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
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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:
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
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
(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:
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
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
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
References
Al-Ghawas, H.A., PhD Dissertation, University of California, Santa Barbara, CA, 1988.
Al-Ghawas, H.A., Hagewiesche, D.P., Ruiz-Ibanez, G., Sandall, O.C., J. Chem. Eng. Data (1989), 34, pp. 385-391.
Pellegrino, J.J., Ko, M., Nassimbene, R., Einert, M., Proceedings of the International Symposium on Gas Separation Technology, Antwerp, Belgium, 1989, pp. 445-456.
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.