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In situ studies of the thermal evolution of the structure and sorption properties of Mg-Al-CO3 layered double hydroxide
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In situ studies of the thermal evolution of the structure and sorption properties of Mg-Al-CO3 layered double hydroxide
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Content
IN SITU STUDIES OF THE THERMAL EVOLUTION
OF THE STRUCTURE AND SORPTION PROPERTIES OF MG-AL-CO
3
LAYERED DOUBLE HYDROXIDE
by
Yongman Kim
A Dissertation Presented to the
FACULTY OF THE GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(CHEMICAL ENGINEERING)
December 2006
Copyright 2006 Yongman Kim
ii
Acknowledgements
My advisor Dr. Theodore T. Tsotsis and Dr. Muhammad Sahimi have
supported me throughout my graduate student career. I would particularly like to
acknowledge them for their emphasis on the troubleshooting for all kinds of
happenings. Also, I would like to thank Dr. Paul KT Liu at Media & Process
Technology, he had generously sent hydrotalcite samples along with lots of other
samples to analyze. It was a great challenge to keep up the analysis schedule on
time. I would like to acknowledge Dr. Weishen Yang, I had learnt many things
from him during his short stay at USC. And I would like to thank Dr. Edward Goo,
Dr. Katherine Shing, and Dr. Ted Lee for their helpful advices as a committee
member of my thesis.
I would like to acknowledge Karen Woo and Brendan Char for their
excellent administrative works and also for lots of friendly helps as well. Polly
Chan and Tina Silvar, whenever I became a TA of the undergraduate lab courses,
their helps with big smile have made all lab works joyful. John Curulli gave
generously of his time to get me started with XRD, SEM, and TEM. Also,
whenever I met problems with instrument, John worked so hard to fix it out.
My fellows, Babak Fayyaz, Mahnaz Firouzi, Mayur Ostwal, Nayong Kim,
Taewook Kim, Hyuntae Hwang, and Asdesh Harale, they have been great friends
and colleagues. During my studying years, I spent time with my fellows even more
than my wife, and I had many delight memories with them. I’m sure I’m going to
iii
miss them. I would particularly like to acknowledge Seong Yun Lim for helping me
through some difficult times, and his advices for my research projects.
My senior friends, Daniel Yu and Juhyeon Choi, in my first year of living in
USA, I totally depended on their helps. Also, whenever I had hard times, they
shared my burdens without hesitating. They have an unsated passion with great sprit
Far and away the biggest thanks go to my parents, Ki Jun Kim and Sang Jo
You. By valuing education so highly, acting as my financial safety net, and showing
endless trust to me, they have made me possible to stand five long years.
Lastly, I thank my wife Heejung Han for her support and patience. She
walked through all difficult times with unwavering support, and she has inspired my
emotions as my best friend.
iv
Table of Contents
Acknowledgements
ii
List of Tables
vi
List of Figures viii
Abstract xiii
1 Introduction
1.1 Background……………………….………………….…………………….. 1
1.2 Motivation and Overview of Study………………………………………... 8
Chapter 1 References…...…..……………………………………………… 16
2 In-situ Thermal Evolution Study of Mg-Al-CO3 LDH
2.1 Introduction ………………………………………………………………. 22
2.2 Experimental …………………………………………………………...… 23
2.3 Results and Discussion …………..………………………………………. 27
2.4 Conclusion……………..…………………………………………………. 50
Chapter 2 References……………………………………………………… 52
3 The Study of Sorption Properties of Mg-Al-CO3 LDH
3.1 Introduction ……………………………………………………………….. 53
3.2 Experimental ……………………………………………….………...…… 53
3.3 Results and Discussion……………………………………………….…… 54
3.4 Conclusion ………………………………………………………………… 92
Chapter 3 References ……………………………………………………… 94
4 Diffusivity and Adsorption Isotherm of Carbon Dioxide in Mg-Al-
CO3 LDH at Elevated Temperatures
4.1 Introduction ……………………………………………………………..… 95
4.2 Experimental ………………………………………………………………. 100
4.3 Theory and Results…..……………………………………………………. 103
4.4 Conclusion ………………………………………………………………… 137
Chapter 4 References……………………………………………………… 139
v
5 Preparation and Characterization of Mg-Al-CO
3
LDH membranes
5.1 Introduction ……………………………………………………………..… 142
5.2 Experimental ……………………………………………………………….147
5.3 Results and Discussion
5.3.1 The co-precipitation method………………………………………… 152
5.3.2 The sol-gel method…………………………………………………... 156
5.3.3 Membrane preparation by separate nucleation and aging steps……… 161
5.3.4 Using sulfuric acid as a binder……………………………………….. 162
5.4 Conclusion ………………………………………………………………… 164
Chapter 5 References………………………………………………………. 166
Alphabetized Bibliography 168
vi
List of Tables
Table 1.1 Ionic radius of some cations.
