mol% DEA at 30, 50, and 70°C
45
IO^P
E 4-J
m
u Q.
5mol% MEA
0.01 -:
0.001 i i i
0.20 y [mol C02/mol amine]
Figure 4.9 Solubility of C02 in DEG solutions containing 5mol% MEA and 10mol% MEA at 40°C
46
4.2 pH MEASUREMENTS
4.2.1 AQUEOUS MDEA SOLUTIONS
Results of the pH measurements in aqueous solutions of 4M MDEA
at 30, 40, 50 and 60°C, given in Appendix C, are presented in
Fig. 4.10. The linear lines are obtained from a semiempirical
formula established in section 5.2.3.
x a.
10.5
10.0
9.5
9.0
8.5
8.0
7.5
-
-
-
-
-
: >
^
i
i
i
^^^^\^°
- - ^ > ^
-
-
-
0.1 10 100 (l-y)/y
Figure 4.10 Experimental pH data for aqueous 4.00M MDEA solutions at 30, 40, 50, and 60°C. The lines follow Eqn. (5.18)
47
4.2.2 AQUEOUS AMP SOLUTIONS
Results of the pH measurements in aqueous solutions of 3M AMP at
20, 30, 40, and 50°C are summarized in Appendix C and are
presented in Fig. 4.11 together with linear lines obtained from
a semiempirical relation established in section 6.2.3.
Figure 4.11 Experimental pH data for aqueous 3.00M AMP solutions at 20, 30, 40, and 50°C. The lines follow Eqn. (6.21)
48
4.3 KINETIC MEASOKEMENTS
The data here presented have emerged from some screening
measurements undertaken to investigate the influence of the
solvent on the absorption rate of C02 into MEA solutions. Water,
triethyleneglycol (TEG), n-methyl-pyrrolidone (NMP), ethanol
(EtOH), and diethyleneglycol monomethylether (DEGMME) were
investigated as solvents. The results are tabulated in Appendix
D. In Fig. 4.12 and 4.13 the absorption rates of C02 are given
as a function of the wetting rate (see for example Morris and
Jackson (1953)), for each of the solvents, and for 5mol% MEA
solutions of the different solvents, respectively.
The string-of-discs column, on which the experiments were
performed, have such operating characteristics that the best
range for comparing absorption rates is at wetting rates between
0.4 and 0.5 cm3/cm s. The column must have a high enough wetting
rate to ensure complete wetting of the discs, and the wetting
rate must be kept below the point where ripples are formed on the
surface.
The results of these experiments are to be used with caution,
since as remarked in section 3.3.1, there was a pronounced
increase in the temperature of the solution along the column, see
Appendix D. A more complete analysis should have taken this into
account, when the different solvents are compared. Fig. 4.13
indicates that the absorption rates are highest for aqueous and
alcoholic MEA solutions. Also, a comparison between Figs. 4.12
and 4.13 indicates that the addition of MEA has the most effect
on aqueous solution.
49
WATER
'." I—
0 . 2 5 0 . 5 0 !*• (cm 3 /cm a)
Figure 4.12 Rate of absorption of C02 in the physical solvents: water, n-methyl-pyrrolidone, ethanol, triethyleneglycol, and diethyleneglycol monome thy lether, as a function of wetting rate at 20°C
50
I-' I cm 1/rin g j
Figure 4.13 Rate of absorption of C02 in the physical solvents: water, n-methyl-pyrrolidone, ethanol, triethyleneglycol, and diethyleneglycol monomethylether, containing 5mol% MEA, as a function of wetting rate at 20 °C. The temperatures of the solvents out of the absorption column, are given in Appendix D
51
Chapter Five
A Model for Equilibrium Solubility of Carbon Dioxide
in Aqueous Solutions of the Tertiary Amine MDEA
In this chapter, a serai empirical gas-liquid equilibrium model for
C02 in aqueous methyldiethanolamine (MDEA) solutions, is
presented. The equilibrium model is based on experimental
solubility and pH determinations. It gives the equilibrium
partial pressure of C02 as a function of three variables: the
amine concentration, the C02 loading, and the temperature.
In section 5.2, a model based on experimental data obtained
solely in our laboratory, is presented. Both equilibrium and pH
measurements were undertaken at temperatures between 30 and 60°C.
In section 5.3, a model based on the same pH measurements and a
more comprehensive set of equilibrium data from the literature
with temperatures up to 120°C (Jou et al. (1982)), is presented.
In this case one have succeeded in modelling the partial pressure
of C02 over a range of seven decades, the C02 loading over more
than three decades and covering a temperature range between 25
and 140°C with very good accuracy. The model is shown to be
accurate for amine molarities between 1.69 and 4.28M.
5.1 INTRODUCTION
The detailed chemistry of CO2 absorption in tertiary amine
solutions is discussed in section 2.2, and we will in this
chapter restrict ourselves to comment on those parts of the
chemistry having direct influence on the present modelling
52
procedure.
In aqueous solutions of tertiary amines, the overall reaction to
take place with C02 is the bicarbonate formation, Eqn. (5.1).
C02 + R2NCH3 + H20 = R2NHCH3+ + HC03~ (5.1)
Carbamate formation does not occur in the case of tertiary amines
like MDEA, because the MDEA molecule does not have a hydrogen
atom attached to the nitrogen atons. This leads to a slower
absorption rate than for primary and secondary amines, since the
bicarbonate reaction, Eqn. (5.1), occurs relatively slowly. In
order to speed up the reaction, activating agents can be added
to the MDEA solution. Such agents are for instance other amines
with higher reaction rates. Investigations on the kinetics of C02
absorption in activated MDEA solutions have been made by Xu et
al. (1992).
With H2S, MDEA will react directly following the same fast
reaction mechanism as for primary and secondary amines. This is
the reason why MDEA as a tertiary amine is heavily used in
selective absorption of H2S when C02 is present. However, MDEA is
also well suited for bulk C02 removal, see Bullin et al. (1990).
One obvious reason for this is that MDEA can be used in high
concentrations up to 50wt%, in combination with high C02
loadings. Also the loss of amine using MDEA solutions will be
small, due to low vapor pressure and slow degradation rates. MDEA
solutions are also less corrosive than other amines such as MEA
and DEA (Bullin et al. (1990)).
For tertiary alkanolamines, the alcoholic group will also have
some reactivity with C02, but according to Yu et al. (1985) the
reactivity of the alcoholic groups in MDEA will be small compared
53
to the reactivity of the amino group at the pH levels of
interest. Thus, in developing the model, only the reactivity of
the amino group is taken into account.
A number of applicable models, describing the gas-liquid equili
bria of C02 in alkanolamines, have been presented in the
literature over the years, such as Danckwerts and McNeil (1967),
Klyamex and Kolesnikova (1972), Kent and Eisenberg (1976),
Deshmukh and Mather (1981 ), Chakma and Meisen (1987), and Austgen
et al. (1989), all of which are discussed in section 2.1.2. Our
approach differs from the previously developed models, in that
we apply measured pH data to describe the effect of temperature
and loading. This allows an estimation of the relation between
the equilibrium partial pressure of CO2 and the solution loading,
without the necessity of knowing several equilibrium constants,
the Henry's law coefficient, the activity coefficients, or
interaction parameters.
5.2 C02 EQUILIBRIUM MODEL FOR AQUEOUS 4M MDEA
5.2.1 APPROXIMATIONS
For all values of y and T encountered in this investigation, the
concentration of free C02(aq) has been assumed to be negligible
in comparison with the HC03~ concentration. The reasons for this
are the low solubility (high Hc02-value) of CO2 as such in
aqueous solution, and the moderate partial pressures of CO2
covered in this study. Because of the relatively low basicity of
MDEA, one can also neglect the formation of the carbonate ion at
all but extremely low values of y (Yu et al. (1985)). Such low
y-values are often outside the region of interest in actual CO2
absorption processes. This leaves HCO3" as the only main C02-
54
source of the liquid phase. With m denoting the amine molarity
and y the C02 loading, the following relations arise:
m*y = CHC03" ( 5' 2 )
m = CR2NCH3 + CR2NHCH3+ (5.3)
The electroneutrality requirement gives:
CR2NHCH3+ = CHC03" = m'V < 5 - 4 )
E q n s . ( 5 . 3 ) and ( 5 . 4 ) c o m b i n e d g i v e :
CR2NCH3 = m ( 1 - y ) ( 5 - 5 )
5.2.2 THE BASIC MODEL
Combining the first dissociation constant of carbonic acid:
K1 = <aHC03-,aH+)/aCO2 = aH+,<fHC03-/fC02)*(CHC03-/CC02) <5'6)
and H e n r y ' s l a w i n t h e f o r m :
PC02 = H * CC02 < 5 - 7 )
and i n t r o d u c i n g Eqn. ( 5 . 4 ) , we g e t :
PC02 = K«aH+-m»y ( 5 . 8 )
w h e r e
K = ( f H C 0 3 - , H > / < f C 0 2 - K 1 > < 5 ' 9 >
I n t r o d u c i n g pH i n s t e a d o f a H +, Eqn. ( 5 . 8 ) c a n b e r e f o r m u l a t e d a s
f o l l o w s :
PC02 = " • v i " 1 1 0 9 ' ' " p H ) ( 5 .10 )
5.2.3 A CORRELATION FOR pH
We start with the expression for the amine protonation constant
on activity basis:
Kp = aR2NCH3*aH+/aR2NHCH3+ ( 5 . 1 1 )
55
I n t r o d u c i n g a = f « C , and s o l v i n g f o r aH+, we g e t :
aH+ = Kp • (CR2NHCH3+/CR2NCH3)
w h e r e
Kp' = Kp» (fR2NHCH3+' fR2NCH3*
Now m a k i n g u s e o f E q n s . ( 5 . 4 ) and ( 5 . 5 ) , Eqn . ( 5 . 1 2 ) c a n b e
r e w r i t t e n i n t h e f o l l o w i n g form:
pH = pKp» + l o g [ ( 1 - y ) / y ] ( 5 . 1 4 )
For 4.00M MDEA, measured pH values are plotted against
log[(1-y)/y] in Fig. 4.10. The figure shows a linear
relationship between pH and log[(1-y)/y] as expected from Eqn.
(5.14), but rather of the form:
PH = DKp' -.- b-loaH1-v)/vl (5.15)
where b is a constant, independent of the temperature.
Fig. 4.10 allows an estimate of the numerical value of b for the
temperature range investigated. From the slope of the parallel
lines we find:
b = 0.88 (5.16)
Also from Fig. 4.10, pK_' can be estimated. Our analysis of the
data shows that pK_' can be regarded as independent of the
loading y, closely following the correlation:
pKp' = 13.38 - 0.0154'T (5.17)
The accuracy of the relationship expressed in Eqn. (5.17) is very
good, as can be seen from Fig. 5.1 where each data point is the
mean value of pK_' calculated for several y-values in the range,
y = 0.015 - 0.67, investigated.
(5.12)
(5.13)
56
Eqns. (5.15)-(5.17) yield pH for all buffer compositions covered
in this investigation over the temperature range investigated:
pH = 13.38 - 0.0154'T + 0.88»loqf(1-v)/vl (5.18)
Our experience is that such pH measurements are quite demanding.
However, we now have at our disposal a model covering the T,y-
ranges of most interest in C02 absorption, and it should not be
necessary to undertake new pH measurements for the given amine
concentration (4M MDEA).
9.0 r
8.8 -
8.6 -
n
a.
8.-1 -
8.2 -
8.0 I L
303 313 323 333
T IK1 Figure 5.1 pKp' as a function of the temperature for aqueous
4.00M MDEA solution
57
5.2.4 A CORRELATION FOR logK
Introducing logarithms, Eqn. (5.10) can be rewritten as follows
logK = logpc02 - log(m«y) + pH (5.19)
Now, making use of the pH-model, Eqn. (5.18), and introducing
experimental gas-liquid equilibrium data, we find:
loqK = 2.18 + 0.0188'T (5.20)
logK is here regarded as being independent of the loading y. As
can be seen from Fig. 5.2, where each data point is the mean
value of logK calculated for several y values at five different
temperatures, the linear relationship in Eqn. (5.20) gives a good
description of the experimentally derived data.
B.6
8.4
8.2 I
8.0
7.8
7.fi
3U3 313 323 333 TIKI
Figure 5.2 logK as a function of the temperature for aqueous 4.00M MDEA solution
58
J I i L
5.2.5 A PRELIMINARY FINAL MODEL
Combining Eqns. (5.10), (5.18), and (5.20) gives a preliminary
final model:
PC0. = m»v10'-dtc-T-b-lo9ni-Y)/Yl) (5.21)
where b = 0.88 c = 0.0342 d = 11.20
and T is the absolute temperature in K.
We have found that the value of logK starts to show a minor
dependence on y at low loadings (y<0.15). By taking into account
equilibrium data from the literature, we were able to cover
several decades of C02 concentrations. It was then found that by
adding a y-term in Eqn. (5.20), a satisfactory description of the
logK-expression as a function of both T and y, even for low y
values, could be established. This is further described in
section 5.3.
5.2.6 COMPARISON WITH EXPERIMENTAL EQUILIBRIUM DATA
In Fig. 4.1, equilibrium curves derived from the model are given
for 4M MDEA at 30, 45 and 60°C, together with experimental values
for the same temperatures. There exists good agreement between
model and actual equilibrium data. In Fig. 4.2, modelled
equilibrium curves for 4.28M MDEA, are compared with experimental
data from Jou et al. (1982) and own experimental data at 25, 40,
and 70°C. Considering that the model in section 5.2.5 is based
on data for 4M MDEA, the agreement is seen to be very good even
at relatively low loadings, y<0.4. For higher y-values, the model
overpredicts the equilibrium partial pressures.
59
The MDEA equilibrium model here presented, is simple and easy to
use, since it gives the equilibrium partial pressure of C02 as
an explicit function of only two central and easily determined
process variables, y and T.
5.3 EXTENDED EQUILIBRIUM MODEL, VALID FOR AQUEOUS SOLUTIONS WITH
1-4.5M MDEA AT TEMPERATURES BETWEEN 25 AND 140°C
5.3.1 INTRODUCING VLE DATA FROM THE LITERATURE
The model presented in the previous section is applicable in the
relatively low temperature range (25-70°C), which is encountered
in absorption columns. To be able to simulate the whole
absorption/stripping-process, one has to include higher
temperatures. Equilibrium data, covering the temperature range
25-120°C, are given in the literature (Jou et al. (1982)) for
4.28M MDEA. Combining these data with the pH data presented in
section 4.2.1 for a 4.00M solution, an equilibrium model emerges,
covering partial pressures from 0.00001 to 50 atm. This model,
which can be used over a temperature interval of more than 100°C,
needs one parameter for the amine system, and one parameter for
the C02 system, i.e. a total of two parameters, which is the same
as for the more restricted model given in section 5.2. We shall
now proceed with determining these two parameters from
experimental VLE- and pH-data. The model is tested against
experimental values and found to be accurate at amine molarities
ranging from 1.69 to 4.28M.
5.3.2 A NEW CORRELATION FOR THE PARAMETER K
The basic model and the correlation for pH are assumed identical
60
with the expressions presented for the restricted model in
section 5.2. However, as stated in section 5.2.5, logK starts to
show a dependency on the CO2 loading as the y interval is
broadened. We again start with Eqn. (5.19):
logK = logpC02 - log(m«y) + pH (5.19)
Making use of the pH-model from Eqn. (5.18) and introducing the
complete sets of experimental VLE data from Jou et al. (1982),
give:
loqK = 4.78 + 0.0094-T + 0.29-logr(1-v)/vl (5.22)
5.3.3 THE FINAL MODEL
Combining Eqns. (5.10), (5.18), and (5.22) gives the final model:
p c 0, = m.v.io(-d • c-T-b-log[(1-y)/yI) { 5 - 2 3 )
which is the same as Eqn. (5.21). However, the parameters were
found to assume new values:
b = 0.59 c = 0.0248 d = 8.60
T is the absolute temperature in K.
As shown below, Eqn. (5.23) with the given parameter values has
been found valid for all loadings less than 1 .0 and amine
molarities between 1.69 and 4.28M. It is tested and found
accurate at temperatures between 25 and 140°C, see Figs. 5.3 and
5.4. The model equation can easily be programmed by a scientific
calculator, such as HP-42S. The listing of a program for this
purpose, is given in Appendix E.
61
5.3.4 COMPARISON WITH EXPERIMENTAL EQUILIBRIUM DATA
Modelled equilibrium curves from Eqn. (5.23) together with given
parameter values are compared with experimental data for 4.28M
MDEA at 25, 40, 70, 100, 120, and 140°C in Fig. 5.3, 5.4, and
5.5. Both literature data (Jou et al. (1982), Chakma and Meisen
(1987), and Austgen (1989)), as well as own experimental data,
show good agreement with the model. However, at low loadings
(y<0.01) at the lowest temperature, 25°C, the model differs
somewhat from data given by Jou et al. (1982), see Fig. 5.3. This
is outside the region of interest in most acid gas treating
processes.