5
Table 1.2 Values of c’ for some inorganic anions.
7
Table 2.1 The changes in the basal spacing of Mg-Al-CO
3
LDH with
temperature calculated from the HTXRD patterns.
48
Table 2.2 The changes in the basal spacing of Mg-Al-CO
3
LDH with
temperature calculated from the HTXRD patterns.
48
Table 3.1 Weight-loss from the TG/MB-MS studies, and calculated
weight-loss based on the ICP data for the samples. (a) LDH1;
and (b) LDH2.
55
Table 3.2 The fractions of H
2
O and CO
2
(as % of the total sample
weight) that are evolved in different temperature ranges for
both LDH1 and LDH2.
60
Table 3.3 Weight-gain during the sorption step for the moderate-pressure
flow experiments using dry CO
2
.
89
Table 3.4 Weight-gain during the sorption step for the moderate pressure
flow experiments at various temperatures using dry CO
2
.
90
Table 3.5 Weight-gain during the sorption step for the moderate pressure
flow experiments at various temperatures using humidified
CO
2
.
91
Table 4.1 Diffusivity data (D/r
2
) for CO
2
measured by gravimetric
method.
96
Table 4.2 Diffusion coefficients measured by experiment and calculated
by molecular dynamic simulation.
109
Table 4.3 Langmuir adsorption parameters of CO
2
in LDH2. 115
vii
Table 4.4 Sips equation parameters for CO
2
in LDH2.
116
Table 4.5 The Toth equation parameters for CO
2
in LDH2.
118
Table 4.6 The chi
2
values of nonlinear fittings for different models 120
Table 4.7 The characteristic energies for CO
2
in LDH2 with DR
equation.
121
Table 4.8 The uptake amount and BET surface area of LDH3 particles. 127
Table 4.9 The Langmuir adsorption parameters of CO
2
in LDH3 at
200
o
C with different particle size.
128
Table 4.10 The Sips isotherm equation parameters of CO
2
in LDH3 at
200
o
C with different particle size.
129
Table 4.11 The raw data of CO
2
uptake of sieved LDH3 132
Table 4.12 Polynomial form of particle size distribution function, f(R)=
Ax
6
+ Bx
5
+ Cx
4
+ Dx
3
+ Ex
2
+ Fx+ G.
136
viii
List of Figures
Figure 1.1 (a) Brucite [Mg(OH)
2
] single sheet (b) Hydrotalcite
[Mg
6
Al
2
(OH)
16
][(CO
3
)
.
4H
2
O], structure, Top view without
showing H
2
O molecules.
3
Figure 2.1 Comparison between the DRIFTS and FTIR results of Mg-
Al-CO
3
LDH.
26
Figure 2.2 In-situ DRIFTS of Mg-Al-CO
3
LDH as a function of
temperature.
29
Figure 2.3
In-situ DRIFTS of MgCO
3
as a function of temperature.
32
Figure 2.4 Fraction of species removed from Mg-Al-CO
3
LDH as a
function of temperature.
33
Figure 2.5 In-situ MS of Mg-Al-CO
3
LDH as a function of temperature.
35
Figure 2.6 In-situ TG/MS of Mg-Al-CO
3
LDH as a function of
temperature.
36
Figure 2.7 In-situ TG/DTA of Mg-Al-CO
3
LDH as a function of
temperature.
39
Figure 2.8
In-situ TG/DTA of Al(OH)
3
as a function of temperature.
40
Figure 2.9 In-situ TG/DTA of Mg(OH)
2
as a function of temperature.
41
Figure 2.10 The thermal evolution of Mg-Al-CO
3
LDH as a function of
temperature.
45
Figure 2.11 In-situ HTXRD of Mg-Al-CO
3
LDH as a function of
temperature.
46
ix
Figure 3.1 The XRD spectra of the LDH samples. (a) LDH1 sample; (b)
LDH2 sample.