The model is based on equilibrium data for 4.28M and pH data for
4.00M solutions. Comparison with experimental data at other amine
concentrations shows that the model can be used at a wide range
of amine molarities. In Fig. 5.6 and 5.7 equilibrium curves
derived from the model for 2.00M MDEA at 25, 40, 70, 100, and
120°C, shows good agreement with experimental values from Jou et
al. (1982) and Austgen (1989), except for the 25°C-curve where
considerable deviation from experimental data occurs at loadings
below 0.05 mol/mol. Also the 40cC-curve shows deviations for some
data points at low loading. In Fig. 5.8 equilibrium curves
derived from the model for 3.04M MDEA at 40 and 100°C, are
compared with experimental values from Jou et al. (1986). A nice
agreement can be seen at 40°C, while the model underpredicts the
partial pressure somewhat at 100°C. A comparison, showing a good
description of the experimental data for 1.69M MDEA at 100°C,
collected from Chakma and Meisen (1987), is given in Fig. 5.9.
Measurements undertaken in this work for 4M MDEA at 30 °C,
presented in section 4.1.1, are compared with the modelled
equilibrium curve in Fig. 5.10. The agreement can be seen to be
62
acceptable, but all experimental values lie above the modelled
It should be emphasized, that all the presented equilibrium
curves have emerged from pH data undertaken for 4M solutions,
only. Despite of this, and quite surprisingly, the model exhibits
good agreement for a broad range of amine concentrations. This
could be partly explained by the fact that pH is related to COj
loading, and not C0 2 concentration, in the pH-expression
established in Eqn. (5.18). The influence of the amine
concentration on the pH values, is thereby reduced. Furthermore,
the sound principles on which the model is built, is believed to
contribute to the good agreement achieved.
I'-Df
10-d
O
^i
O.i -
0.01 i
0.001 i
0.0001 -s
0.00001
Figure 5.3
0.0001 0.001 0.01 0.1 1 y (mol C02/mol MDEA)
Comparison of the present model (solid lines) with experimental data from the literature (Jou et al. (1982)) on the system of 4.28M MDEA aqueous solution at 25, 40, 70, 100, and 120°C
63
100
E - 4 - *
o CM O O
0.01
0.001 -
0.0001 i
0.00001 "T i i i i r~p
0.1 y (mol C 0 2 / m o l MDEA)
Figure 5-4 Comparison of the present model (solid line) with experimental data from the literature (Chakma and Meisen (1987)) on the system of 4.28M MDEA aqueous solution at 140°C
64
100 -a
c o
O '_> a
0.00001
0.01 -
0.001 -2
0.0001 •=
y (mol C02/mol MDEA)
Figure 5.5 Comparison of the present model (solid line) with present experimental data ( A ) and data taken from the literature (Jou et al. (1982) ( * ) and Austgen (1989) ( Q )) on the system of 4.28M MDEA aqueous solution at 40°C
65
E o
O O
100 3-
10 =
1 =
0.1 =
0.01 ^
0.001 =
0.0001 -
0.00001 0.001
-T 1 — I t I I I I 1 1 — I — I I I I I t ~l 1 — I — I I I I I
0.01 0.1 y (mol C02/mol MDEA)
Figure 5.5 Comparison of the present model (solid lines) with experimental data from the literature (Jou et al. (1982)) on the system of 2.00M MDEA aqueous solution at 25, 40, 70, 100, and 120°C
66
£ "a
CJ O o
100 -a
1 0 -
0.1 -
0.01 =
0.001 -
0.0001 =
0.00001 0.001 0.01 0.1
y (mol C02 /mo l MDEA)
Figure 5.7 Comparison of the present model (solid line) with experimental data from the literature (Jou et al. (1982) ( * ) and Austgen (1989) ( D )) on the system of 2.00M MDEA aqueous solution at 40°C
67
CM O o
1 0 0 *
1 0 =
0.1 =
0.01 =
0.001 =
0.0001 =
0 . 0 0 0 0 1 ~v* 1 — i — i — i i i 111 1 — i — i — I ' I i 111 ; — i — i — i i 11 i
0.001 0.01 0.1 y (mol C02/mol MDEA)
F i g u r e 5 .8 Comparison of the p r e s e n t model ( s o l i d l i n e s ) wi th exper imenta l data from the l i t e r a t u r e (Jou e t a l . (1986) ) on the system o f 3.04M MDEA aqueous s o l u t i o n a t 40 and 100°C
68
O o
100 f
10-
0.1 -
0.01 -;
0.001 =
0.0001 i
0.00001
y (mol C02/mol MDEA)
Figure 5.9 Comparison of the present model (solid line) with experimental data from the literature (Chakma and Meisen (1987)) on the system of 1 .69M MDEA aqueous solution at 100°C
69
100
10 =
E a
CM O O
0.1 =
0.01 -
0.001 z
0.0001
0.00001 ~t 1—i i 111 m—'-i—i i i 11 m 1—i i 11 nn 1—i i 11 ui 0.0001 0.001 0.01 0.1
y (mol C02 /mo l MDEA)
Figure 5.10 Comparison of the p resen t model ( s o l i d l i n e ) with p r e s e n t exper imental da ta for aqueous s o l u t i o n s of 4.00M MDEA a t 30°C
70
5.4 ACCURACY OF THE MODEL
The present equilibrium models are based on pH measurements. The
accuracy of the prediction the model gives, is therefore very
much dependent on the accuracy of these pH measurements. As an
example, uncertainties in the estimation of pH on ±0.02, ±0.05,
and ±0.1, would result in uncertainties in the predicted C02
partial pressure of 4.5%, 10.9%, and 20.6%, respectively. This
shows the importance of having reliable pH data available. We
believe to have established a pH measurement procedure which
yields consistent data with high accuracy.
The largest uncertainty is regarded to be the C02-analysis
related to the liquid phase.
5.5 CONCLUSIONS
For an aqueous solution of the tertiary amine methyl-
diethanolamine, a semiempirical thermodynamic approach has been
developed to model the relation between the equilibrium partial
pressure of C02, the C02 loading, the absolute temperature, and
the amine molarity.
It is demonstrated that the model fits experimental data very
well. The model shows excellent agreement with experimental data
at all temperatures between 25 and 140°C at C02 loadings and
amine molarities usually encountered in acid gas treating plants.
71
Chapter Six
A Model for Equilibrium Solubility of
Carbon Dioxide in an Aqueous Solution
of the Sterically Hindered Amine AMP
Following the same lines as described for tertiary amines in
Chapter 5, a semiempirical gas-liquid equilibrium model for C02
in aqueous 3M AMP (2-amino-2-methyl-1-propanol), is presented.
It applies to high CO2 loadings (y>0.5) in the temperature range
between 20 and 50°C, and is based on experimental solubility and
pH determinations. For a given amine concentration, it yields the
equilibrium partial pressure of CO2 as a function of only two
variables: the C02 loading and the temperature.
6.1 INTRODUCTION
The growing interest in aqueous solutions of sterically hindered
amines for the use in acid gas treating processes is due to their
high cyclic capacity, and relatively high absorption rates at
high CO2 loadings (Sartori and Savage (1983)).
The primary amine, 2-amino-2-methyl-1-propanol (AMP), is regarded
as sterically hindered because the amino group is attached to a
tertiary carbon atom. Aqueous solutions of AMP show low carbamate
stability, and therefore larger cyclic capacity may be obtained
than for conventional amines such as MEA. AMP is used here to
demonstrate that the same new C02 VLE modelling technique as
demonstrated for aqueous MDEA solutions, is applicable also for
a hindered amine solution.
72
The main reactions between C02 and primary amines are earlier
presented in section 2.2, and are here briefly recapitulated to
give basis for the modelling procedure.
The bicarbonate formation reaction (BF) occurs during absorption
of C02 in primary, secondary and tertiary amine solutions. We
have taken a primary amine as an example.
BF: C02 + RNH2 + H20 = RNH3+ + HCO3" (6.1)
In the case of primary and secondary amines, carbamate formation
(CF), and carbamate reversion (CR), also need to be considered.
CF: C02 + 2RNH2 = RNH3+ + RNHCOO- (6.2)
CR: RNHCOO" + H20 = RNH2 + HC03" (6.3)
Reactions (6.2) and (6.3) are governed by the carbamate stability
constant, Kc, and the amine protonation constant, K-:
Kc = CRNHC0O~/(CRNH2"CHCO3~) (6.4)
Kp = (CRNH2-CH+)/CRNH3+ (6.5)
When Kc is very small, one can neglect carbamate formation, and
only consider the bicarbonate formation, Egn. (6.1).
Kc for AMP was reported by Sartori and Savage (1983) as less than
0.1 1/mol at 40°C. This low value implies that the BF-reaction
is often predominant at absorption conditions.
In section 2.2.2, it is reported that a possible mechanism of
bicarbonate formation is given by Astarita et al. (1983). The
73
mechanism proposed goes via the formation of an intermediate
"zwitterion" which reacts with water more easily to form
bicarbonate:
C02 + RNH2 = RN+H2COO~ (6.6)
RN+H2COO" + H20 = RNH3+ + HCO3" (6.7)
This zwitterion path applied in the case of sterically hindered
amines, is an extension of the reaction mechanism for carbamate
formation proposed by Caplow (1968) and later supported by
Danckwerts (1979). Yin and Shen (1988) investigated the kinetics
of the C02 reaction in an AMP-solution, and ended up with this
mechanism. Bosch et al. (1990) also propose a zwitterion
mechanism, but they point out that the zwitterion will react with
all bases present in the solution, not only the water. This
reaction pattern is not included in the present model. Further
discussion of the chemistry of C02 absorption in sterically
hindered amine solutions is given in section 2.2.4.
According to Sartori and Savage (1983), the relatively high
absorption rates for hindered amines are due to the low carbamate
stability, which leads to high free amine concentration. The
absorption rates are in many cases appreciably higher than those
for conventional amines, at least at high C02 loadings. This
happens despite that there will generally be some reduction of
the rate constant due to steric interference.
74
6.2 THE EOUILIBRIOM MODEL FOR CQ2
6.2.1 APPROXIMATIONS
At the pH levels encountered in this investigation, in the
temperature range 20 to 50°C, the concentration of free C02(aq)
is negligible in comparison to the HC03" concentration. Also,
except for very low C02 loadings, the concentration of C032~ can
be neglected compared to HC03~ (Sartori and Savage (1983)). This
leaves HC03" as the dominating C02-compound in the liquid phase.
With m denoting the amine molarity and y the C0 2 loading,
assuming that the carbamate concentration is zero, we obtain the
following relationships:
m*y = CHC03" < 6- 8 )
m = CRNH2 + CRNH3+ (6.9)
Electrical neutrality requires that:
CRNH3 + = CHC03" = «n-y (6.10)
Combination of Eqns. (6.9) and (6.10) yields:
CRNH2 = n»<1-y> (6.11)
Eqns. (6.8)-(6.11) are equivalent to Eqns. (5.2)-(5.5), with the
AMP compounds taking the place of the MDEA compounds.
6.2.2 THE BASIC MODEL
As a concequence of the similarity in the assumptions introduced
regarding the concentration of the participating C02 compounds,
the development of the VLE formula can follow the same procedure
as for MDEA. Combining the first thermodynamic dissociation
constant of carbonic acid:
75
Kl = (aHC03" * aH+ > /aC02 = aH+ ' < fHC03~/ f C02 > * <CH C 0 3- /CC 0 2) ( 6 . 1 2 )
and H e n r y ' s l aw i n t h e fo rm:
PC02 ~ H * Co^ (6.13)
and introducing Eqn. (6.10), we obtain:
PC02 = K«aH+-m-y (6.14)
where
K = tfHC03--H>/<£co2'Ki> (6.15)
On introducing pH instead of aH+, Eqn. (6.14) can be reformulated
as follows:
pCQ, = m-v10' lo9K ~ PH> (6.16)
6.2.3 A CORRELATION FOR pH
We s t a r t with an expression for the amine protonation constant,
based on ac t iv i ty :
Kp = aRNH2,aH+/aRNH3+ ( 6 . 1 7 )
Introducing a = f-C and solving for aH+, we obtain:
aH+ = Kp'-fCRNHj-f/CRNHj) (6.18)
where
V = V(fRNH3^fRNH2) (6.19)
Now making use of Eqns. (6.10) and (6.11), Eqn. (6.18) can be
rewritten in the following form:
pH = pKp' + log((1-y)/y) (6.20)
pH values here measured are plotted against log((1-y)/y) in Fig.
4.11. The diagram shows a linear relationship between pH and
log((1-y)/y) as expected from Eqn. (6.20), but rather of the
form:
pH = PKp' + h'loa(n-vWv) (6.21)
76
where b is a constant, independent of temperature.
Fig. 4.11 allows the numerical value of b to be estimated for the
investigated temperature range, viz.,
b = 1.31 (6.22)
pKp' can also be estimated from Fig. 4.11. Our analysis of the
data shows that pKp' is independent of the loading y, and closely
follows the correlation:
pKp' = 1950/T + 3.13 (6.23)
The accuracy of this relationship is brought out by Fig. 6.1
where each point represents the mean value of pK_' for several
y-values in the range 0.47 - 0.90 investigated. For such a
restricted temperature range, using T, as used for MDEA, or using
1/T, does not have much influence on the accuracy of the model.
9.8
9.6
a.
9.4
9.2
3.1-10 3 3.2-II)3 3.3-10 3 3.4- IO'3
1/rlK'l
Figure 6.1 pKp' as a function of 1/T for aqueous 3.00M AMP solution
77
6.2.4 A CORRELATION FOR logK
Taking logarithms, Eqn. (6.14) can be rewritten as follows:
logK = logpco2 - log(m«y) + pH (6.24)
On introducing experimental gas-liquid equilibrium data from this
work and also from Komorowicz and Erga (1987), we found that:
logK = 7.17 + 0.52-v (6.25)
irrespective of T. This is in contrast to what was found in the
case of MDEA, where logK was found to depend on the temperature.
The linear relationship for logK is presented in Fig. 6.2. The
data points cover the temperature range 20 - 50°C.
We have observed that, above 50°C, logK is somewhat affected by
the temperature.
0.4 0.6 0.8
y Imol C02/mol AMPl
1.0
Figure 6.2 logK as a function of C02 loading for aqueous 3.00M AMP solution
78
6 . 2 . 5 THE FINAL MODEL
Combination o f Eqns . ( 6 . 1 6 ) , ( 6 . 2 1 ) , ( 6 . 2 2 ) , ( 6 . 2 3 ) and ( 6 . 2 5 )
y i e l d s t h e f i n a l model:
p c 0 2 = m - v 1 0 < c * d*y - e / T ~ b-log[(1-y)/y1) ( 6 . 2 6 )
where b = 1.31 c = 4 . 0 4 d = 0 .52 e = 1950
and T is the absolute temperature in K.
6.2.6 COMPARISON WITH EXPERIMENTAL EQUILIBRIUM DATA
Fig. 4.3 presents the equilibrium curves derived from the model
for 3.0M AMP at 20, 40, and 50°C, together with experimental
values for 40°C from Sartori and Savage (1983), data for 20 and
40°C from Komorowicz and Erga (1987) and own experimental data
for 40 and 50°C. The agreement between the model and the actual
data is seen to be very good.
It should be noted that the absorption of C02 in acid gas
treating plants, will often take place inside the temperature
range covered by the model.
The advantage of the present equilibrium model is the same as for
the MDEA-model, and lies in its simplicity: It yields Pco2 a s a n
explicit function of only two variables, y and T, which both are
easy to measure. The modelling is based on an analysis of
measured pH and gas-liquid equilibrium data, similar to the
modelling of the C02/MDEA system in Chapter 5, and in much the
same way as earlier achieved for aqueous S02 solutions buffered
with citrate and adipate ions (Erga (1980, 1986)).
79
6.2.7 LIMITATIONS
The present model has certain limitations compared to the more
comprehensive model developed for the MDEA-system, the most
important being that it does not cover stripping conditions- The
difficulty in modelling the AMP-system all the way up to
stripping temperatures, may be due to the fact that the C02
reactions with sterically hindered amines are more complex and
less understood than those with tertiary amines (Bosch et al.
(1990)).
6.3 CONCLUSIONS
A semiempirical thermodynamic model has been developed which
describes the equilibrium partial pressure of C02 as a function
of only the CO2 loading and the absolute temperature/ for a given
concentration of the sterically hindered amine, 2-amino-2-methyl-
1-propanol.
Equilibrium curves derived from this model compare very well with
the equilibrium data found in the literature. The model is at
present restricted to the low temperature range of 20 - 50°C and
to high loadings, y = 0.50 - 0.95. However, these are conditions
which are often encountered in CO2 absorption units.
80
Chapter Seven
Vapor-Liquid Equilibria of Mixed Nonaqueous Solvents
In this chapter we -will discuss some aspects of the VLE
measurements reported in section 4.1.3, where the TEG/MEA,
TEG/DEA, and DEG/MEA systems were investigated. The measurements
were undertaken to compare the solubility of CO2 in glycol-amine
solutions with the solubility in the more frequently used
solvents such as aqueous alkanolamine solutions. An equilibrium
model has been developed for predicting CO2 VLE data at elevated
temperatures and at low loadings. Due to difficulties in
obtaining reliable experimental data at these conditions, it is
important to have a predictive model based on sound principles.