56
Figure 3.2 (a) A SEM picture of a membrane tube prepared by the
deposition of a thin LDH layer on an underlying macroporous
support; (b) TEM picture of LDH1; (c) TEM picture of
LDH2.
57
Figure 3.3 The TGA spectra and CO
2
MS signal for the two LDH
samples generated with a scan rate of 5
o
C/min; (b)
cumulative amount of H
2
O evolved.
58
Figure 3.4 The effect of varying (a) the heating rate; and (b) of using
different purging gases on the weight-loss for the LDH2
sample.
62
Figure 3.5 Weight-gain or loss. (a) weight-gain during adsorption for
various temperatures as a function of the cycle number; (b)
weight-loss during desorption for various temperatures as a
function of the cycle number; (c) weight-change due to loss
of water or CO
2
during desorption as a function of
temperature.
67
Figure 3.6 (a) In-situ DRIFTS of LDH2 as a function of temperature; (b)
change in the CO
3
2-
integrated peak area (left), and change in
the OH
-
and H
2
O integrated peak areas as a per cent fraction
of the original peak area (right) during the sorption-
desorption cycles.
71
Figure 3.7 Weight-gain or loss (top) and total sample weight (bottom)
during the sorption/desorption cycles.
74
Figure 3.8 H
2
O and CO
2
MS signals during the heating, and the
desorption part of the cycles as a function of time.
75
Figure 3.9 Weight-gain or loss (top) and total sample weight (bottom)
during the sorption/desorption cycles.
76
Figure 3.10 H
2
O and CO
2
MS signals during the heating, and the
desorption parts of the cycle as a function of time.
77
x
Figure 3.11 Weight-gain or loss (top) and total sample weight (bottom)
during the sorption/desorption cycles.
79
Figure 3.12 Weight-loss/gain during the temperature cycling experiments.
Solid lines are the experiments from room temperature to
150
o
C; Dotted lines are experiments from room temperature
to 200
o
C.
80
Figure 3.13 MS signals for H
2
O (top) and CO
2
(bottom) during the
temperature cycling experiment from room temperature to
150
o
C.
83
Figure 3.14 MS signals for H
2
O (top) and CO
2
(bottom) during the
temperature cycling experiment from room temperature to
200
o
C.
84
Figure 3.15 Weight-loss/gain during the temperature cycling experiments.
(a) from room temperature to 250
o
C; (b) from room
temperature to 300
o
C; (c) from room temperature to 350
o
C.
86
Figure 3.16 MS signals for H
2
O (top) and CO
2
(bottom) during the
temperature cycling experiment from room temperature to
250
o
C.
87
Figure 4.1 A graph of (a) M
t
/M
∞
against t
1/2
, and (b) ln(1- M
t
/M
∞
)
against t (bottom) for the uptake of carbon dioxide at 200
o
C.
104
Figure 4.2 A graph of (a) M
t
/M
∞
against t
1/2
, and (b) ln(1- M
t
/M
∞
)
against t (bottom) for the uptake of carbon dioxide at 225
o
C.
105
Figure 4.3 A graph of (a) M
t
/M
∞
against t
1/2
, and (b) ln(1- M
t
/M
∞
)
against t (bottom) for the uptake of carbon dioxide at 250
o
C.
106
Figure 4.4 Temperature dependence of diffusion coefficient for CO
2
in
LDH.
108
xi
Figure 4.5 Schematic representation of the gradient-echo pulse
sequence. (a) Radiofrequency pulse sequence. (b) Associated
magnetic field gradient pulses (shaded areas). The time
between the start of the π/2 pulse and the center of the
primary echo, during which T
2
relaxation takes place, is
denoted by 2τ; Δ is the time between the leading edges of the
field gradient pulses, g is the magnitude of the magnetic field
gradient pulses (mTm
-1
), and δ is their duration.
111
Figure 4.6 The PFG NMR spectrum of the LDH, which the spectrum of
a blank NMR tube was subtracted.
112
Figure 4.7 The experimental data and nonlinear curve fitting with the
Langmuir equation for adsorption isotherm of CO
2
in LDH2.
114
Figure 4.8 The experimental data and nonlinear curve fitting with the
Sips (Langmuir-Freundlich) equation for adsorption isotherm
of CO
2
in LDH2.
117
Figure 4.9 The experimental data and nonlinear curve fitting with the
Toth equation for adsorption isotherm of CO
2
in LDH2.
119
Figure 4.10 The experimental data and linear fitting with linearized DR
equation for adsorption isotherm of CO
2
in LDH2.
122
Figure 4.11 SEM images of LDH3 particles for fractionated sections. The
range of particle diameter: (a) 43~63 μm (b) 63~75 μm (c)
75~90 μm (d) 90~125 μm (e) 125~180 μm (f) 180~215 μm.
124
Figure 4.12 Particle size distribution of each sieved section of LDH3. The
range of particle diameter: (a) 43~63 μm (b) 63~75 μm (c)
75~90 μm (d) 90~125 μm (e) 125~180 μm (f) 180~215 μm.
125
Figure 4.13 The uptake amount of CO
2
with different LDH3 particle sizes
at 200
o
C
126
Figure 4.14 (a) The experimental data and nonlinear curve fitting with the
Langmuir equation, and (b) the experimental data and
linearized Langmuir equation for adsorption isotherm of CO
2
in LDH3.
130
xii
Figure 4.15 (a) The experimental data and nonlinear curve fitting with the
Langmuir-Freundlich equation, and (b) the parameter values
of Langmuir-Freundlich equation for adsorption isotherm of
CO
2
in LDH3.
131
Figure 4.16 Nonlinear relation between diffusion constants and particle
radii of LDH3.
134
Figure 5.1 The XRD spectrum of LDH powder synthesized by sol-gel.
144
Figure 5.2 The mechanism of the binding reaction between LDH and
sulfate groups.
147
Figure 5.3 (a) FT-IR spectrum and (b) XRD spectrum of LDH powder
synthesized by coprecipitation method.
149
Figure 5.4 The effect of aging time on (a) the flow rate of CO
2
via LDH
film on tubular α-alumina support (b) the separation factor
for N
2
/CO
2
with the LDH membrane prepared by
coprecipitation method.
153
Figure 5.5 The effect of hydrothermal aging temperature on the
separation factor for N
2
/CO
2
with the LDH membrane
prepared by coprecipitation method.
154
Figure 5.6 SEM image of (a) cracked surface, and (b) cross section of
LDH coated α-alumina support; (c) not covered surface of
LDH coated α-alumina support by coprecipitation method.
155
Figure 5.7 Separation factor for CO
2
/N
2
with LDH membrane prepared
by sol-gel method at room temperature.
158
Figure 5.8 The effect of (a) aging time, and (b) aging temperature on the
separation factor for CO
2
/N
2
with the LDH membrane
prepared by sol-gel method.
159
Figure 5.9 Separation factor for CO
2
/N
2
with 3
rd
layer on tubular α-
alumina support at room temperature and at 220
o
C.
160
Figure 5.10 The effect of (a) aging time, and (b) aging temperature on the
separation factor for CO
2
/N
2
with the LDH membrane
prepared by the separation of nucleation and aging method.
163
xiii
Abstract
Several in situ techniques have been used to investigate the thermal
evolution of the structure of Mg-Al-CO
3
layered double hydroxide under inert
atmosphere. Based on the results of the study, a model was proposed to describe the
structural evolution of the Mg-Al-CO
3
LDH. The sorption characteristics and
thermal reversibility of Mg-Al-CO
3
LDH were also investigated with in situ
techniques under both inert and reactive atmospheres. The experimental
observations are shown to be consistent with the structural model proposed. The
structure, sorption characteristics, and thermal reversibility of Mg-Al-CO
3
LDH
materials are important in their use for the high temperature applications.
Diffusivity constants and adsorption isotherms for carbon dioxide in Mg-
Al-CO
3
LDH were investigated by the gravimetric method at elevated temperatures.
The experimental estimates of diffusivity constants and those obtained by molecular
dynamic simulations are in good qualitative agreement. The experimental
adsorption isotherms for CO
2
in LDH have been studied with the Langmuir
isotherm equation and various empirical adsorption isotherm equations. It is
observed that the heterogeneity of the system and the interaction between CO
2
and
Mg-Al-CO
3
LDH increases with temperature. Particle size effect on CO
2
uptake
and adsorption isotherm was also investigated. It was observed that both the amount
of CO
2
uptake and BET surface area increases as the particle size decreases. When
the uptake amount was normalized with BET surface area, the uptake amount was
fairly constant for all the ranges of particle sizes.
xiv
Various methods for the preparation of LDH membranes were studies, and
the transport properties of prepared membranes were investigated. The values of
separation factor were similar to Knudsen factor, and, to prepare perm-selective
microporous LDH membranes, it is required to find an effective binder or a method
for LDH crystals to intergrow each other.