7.1 EQUILIBRIUM SOLUBILITY MODEL FOR CO-i IN TEG/MEA SOLUTIONS
7.1.1 BACKGROUND
For the TEG/MEA system, experimental equilibrium data at stripper
and lean end absorber conditions are scarce, and a predictive
model based on available data is desired. A model is here
presented which describes the measured experimental data very
well. It is believed that this model might even be useful for
predicting the equilibria outside the experimentally investigated
ranges of temperature and partial pressure.
The objective was to establish a model equation describing the
C02 equilibria at the following conditions: C02 loadings between
0.005 and 0.45, temperatures between 30 and 150°C, and amine
81
concentrations between 0.60 and 1.0M. The following modelling
procedure is based on the experimental equilibrium data given in
section 4.1.3 for 10mol% MEA (0.79M) in TEG.
7.1.2 MODELLING PROCEDURE
Comprehensive VLE data for aqueous MDEA system obtained by Jou
et al. (1982) and Chakma and Meisen (1987), indicate that a plot
of equilibrium partial pressure of CO2 against 1/T, at constant
loading, yields a linear relationship. Such a plot is given in
Fig. 7.1. for the aqueous MDEA system for temperatures in the
range 25 to 140°C (298-413K), and for loadings in the range 0.004
to 0.5 mol /mol. As can be seen, the curves are approximately
parallel. Assuming that the nonaqueous system has a similar
behavior, we have a convenient way of extrapolating existing
data.
82
100 i
e •>-> n
Q.
0.00001
0.01 =
0.001 =
0.0001 =
0.5 0.4 0.3
0.2
0.1
0.04
0.01
y
i i 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 M i M r 1 1 i i 1 1 1 1 M 1 1 1 1 1 1 1 1 1 1 1 1
2.40 2.60 2.80 3.00 3.20 3.40 1 0 0 0 / T [ 1 0 0 0 / K ]
0.004
Figure 7.1 Equilibrium partial pressure of C02 presented as a function of 1000/T for eight different C02 loadings in aqueous solutions of 4.28M MDEA. The data are taken from figures presented by Jou et al. (1982) and Chakma and Meisen (1987)
83
•A
1 ;
0.1 1
0.01 -=
"
0.001 i
0.0001 -
"».. \ ^ v ...
% -• N
X x X N
X X
X v. X X N
\ X
—i—i—i—n—i—i—i—i—i—i—i—r~
x X
- X -v \ \ X
x v X X 3v X.
X x ^ X x x \ \Xx X xxx
X X X * X. X \ , X^
T-i—i—i—i i i i i—i i i f j—i—i—i
0.4
0.3
0.2
0.1
2.20 2.60 3.00 X ^ . 4 0 1000/T [1000/K] X 0.05
Figure 7.2 Equilibrium partial pressure of C02 presented as a function of 1000/T for five different C02
loadings in a solution of TEG and 10mol% MEA
Fig. 7.2 shows the same plot for the TEG/MEA system. The
corresponding data points of partial pressure and inverse
temperature for different loadings, are obtained by smoothing the
data given in Table 5, Appendix B, and presented in Fig 4.5. The
approximate parallel lines suggest that the assumption of
linearity, is valid. This implies that the equilibrium curves for
a certain loading may be described with an equation of the form:
p c 0 2 = exp(A • 1000/T) • B (7.1)
Values for A and B are obtained for several different loadings,
ranging from 0.005 to 0.4 mol C02/mol MEA. The parameter B
follows closely an equation of the form:
84
InB = C • ln(y/1-2y)2 + D (7.2)
This implies that the equilibrium partial pressure is
proportional to (y/l-2y)2. For MEA, which is a primary amine
forming a stable carbamate, the same dependence is known to apply
also for aqueous systems (Astarita et al. (1983) and Sartori and
Savage (1983)).
The parameter A can be regarded as being constant irrespective
of the loading:
1000'A = E (7.3)
The coefficients C, D, and E are adjusted to obtain a best
possible fit to actual experimental data. The final equation,
describing the C02 equilibria in a TEG solution with 10mol% MEA,
then emerges:
PCQ2 ~ expfc ' ln(v/l-2y)2 4 d - e/Tl (7.4)
where c = 0.57 d = 23.88 e = 8570
and T is the absolute temperature in K.
7.1.3 COMPARISON WITH EXPERIMENTAL EQUILIBRIUM DATA
Literature data for comparison are to our knowledge not available
for this particular system. Fig. 7.3 presents the equilibrium
curves derived from the model at 30, 50, 70, 100, and 150°C,
together with present equilibrium data for temperatures up to
70°C.
85
e JJ
•.'•J -a
3 -i
i -
0 . 1 •=
£ 0.01 i
0.001 i
0.0001 -
0.00001 - i — i i r i i i 1 1 — i — r i i i i
0.01 0.1 y [mol C02/mol amine]
i 1 1—i—n
Figure 7.3 Comparison of the present model (solid lines) with present experimental data for a solution of TEG and 10mol% MEA at 30, 50, and 70°C, and predicted equilibrium curves for 100 and 150°C
7.2 COMPARISON WITH AQUEOUS AMINE SOLUTIONS
Fig. 7.4 compares the equilibrium curves for the TEG/MEA and
DEG/MEA solution with an aqueous MEA solution (Lee et al.
(1976)). All solutions contain 5mol% MEA. The amine
concentrations in the solutions are as follows: TEG-0.39M, DEG-
0.54M, and water-2.5M.
86
At C02 partial pressures below 1 atm the solubility of the acid
gas is highest in the aqueous solution. Near 1 atm we have a
crossover in the figure, but since the amine strength of the
aqueous solution is about 5 times the strength of the nonaqueous
solutions, the aqueous MEA solution will still exhibit the best
C02 pick up even at partial pressures above 1 atm. However, Fig.
7.5 suggests that the shape of the equilibrium curves does not
change markedly with amine concentration, as long as the partial
pressures are given as a function of C02 loading and we operate
below the area where the physical absorption starts to
predominate. This implies that the C02 pick up in the glycol-
amine solutions can be improved essentially by using solutions
with much higher amine concentrations.
The equilibrium curves in Fig. 4.4, 4.5, 4.6 and 4.7 show that
the C02 solubility in these mixed solvents varies favorably with
temperature. This investigation does not include solubility
measurements at temperatures encountered in stripper columns.
However, it can be seen, from the difference in solubility at 30
and 70°C, that in a typical C02 removal process one can attain
a C02 pick up well above 0.5 mol C02/mol amine, based on
equilibrium considerations. Restrictions due to slow reactions
at high loadings may however result in a somewhat lower maximum
attainable C02 pick up.
In Fig. 7.6, the C02 solubility in a TEG solution with MEA is
compared with the solubility in a TEG solution with DEA, showing
some deviation in solubility at medium and low partial pressures.
87
10-3=
" _l _
— -J I 1" 1
, _ 1- _
u _ 1 1 1
1
1 1 1
»/ /
e 4-1
o u
' _ ^ J ^ ° / '
0.01 --
0.001 I I I I I I I I I
DEG i
i i i i i i i i i i i r i i i i i
0.00 0.20 0.40 0.60 y [mol C02/mol amine]
0.80
Figure 7.4 Comparison of equilibrium curves at 40°C for three different solvents, all containing 5mol% MEA * water, data from Lee et al. (1976) Q DEG, present data from Fig. 4.9 A TEG, present data smoothed from Fig. 4.4
88
10̂ =
13.6mol% DEA
6 XI 10 " 0.1 -:
o
0.01 - =
0.001 0.00 0.20 0.40 0.60 0.80
y [mol C02/mol amine]
Figure 7.5 Comparison of equilibrium curves for the TEG/DEA system at different amine concentrations at 30°C A 5mol% DEA * 10mol% DEA D 13.6mol% DEA
89
10--h
(0
- 0.1
£
0.01 -: =
0.001
. - - _ - 1 -
_ - J .
1 1
1
— — — — — — ^ .
~ ~1" ~ l '
*y
_ _ / _ J —/— -1
/ i
/DEA !>
7 / "• / / '
^a
/a
'MEA
t 1
1 1_ _ 1 1 1 ^
— C^AT^
i — y ryr~
. i i
i i i
i —
- r~ -i
_i _ _i _* -
y^\y ^ ^ ~ " K J T ^
— %*-=1 - — = :
i
i j
, _i _ _
i i i
~i ~ _ _j _ _
1
- * * " * " ^ l
_ I
0.00 i i r i i i i I i i i
0.20 i i i I i i i i i
0.40 i i i~i | r
0.60 r i i i i i i i I
0.80 y [mol C02/mol amine]
Figure 7.6 Comparison of equilibrium curves at 50°C for TEG solutions containing 10mol% MEA and 10mol% DEA
90
7.3 COMPARISON WITH OTHER MIXED SOLVENTS
C02 solubilities in solutions of TEG and n-methyl-pyrrolidone
(NMP) containing comparable concentrations of MEA, are compared
in Fig. 7.7. Data for the NMP/MEA system are collected from
Murrieta-Guevara and Trejo Rodriguez (1984) and Dimov et al.