1
Chapter 1
Introduction
1.1 Background
Hydrotalcites, also known as layered double hydroxides (LDHs) or anionic
clays
1,2
, are hydroxycarbonate compounds of magnesium and aluminum, discovered
in Sweden around 1842. The exact chemical formula for hydrotalcite,
[Mg
6
Al
2
(OH)
16
][(CO
3
)
.
4H
2
O], was reported first by E. Manasse in 1915, who was
also the first to recognize that carbonate ions were essential for its structure
3
. In
1930 the existence of two different polytypes of hydrotalcite was reported by
Aminoff and Broomé based on their X-ray investigation, with one having
rhombohedral symmetry, and the other possessing hexagonal symmetry, and called
manasseite in honor of Manasse
4
. However, the layered structure of hydrotalcite
remained uncertain and confused at that time because of the lack of adequate
crystallographic data due to the complex and unusual composition of hydrotalcite.
In 1942 a large amount of a hydrotalcite-like compound was synthesized by
Feitknecht, who proposed that the synthesized compound consisted of one layer of
2
hydroxide with one cation, intercalated with a layer of a second one
5,6
. However,
Feitknecht’s hypothesis was refuted by Allmann and Taylor with their X-ray
analysis of monocrystals. They concluded that two cations exist on the same layer,
and only the carbonate anions and the waters are located in the interlayer
7,8
. The
earlier works of Allmann and Taylor utilized the minerals sjögrenite and pyroaurite
with the approximate composition [Mg
6
Fe
2
(OH)
16
][(CO
3
)
.
H
2
O], due to the
availability of monocrystals for such compounds. Later the structural investigation
of hydrotalcite was carried out by Allmann
9
, and the progress was reviewed by
Allmann
10
and by Taylor
11
. Actually, hydrotalcites found in nature are categorized
as the pyroaurite-sjögrenite group.
To understand the structure of hydrotalcite it is necessary to study the
structure of brucite, Mg(OH)
2
, where octahedra of Mg
2+
in a six-fold coordination
form with OH
-
groups share the edges to form infinite layers (see Fig. 1.1a), and
these layers are held together by hydrogen bonding. When Mg
2+
ions are substituted
by trivalent ions, the positive charges are created by heterovalent cations present in
the layer. The net positive charge is compensated by anions located in the interlayer
region between two positively charged metal hydroxide layers. For example, Mg
2+
(a) (b)
Figure 1.1 (a) Brucite [Mg(OH)
2
] single sheet (b) Hydrotalcite [Mg
6
Al
2
(OH)
16
]
[(CO
3
)
.
4H
2
O], structure, Top view without showing H
2
O molecules.
12
3
4
ions are substituted by Fe
3+
in the pyroaurite and by Al
3+
in the hydrotalcite
respectively. The net positive charge is compensated by CO
3
2-
anions (Fig. 1.1.b).
The compensating CO
3
2-
anions and water are randomly dispersed in the interlayer
like in the liquid water. They are free to move in the interlayer by breaking their
hydrogen bonds and forming new ones. The hydroxyl groups of the brucite-like
layer are tied to the CO
3
2-
anions directly or via intermediate water by hydrogen
bridges, and CO
3
2-
groups lying flat in the interlayer
10
.
However, interlayer water
molecules are loosely bound, and they can be removed without destroying the
structure of the material.
In general, it is possible to have LDH’s with the formula of [M(II)
1-x
M(III)
x
(OH)
2
]
x+
[A
n-
x/n
mH
2
O], indicating that it is possible to synthesize compounds with
different stoichiometries, with more than two metals or with two anions. M(II) and
M(III) ions can form hydrotalcite-like compounds as long as they can be
accommodated in the holes of the close-packed configuration of OH groups in the
brucite-like layers. For example, Be
2+
is too small, and Ca
2+
is too big for
octahedral coordination in brucite-like layers. However, it was reported by Allmann
that there are natural and synthetic LDHs with small amounts of Ca
2+
inside the
5
brucite-like layer
14
, which it was later confirmed by Drits et al
15
. All trivalent
cations with atomic radii between 0.5 and 0.8 Å (except V
3+
and Ti
3+
, due to their
instability in air) can form LDH’s (the ionic radii of some divalent and trivalent
cations are shown in Table 1.1).
Table 1.1 Ionic radius of some cations
13
, in Å.
Be
2+
Mg
2+
Cu
2+
Ni
2+
Co
2+
Zn
2+
Fe
2+
Mn
2+
Cd
2+
Ca
2+
M
2+
0.30 0.65 0.69 0.72 0.74 0.74 0.76 0.80 0.97 0.98
Al
3+
Ga
3+
Ni
3+
Co
3+
Fe
3+
Mn
3+
Cr
3+
V
3+
Ti
3+
In
3+
M
3+
0.50 0.62 0.62 0.63 0.64 0.66 0.69 0.74 0.76 0.81
The value of x in the above formula is typically found to be in the range of
0.2 and 0.33, but significantly higher values have also been reported
16
. Due to the
repulsion of positive charges, the Al
3+
ions in the brucite-like layer remain widely
separated. The Al octahedral can not be neighboring for x < 0.33
17
, and the
increased number of Al octahedra, for higher x values, result in the formation of
Al(OH)
3
. Similarly, low value of x lead to high densities of Mg octahedra in the
brucite-like layer which act as nuclei for the formation of Mg(OH)
2
. Therefore,
although LDH’s can exist for 0.1 ≤ x ≤ 0.5, many reports show that the value of x
might be in the range between 0.2 and 0.33 in order to obtain pure LDH’s
17-20
.
There is no limitation practically to the nature of anions located in the
interlayer and compensating for the positive charge of the brucite-like layer.
However, the thickness of the interlayer is determined by the number, size,
orientation and the strength of the bond between the anions and the hydroxyl groups
of the brucite-like layers. Though there is no limitation for interlayer anions, it is
very difficult to avoid contamination by the CO
2
in the aqueous solution when
hydrotalcite-like compounds are prepared with anions different from carbonate ions.
The following sequence of the affinity for the interlayer anions can be derived for
both mono- and divalent anions (the divalent anions are more strongly held in the
interlayer than the monovalent anions, and carbonates are the most strongly held
anions):
18, 21-24
CO
3
2−
SO
4
2 −
OH
−
> F
−
> Cl
−
> Br
−
> NO
3
−
> I
−
Carlino also reported that carboxylic acid did not readily undergo intercalation into
hydrotalcite-like compounds, since carbonate ions are tenaciously held in the
6
7
interlayer region
25
. The thickness of the interlayer region is the difference between
c` (calculated from the first basal reflection d
003
; hexagonal unit cell distance c =
3c`) and 4.8Å (the thickness of the brucite-layer)
21
. Values of c` for some inorganic
anions are presented in Table 1.2. The low value of c` for OH
-
is related to the
similarity of its ionic diameter with that of water molecule, and to the strong
hydrogen bond between water and hydroxyl group of the brucite-like layers; this
leads to the most close-packed arrangement
21,22
.
Table 1.2 Values of c` for some inorganic anions
22
, in Å.
Anion OH
-
(CO
3
)
2-
F
-
Cl
-
Br
-
I
-
(NO
3
)
-
(SO
4
)
2-
(ClO
4
)
-
C` 7.55 7.65 7.66 7.86 7.95 8.16 8.79 8.58 9.20
Recently, LDH’s have received much attention since these materials have a
well-defined layered structure with nanometer (0.3 ~ 3 nm) interlayer distances and
contain important functional groups. They have been used either fresh or mainly
after calcination, since the metal oxides obtained by calcination have high surface
area, are basic, stable to thermal treatment, and exhibit a memory effect. The
8
memory effect of LDH is the facile reconstruction of the layered structure after
exposure of the calcinated LDH to water
18, 38
. The first literature report referring to
the use of hydrotalcite-like compounds as basic catalysts appeared in 1971 by
Miyata et al.
26
. Earlier than 1971, there were a number of reports of catalysts,
whose precursors were similar to hydrotalcite-like compounds, but they were not
specially mentioned as relating to hydrotalcites. The memory effect of LDH’s is
mostly utilized in applications such as adsorption for liquid ions
27-29
and gas
molecules
30, 31
. Based on their other properties LDH’s also find use as catalysts for
oxidation
32-34
, reduction
35
, and other catalytic reactions
36-38
. LDH’s have also
recently found potential use in novel reactive separation applications for increasing
the conversion of catalytic reactions by removing one of the products from the
reactor
39-41
.