(1976). The equilibrium curves indicate a higher solubility in
the NMP/MEA system. NMP without amine is used in the Purisol
process licensed by Lurgi and described by Grunewald (1989). At
absorption conditions the solubility of H2S in NMP is about ten
times higher than the solubility of C02. This makes it an
attractive solvent for removing H2S selectively (Kohl and
Riesenfeld (1985)).
0.1 J= = = = = = b = = = j : = fc=.é = =
0.01 - :
0.001 - : =
0.00
1 1
u 1 1
1
III 1 III 1
JU
U
nn
IM
i u
n
ni i
un
n
i i
: : : : : : # :
IHiplff 7 \_f~
1 mt-TEG7 /NMP
I I I IM I 7 I I I ' > ' '
^ - z zzycz, r̂ /" ,
l i sS i ^ - o i
_ _ Z r Z -ZZ r^Z Z _ i • / f
^ : : ^ : " : : : c : ~ : :
- / - * • '
/?..:„.. .: . . . . i i i i
11 M
t 1
Ull
nm
i
n rr
—
i t
IMI
I U
ll n
m
II n
i ll
lll
•in
n-
—i
i n
u
Ulli
inn
1
Ull
'II 1 M | 1 1 II II 1 1 1 | 1 1 1 1 1 1 1 1
i i
i i
i i
ni i
ii11
UU
L
lill
1 1
1 1
1 1
1 1
~ r 1 i i
1 ] ! 1 1 1 1 1 1 l-T |
0.20 0.40 0.60 0.80 y [mol C02/raol amine]
1.00
Figure 7.7 Present C02 solubility data in a mixed TEG/MEA solution ( * ) compared with the solubility in NMP/MEA solutions at 50°C. Data are taken from Dimov et al. (1976) ( ffl ) and Murrieta-Guevara et al. (1984) ( • )
91
7.4 COMPARISON WITH PURE PHYSICAL SOLVENTS
A comparison between the pure physical solvent (TEG) and the same
physical solvent with an amine (MEA) added, is here given. In
Fig. 7.8 we look at the C02 solubility in the TEG/MEA system at
50°C, compared with the physical solubility in a pure TEG
solution (data from Jou et al. (1987)). The figure shows a strong
increase in CO2 solubility with increasing amine concentration.
1.50
1.00 -
0.50 -
0.00 r 1 1 1 0.00
*TT 1 i i 1 1 1 1 1 1 r 1 1 1 r 1 ri 1 1 i~| 1 1 1 1 1 1 1 1 1 0.02 0.04 0.06 0.08
x [mol C02/mol tot]
Figure 7.8 C02 solubility data for 5mol% and 10mol% MEA in TEG, compared with the solubility in pure TEG (Jou et al. (1987))
92
Chapter Eight
Conclusions and Recommendations
The results of this work are discussed within the different
chapters of this thesis. The most important findings are
summarized in this chapter, and recommendations are given for
future work.
8.1 CONCLUSIONS
The development of a simple and reliable modelling technique to
describe the vapor-liquid equilibria of C02 in aqueous
alkanolamine solutions, is regarded as the main contribution of
this work. By making use of measured pH data, we have
circumvented the problem of estimating interaction parameters,
activity coefficients, and equilibrium constants, in the
prediction of vapor-liquid equilibria. The applicability of the
model is best demonstrated on the tertiary amine system using
MDEA. For this system, the VLE is accurately represented for
temperatures in the range 25 to 140°C, for C02 loadings from
0.0U1 to 1 mol/mol, and for amine molarities usually encountered
in acid gas treating processes. The absorption of C02 into
solutions containing the sterically hindered amine AMP, is also
well described by the model.
The equilibrium solubility of C02 in mixed solvents containing
a glycol i.TEG, DEG) and an alkanolamine (MEA, DEA) has been
measured at temperatures encountered in absorption units. An
equilibrium model, following much the same lines as the model
93
described for the aqueous systems, has been developed for the
C02/TEG/MEA system. This model enables estimation of CO2 partial
pressures, covering loadings and temperatures for both absorption
and desorption conditions.
The rates of absorption of C02 have been measured and compared
for five physical solvents, and for the same solvents containing
5mol% MEA. The measurements indicate that aqueous and alcoholic
solutions of MEA absorb CO2 considerably faster than solutions
of NMP, TEG, or glycolalkylether.
An important spin-off of the work described in this thesis, is
that two new experimental set-ups have been designed and built.
These are an apparatus for equilibrium solubility measurements
at higher temperatures, and a one-sphere apparatus for
measurements of reaction kinetics. Both of these set-ups are now
in use.
8.2 RECOMMENDATIONS
In future works, it is recommended that H2S absorption into the
same aqueous systems (MDEA and AMP) should be investigated with
the objective of developing a similar VLE model. This could
enable approximate process calculations on both simultaneous and
selective removal of H2S and C02. Since the chemistry of the
reaction between H2S and alkanolamines is quite simple and well
understood, the development of such a model should be attainable,
given that a pH electrode is used that is not polluted by the H2S
present in the solution.
An extension of the VLE model for the aqueous AMP system to
include desorption conditions, is also desirable. To accomplish
94
this, additional measurements on the new solubility apparatus,
as well as pH measurements, should be undertaken.
More work should be done to unveil both reaction kinetics and
equilibrium solubility of C02 and H2S in nonaqueous systems
containing amines. These are systems for which literature data
are scarce.
Measurements of C02 solubility in the TEG/MEA system at elevated
temperatures/ using the new solubility apparatus, should be
undertaken to validate the model developed in this work.
Consistent model equations, correlating the VLE in other
nonaqueous amine solutions, should be developed for a number of
systems. To do this, investigations giving rise to a better
understanding of the chemistry encountered in such systems, are
recommendable.
95
Nomenclature
AMP b,c,d,e
BF CF CR CHA DEA DEG DEGMME DIPA DGA ETG EtOH LNG MDEA MEA NMP NRTL PC PE R RNH2
R2NCH3
TEA TEG TBE VLE
a C
6P
f G H K
Kl
Kc Kp
Kp' L
[mol/l] [mol/1 = M] [atm]
[1/h] [atm'1/mol] tatm-l2/mol [mol/1]
[1/mcl] [mol/1] [mol/1] [1/h]
2-amino-2-methyl-1-propanol experimentally determined constants bicarbonate formation carbamate formation carbamate reversion cyclohexylamine diethanolamine diethyleneglycol diethyleneglycol monomethylether diisopropanolamine diglycolamine ethyleneglycol ethanol liquefied natural gas methyldiethanolamine monoethanolamine n-methyl-pyrrolidone nonrandom-two liquid propylene carbonate 2-piperidine ethanol alcoholic alkyl group primary amine, for example AMP tertiary amine, for example MDEA triethanolamine triethyleneglycol 2-(tertbutylamino) ethanol vapor-liquid equilibrium
activity concentration difference in water vapor partial pressure between the gas leaving the buffer solution and the condenser activity coefficient gas flow rate Henry's law coefficient (fHC03--H)/(fC02.K1) first dissociation constant of carbonic acid
carbamate stability constant amine protonation constant
(fRNH3+/fRNH2)#Kp liquid flow rate
96
L' m N P P ppmv
r t T X
y
[cn3/cti\'s] [mol/1] [kmol/m2«s] [atm] [atm] [ppm]
[mol/l's] [°C] [K] [mol C02/mol [mol C02/mol
total] amine]
liquid wetting rate amine molarity absorption rate total pressure partial pressure instrument reading from gas-analyzer converted from mA to volumetric ppm reaction rate temperature temperature CO2 molefraction in liquid phase C02 loading in liquid phase
97
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107
Appendix A
Tabulated Data of C02 Solubility in Aqueous Systems
Table 1 Solubility of C02 in aqueous solutions of 4.00M MDEA at 30, 45, and 60°c
y [raol C02/mol amine] pco2 [atm] T [°C]
0.066 0.129 0.136 0.213 0.250 0.256 0.331 0.387 0.444 0.499
0.055 0.081 0.130 0.153 0.185 0.226 0.298 0.338
0.044 0.051 0.066 0.094 0.106 0.158 0.176
0.0059 0.0170 0.0216 0.0390 0.0481 0.0580 0.0915 0.130 0.184 0.276
0.0118 0.0290 0.0467 0.0749 0.0898 0.1500 0.195 0.283
0.0210 0.0290 0.0607 0.0925 0.105 0.196 0.250
30 30 30 30 30 30 30 30 30 30
45 45 45 45 45 45 45 45
60 60 60 60 60 60 60
108
Table 2 Solubility of C02 in aqueous solutions of 4.28M MDEA at 25, 40, and 70°C
y [mol CO^/mol amine] pC02 [atm] T [°C]
0.096 0.143 0.334 0.430 0.523
0.061 0.132 0.186 0.261 0.314 0.383
0.061 0.107
0.0067 0.0155 0.0603 0.110 0.161
0.0090 0.0390 0.0640 0.121 0.170 0.256
0.080 0.220
25 25 25 25 25
40 40 40 40 40 40
70 70
Table 3 Solubility of C02 in aqueous solutions of 3.00M AMP at 40 and 50°C
y [mol C02/mol amine] Pc02 Catml T t°C]
0.670
0.596 0.688 0.746 0.764 0.793
0.0770
0.0786 0.136 0.236 0.299 0.316
40
50 50 50 50 50
109
Appendix B
Tabula ted Data of C02 S o l u b i l i t y i n Nonaqueous Systems
Table 4 S o l u b i l i t y of C02 i n s o l u t i o n s o f TEG and 5mol% MEA (0.39M) a t 30 , 50 , and 70°C
T [°C] y [mol C02 /mol amine] p C 0 2 [atm]
30 0 .418 0 .040 30 0 .462 0 .113 30 0 .495 0 .169 30 0 .500 0.181 30 0 .790 1 .000 ( e x t r a p o l a t e d )
50 0 .263 0 .030 50 0 .289 0 .047 50 0 .289 0 .087 50 0 .295 0 .065 50 0 .308 0 .171 50 0 .340 0 .134 50 0 .341 0 .147 50 0 .359 0 .171 50 0 .367 0 .196 50 0 .404 0 .231 50 0 .462 0 .286 50 0 .673 0 .995
70 0 .058 0 .029 70 0 .096 0 .054 70 0 .109 0.075 70 0 .122 0 .106 70 0 .160 0 .176 70 0 .212 0 .223 70 0 .224 0 .275 70 0 .237 0 .335 70 0 .436 1.011
110
Table 5 Solubility of C02 in solutions of TEG and 10mol% MEA (0.79M) at 30, 50, and 70°C
T [°C] y [mol C02/mol amine] p c o 2 [atm]
30 0.308 0.012 30 0.392 0.034 30 0.475 0.080 30 0.519 0.180 30 0.532 0.206 30 0.392 0.022 30 0.446 0.039 30 0.443 0.061 30 0.472 0.084 30 0.472 0.115 30 0.484 0.158 30 0.516 0.198 30 0.503 0.248 30 0.563 0.294 30 0.627 1.010
50 0.215 0.022 50 0.304 0.058 50 0.418 0.220 50 0.428 0.230 50 0.222 0.021 50 0.298 0.052 50 0.345 0.075 50 0.392 0.117 50 0.402 0.156 50 0.449 0.199 50 0.446 0.249 50 0.440 0.291 50 0.462 0.344 50 0.535 1.011
70 0.048 0.012 70 0.089 0.028 70 0.174 0.072 70 0.247 0.123 70 0.250 0.162 70 0.275 0.207 70 0.278 0.265 70 0.323 0.318 70 0.329 0.405 70 0.424 1.011
111
T a b l e 6 S o l u b i l i t y o f C02 i n s o l u t i o n s o f TEG and 5mol% DEA (0.38M) a t 3 0 , 50 , and 70°C
T [°C] y [mol CO2/1110I amine] pC02 ta t ro]
30 0 .139 0.023 30 0 .126 0.019 30 0 .212 0.056 30 0 .284 0.090 30 0 .317 0.139 30 0 .344 0.177 30 0 .390 0.225 30 0 .443 0 .315 30 0 .602 0.992
50 0.033 0.021 50 0.099 0.036 50 0 .086 0 .058 50 0 .132 0.103 50 0.192 0.141 50 0 .152 0.165 50 0.191 0.221 50 0.225 0.289 50 0.258 0.319 50 0 .377 0 .994
70 0.026 0 .088 70 0.072 0 .146 70 0.138 0.155 70 0.151 0.220 70 0.290 0.994
Table 7 Solubility of C02 in solutions of TEG and 10mol% DEA (0.77M) at 30 and 50°C
T [°C] y [mol C02/mol amine] pco2 [atm]
30 0.293 0.056 30 0.368 0.105 30 0.394 0.141 30 0.498 0.169 30 0.521 0.266 30 0.512 0.327 30 0.512 0.337 30 0.564 0.996
50 0.121 0.051 50 0.134 0.067 50 0.140 0.090 50 0.192 0.128 50 0.231 0.177 50 0.218 0.195 50 0.277 0.276 50 0.283 0.373 50 0.453 0.996
113
Table 8 S o l u b i l i t y o f C02 i n s o l u t i o n s of TEG and 13.6mol% DEA (1.06M) a t 3 0 , 5 0 , and 70°C
T [°C] y [mol C02 /mol aminel p C 0 2 [atm]
30 0 . 2 4 3 0 .032 30 0 . 2 8 8 0 .046 30 0 . 2 7 6 0 .058 30 0 . 3 1 6 0 .080 30 0 . 3 7 3 0 .112 30 0 . 3 8 7 0 .187 30 0 . 4 1 0 0.251 30 0 . 4 3 2 0 .307 30 0 . 4 4 3 0 .342 30 0 . 4 4 8 0.371 30 0 . 5 2 6 0 .989
50 0 . 1 0 4 0 .030 50 0 . 1 4 6 0 .048 50 0 . 1 5 3 0.074 50 0 . 1 8 2 0.111 50 0 . 2 5 7 0 .162 50 0 . 2 5 2 0 .210 50 0 .281 0.285 50 0 . 3 1 8 0.343 50 0 . 3 5 6 0 .368 50 0 .429 0 .992
70 0 .017 0.O32 70 0 .045 0 .067 70 0 .057 0 .099 70 0 .054 0 .126 70 0 .087 0 .162 70 0 .109 0.227 70 0 .125 0.267 70 0 .132 0 .290 70 0 .168 0 .408 70 0 .252 0 .992
Table 9 Solubility of C02 in solutions of DEG and 5mol% MEA (0.54M) at 40°C
T [°C] y [mol C02/mol amine] pco2 [atm]
40 0.461 0.044 40 0.405 0.036 40 0.395 0.069 40 0.405 0.092 40 0.447 0.055 40 0.470 0.120 40 0.461 0.147 40 0.479 0.091 40 0.479 0.194 40 0.516 0.130 40 0.526 0.186 40 0.577 0.341 40 0.554 0.251 40 0.684 0.980 40 0.702 1.007
Table 10 Solubility of C02 in solutions of DEG and 10mol% MEA (1.10M) at 40°C
T [°C] y [mol C02/mol amine] pco2 [atm]
40 0.367 0.027 40 0.418 0.045 40 0.356 0.060 40 0.377 0.081 40 0.409 0.112 40 0.409 0.149 40 0.441 0.188 40 0.477 0.235 40 0.472 0.318 40 0.527 0.993
115
Appendix C
Tabula ted pH Data fo r Aqueous Systems