1.2 Motivation and Overview of Study
The “Green House” effect is one of the most important issues facing
researchers in the environmental and ecological areas. It is well known that the
main contributor to the green house effect is CO
2
. The 1995 Intergovernmental
9
Panel on Climate Change predicted that the atmospheric CO
2
would double by the
middle of next century unless additional measures are taken. Internationally, the
climate change problem has received more attention in recent years since its
potential impacts are significant. Therefore, efforts to capture and sequester CO
2
in
flue gas emissions have also increased. The field of CO
2
sequestration is still at an
early state with relatively little research being carried out for developing efficient
and economic technologies. For the long-term, however, it is important to develop
methods for capturing and utilizing CO
2
. In addition to CO
2
sequestration, CO
2
separation is also significant in natural gas purification, and in hydrogen production
by steam reforming of methane and other hydrocarbons.
The potential applications of LDH’s are in a broad range of temperatures,
ranging from room temperature adsorption to high temperature catalytic reactions
(~500
o
C). Also, some of the functional groups that LDH’s contain are known to
be sensitive to high temperatures. Therefore, it is important to understand the
thermal evolution of the structure of LDH’s in order to correlate better their
performance with their structure and functional groups. The common techniques to
study the structure and functional groups in LDH’s are X-ray diffraction (XRD),
10
Fourier transform infrared spectroscopy (FTIR), thermal gravimetric analysis
/differential thermal analysis (TG/DTA), scanning electron microscopy/transmission
electron microscopy (SEM/TEM), and gas chromatography (GC) or mass
spectrometer (MS). Since each technique has its own limitations, combining several
is often needed to obtain a more detailed picture of the structure of LDH’s and their
thermal evolution. One good example of such an investigation is the work of Hibino
et al.
42
, who used XRD, IR, TG/DTA, TEM, and GC to characterize the
decarbonation behavior of Mg-Al-CO
3
LDH’s. In their study they acquired the
XRD and IR spectra of samples treated at different temperatures, ex-situ (room
temperature). Hence, the results obtained this way may not really reflect the actual
LDH structure or the state of its functional groups, as they exist at any given
temperature. This is because upon cooling and exposing the LDH’s to the laboratory
environment (which contains H
2
O and CO
2
), reconstruction of their structure is
easily achieved by rehydration and CO
2
adsorption
43, 44
. Hibino et al.
42
, for
example, reported that the OH
-
and CO
3
2-
groups in the Mg-Al-CO
3
LDH could still
be detected by IR even after the sample had been treated at temperatures as high as
1000
o
C; normally, one would have expected all the OH
-
and CO
3
2-
groups in the
11
LDHs to disappear after treatment at this temperature. Their presence in the spectra
of Hibino et al.
42
is, likely, because these data were acquired ex-situ (room
temperature) rather than in-situ. Clearly, to study the changes in the LDH structure
and their functional groups at different temperatures, the use of in-situ techniques is
essential.
Recently, several studies have appeared using in-situ techniques to
characterize the thermal evolution of the structure of LDH’s, including such
techniques as high-temperature XRD (HTXRD)
45-47
, and infrared emission
spectroscopy (IES)
48-50
. For example, Kanezaki
45-47
in his studies using HTXRD
described the phase changes of Mg-Al-CO
3
LDH. In his experiments the sample
temperature was raised at 5
o
C/min, and HTXRD spectra were recorded at
temperature intervals of 20
o
C. Before each measurement, the sample temperature
was held constant for 15 min. According to the HTXRD observations, the structure
of the LDH that prevails at room temperature (Phase I) has a basal spacing of 7.8 Å.