T a b l e 11 pH v a l u e s a s a f u n c t i o n of C02 l o a d i n g i n aqueous s o l u t i o n s of 4.00M MDEA a t 30 , 40 , 50 , and 60°C
y tmol C0 2 /mol amine] ( 1 - y ) / y pH T [°C]
0 .015 0 .063 0.229 0 .486 0 .668
0.190 0 .449 0 .599
0.030 0.114 0 .239 0 .468
0.0313 0.0850 0.253 0.346
6 5 . 7 1 4 . 9 3 .37 1 .057 0 . 4 9 8
4 . 2 6 1 .227 0 .669
3 2 . 3 7 .77 3 .18 1 .137
31 .0 1 0 . 7 6 2 .95 1.890
10 .316 9 .700 9 .187 8 .713 8 .484
9 .165 8 .600 8 .405
9 .738 9 .174 8 .842 8 .483
9 .547 9 .136 8 .600 8 .429
30 30 30 30 30
40 40 40
50 50 50 50
60 60 60 60
116
Table 12 pH values as a function of C02 loading in aqueous solutions of 3.00M AMP at 20, 30, 40, and 50°C
y [mol C02/mol amine] (1-y)/y pH T [°C]
0.512 0.631 0.780 0.882
0 . 4 6 9 0 . 6 4 6 0 .791 0 .900
0.493 0.645 0.717 0.840
0.499 0.616 0.700 0.822 0.861
0.953 0.780 0.282 0.134
1 .132 0.548 0.264 0.111
1 .030 0.550 0.395 0.190
1 .004 0.623 0.429 0.217 0.161
9.809 9.516 9.041 8.575
9.651 9.243 8.763 8.315
9.338 8.988 8.784 8.400
9.099 8.880 8.660 8.320 8.180
20 20 20 20
30 30 30 30
40 40 40 40
50 50 50 50 50
117
Appendix D
Tabulated Results of Kinetic Measurements
Table 13 Rate of absorption of C02 in water at 20°C
L ,L* G k, k'i N [1/h] [cm3/cm«s] [1/h] [cm/s] [cm/s] [kmol/m2-s]
(estimated 106 from equation given by Morris & Jackson (1953))
4 . 1 3 4 . 1 4 4 . 2 5 4 .47 4 . 8 5 4 . 9 4 5 . 0 7 5 .47 6 .11 6 .12 6 .31 6 .58 6 .79 7 .14 7 . 6 8 7 .70 7 .75 8 .65 9 .17 9 . 4 6
10 .50
0 .312 0 .313 0 .317 0 .337 0 .366 0 .368 0 .383 0 .413 0 .461 0 .456 0 .476 0 .497 0-513 0 .539 0 .580 0 .573 0 .585 0 .653 0 .692 0 .714 0 . 7 8 2
2 .69 2 .76 2 .81 2 . 8 2 2 .98 3 .15 3 .16 3 .28 3 .32 3 .63 3 .64 3 .95 3 .82 3 .83 4 .28 4 .65 4 .34 4 .86 5 .03 5 .44 6 .59
0 .013 0 .015 0 .014 0 .013 0 .013 0 .015 0 .014 0 .013 0 .012 0 .014 0 .014 0 .016 0 .014 0 .013 0 .016 0 .019 0 .016 0 .018 0 .018 0 .021 0.029
0 .0088 0 .0089 0 .0089 0 .0093 0 .010 0 .010 0 .010 0 .011 0 .012 0 .012 0 .012 0 .012 0 .013 0 .013 0 .014 0 .014 0 .014 0 .015 0 .015 0 .016 Q.Q17
1 .43 1.47 1.44 1.50 1.59 1.62 1.68 1.75 1.77 1.86 1.94 2 .10 2 .03 2 .04 2 .28 2 .39 2 .31 2 . 5 9 2 .68 2 . 8 9 3 . 3 8 ( R i p p l e s )
Table 14 Rate of absorption of C02 in a solution of water and 5mol% MEA at 20 °C. Temperature at the end of the experiment is also given
L [ 1 / h ]
1 . 7 0 4 . 6 3 5 . 7 1 7 . 3 3
L' [cm3/cm-s]
0 . 1 2 7 0 . 3 4 5 0 . 4 2 5 0 . 5 4 6
G [ 1 / h ]
50 65 71 77
. 0
. 0
. 3
. 8
N [kmol/m2 's]
1 0 5
25 3 3 , 36 39 ,
. 7
. 4
. 6
. 9
f
[e
3 5 3 3 31 31
r 'C]
. 0
. 3
. 3
. 0
118
Table 15 Rate of absorption of C02 in n-methyl-pyrrolidone at 20°c
L Il/h]
1 .77 2.52 3.72 4.72 6.37 7.36
-L' [cnr/cm-s]
0,134 0.190 0,280 0,356 0.481 0.554
G [1/h]
4.74 5.30 6.89 7.76 9.84 11 .02
N [kmol/m2,s]
105
2.52 2.82 3.67 4.13 5.26 5.86
Table 16 Rate of absorption of CO2 in a solution of n-methyl-pyrrolidone and 5mol% MEA at 20°C. Temperature at the end of the experiment is also given
L [1/h]
2.33 3.29 3.89 4.01 4.90 5.90 5.96
,L' [cm3/cm's]
0.176 0.248 0.293 0.302 0.370 0.446 0.450
G [1/h]
15.79 18.44 19.51 19.65 22.26 24.45 25.62
N [kmol/m2*s]
106
8.40 9.81
10.38 10.46 11 .85 13.01 13.63
T [3C]
29.0 28.3 27.1 28.3 27.1 27.1 27.2
119
Table 17 Rate of absorption of C02 in ethanol at 20°C
I. [1/h] [cm3/cn»'s]
G tl/h]
N [kraol/m2«s]
106
2.34 2.55 2.9T 3.15 3.58 3.85 4.22 6.07 6.17 6.41 7.16
0.177 0.190 0.220 0.238 0.267 0.291 0.314 0.453 0.466 0.483 0.534
5.36 6.23 5.86 6.17 8.62 8.81 9.76
14.63 13.21 13.95 16.57
2.85 3.20 3.12 3.28 4.42 4.69 5.01 7.51 7.03 7.42 8.50
Table 18 Rate of absorption of CO2 in a solution of ethanol and 5mol% MEA at 20°C. Temperature at the end of the experiment is also given
t. [ 1 / h ]
2 . 6 4 3 . 5 4 4 . 1 9 4 . 7 2 6 . 0 2 6 . 0 4 6 . 7 5 7 . 3 6
L' [ c m 3 / c m - s ]
0 . 1 9 7 0 . 2 6 3 0 . 3 1 2 0 . 3 5 1 0 - 4 4 8 0 . 4 5 0 0 . 5 0 4 0 . 5 4 7
G [ 1 / h ]
3 0 . 2 2 4 1 . 1 4 4 3 . 3 7 5 2 . 1 7 6 3 . 4 4 6 5 . 1 6 6 9 . 5 7 7 7 . 8 4
N [ k m o l / m 2 * s ]
1 0 6
1 5 . 5 1 2 1 . 1 1 2 2 . 2 6 2 6 . 7 8 3 2 . 5 6 3 3 . 4 4 3 5 . 7 1 3 9 . 9 5
T l ° C ]
3 2 . 8 3 2 . 8 3 2 . 8 3 2 . 6 3 3 . 6 3 3 . 0 3 3 . 8 3 3 . 8
120
Table 19 Rate of absorp t ion of C02 i n t r i e t h y l e n e g l y c o l a t 20°C
L [1 /h ] [cnr/cm-s]
G [1 /h]
N [kmol /nr ! s ]
106
0,95 3.60 5.93
0.071 0.268 0.442
0.321 0.604 0.759
0.165 0.310 0.390
Table 20 Rate of absorption of C02 in triethyleneglycol and 5mol% MEA
solution of
Ii [ 1 / h ]
1 .73 3 . 3 6 4 . 1 7 6.11
, L ' [cm-Vcm^s]
0 .129 0 .250 0.311 0 .455
G [ 1 / h ]
1 .96 2 . 6 0 2 . 7 3 3 . 4 4
N [kmol/m2»s]
10 6
1 .00 1 .33 1 .40 1 .77
T [°C]
2 2 . 8 2 2 . 8 2 2 . 8 2 2 . 8
121
Table 21 Rate of absorption of C02 in diethyleneglycol monomethylether at 20°C
L L' G N „ [1/h] [cm3/cm«s] [1/h] [kmol/m2«s]
106
1.19 3.11 3.19 4.30
0.089 0.232 0.238 0.320
1 .21 2.66 3.07 3.46
0.62 1 .36 1 .58 1.77
Table 22 Rate of absorption of C02 in a solution of diethyleneglycol monomethylether and 5mol% MEA at 20 °C. Temperature at the end of the experiment is also given
L Il/h]
1 .57 2.10 2.23 3.49 4.49
L' [cm3/cm*s]
0.117 0.156 0.166 0.260 0.334
G [1/h]
4.11 5.32 5.82 6.86 7.76
N [kmol/m2«s]
106
2.11 2.73 2.99 3.52 3.98
T [°C]
25.5 26.4 26.3 26.3 26.4
122
Appendix E
HP-42S Program for C a l c u l a t i o n of Equi l ib r ium P a r t i a l P r e s s u r e of C02 over Aqueous MDEA
The program r e a d s v a l u e s of C02 l o a d i n g , t e m p e r a t u r e , and amine c o n c e n t r a t i o n , and a s e m i e m p i r i c a l model c a l c u l a t e s t h e e q u i l i b r i u m p a r t i a l p r e s s u r e of C02 o v e r aqueous s o l u t i o n s of MDEA.
01 LBL "MDEA" 20 +/-
02 "y=?" 21 1
03 PROMPT 22 +
04 STO 00 23 RCL 00
05 "m=?" 24 /
06 PROMPT 25 LOG
07 STO 01 26 -0.59
08 "T=?" 27 X
09 PROMPT 28 RCL 03
10 STO 02 29 +
11 0.0248 30 10x
12 X 31 RCL 00
13 8.6 32 X
14 - 33 RCL 01
15 STO 03 34 X
16 ". ." 35 "PC02 = "
17 ARCL ST X 36 ARCL ST X
18 AVIEW 37 AVIEW
19 RCL 00 38 END
123