Phase I begins to transform at higher temperatures (>180
o
C) into another phase
(Phase II) with a smaller basal spacing (6.6 Å), which, itself, begins to transform
into a third phase at temperatures exceeding 380
o
C. Kanezaki’s studies, however,
did not provide information on the changes that take place in the functional Mg-Al-
CO
3
LDH groups, and have not shed light on how the various structures and phases
develop with increasing temperature. Kanezaki
46
speculated, for example, that
Phase II developed through the thermal oxidation of interlayer carbonate in Phase I
by a nearby H
2
O molecule by the following reaction,
(CO
3
2-
)
interlayer
+ H
2
O 2(OH
-
)
interlayer
+ CO
2
↑
which, in addition, produces gaseous CO
2
and two interlayer OH
-
. According to
Kanezaki
46
, replacing CO
3
2-
with the less bulky OH
-
explains the reduction of the
basal spacing from 7.8 Å (Phase I, <180
o
C) to 6.6 Å (Phase II, 220 - 380
o
C).
Actually, as it will be discussed in chapter 2, in the temperature range 220 - 380
o
C
it is the OH
-
group rather than CO
3
2-
that disappears, making it doubtful that the
above reaction actually takes place; the reduction of basal spacing in the Mg-Al-
CO
3
LDH is, likely, due to the shrinkage of the interlayer space by the removal of
water.
Kloprogge & Frost
50
used in-situ IES to monitor the changes in the functional
groups during the thermal transformation of LDHs. This study indicates that major
changes, such as the disappearance of the OH
-
band, occur in the temperature range
12
13
350 - 400
o
C. The study of Kloprogge and Frost did not provide other in-situ
information on the structural changes and the products generated from the Mg-Al-
CO
3
LDH as a function of temperature; furthermore, the IES signal is weak at
temperatures below 200
o
C, where important structural changes also take place.
This has led Perez-Ramirez, Mul, Kapteijn & Moulijn
43
to question, (1) Kloprogge
and Frost’s interpretation of the thermogravimetric profiles; (2) the procedure
followed to carry out and interpret the in-situ IES measurements, and (3) their
assessment of the reaction mechanism based on the characterization technique.
Therefore, to elucidate the changes that occur in the structure, type, and number of
functional groups in the Mg-Al-CO
3
LDH at different temperatures, it is required to
utilize a combination of different in-situ techniques, which can provide
simultaneous information on the structural changes, the type/number of functional
groups present, and the various species generated.
The thermal evolution of the Mg-Al-CO
3
LDH structure under an inert
atmosphere is described in chapter 2. The results were obtained with the use of a
combination of several in-situ techniques, including diffuse reflectance infrared
Fourier transform spectroscopy (DRIFTS) for investigating changes in the
14
functional groups, TG/DTA for measuring the weight and energetic changes,
TG/MS for identifying the products generated, and HTXRD for in-situ monitoring
the structure evolution. And a schematic model was proposed to describe the
structural evolution of the Mg-Al-CO
3
LDH.
In Chapter 3, the thermal evolution of the Mg-Al-CO
3
LDH structure was
studied further, and the proposed structural model was validated under more
realistic reactive environments, in which the LDH materials may find eventual
applications. Also the sorption characteristics and thermal reversibility of these
materials were investigated under both inert and reactive atmospheres.
The diffusion characteristics and adsorption isotherm of carbon dioxide in
Mg-Al-CO
3
LDH were studied in chapter 4, with the aid of CO
2
uptake data
measured gravimetrically in the temperature range of 200 ~ 250
o
C. For the
diffusivity constant, the transient CO
2
uptake data were measured gravimetrically at
each elevated temperature, and then the diffusion constant was estimated by fitting
the experimental data to the solution of the relevant diffusion equation. In addition
to providing the experimental results, these observations are compared with
molecular dynamic simulation results investigated in a parallel research project.
15
Adsorption isotherm data were also acquired by the gravimetric uptake
measurements, and the experimental data were fitted with the Langmuir equation
and various empirical adsorption isotherm equations. Also, the particle size effect of
LDH for CO
2
uptake at 200
o
C was investigated as well.
In chapter 5, various methods for the preparation of LDH membranes were
studied, and the transport properties of prepared LDH membranes were investigated.
The LDH membranes are intended to be utilized as facilitated transport membranes
to separate CO
2
from various gas mixtures.
16
Chapter 1 References
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with metal coordination compounds and oxometalates. Coordination
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2. Roy, A. de; Forano, C.; Malki, K. E.; Besse, J. P. Anionic clays: Trends in
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3. Manasse, E., Rocce eritree e di aden della collezione issel. Atti Soc. Toscana
Sc. Nat., Proc. Verb., 1915, 24, 92.
4. Aminoff, G.; Broomé, B., Contributions to the knowledge of the mineral
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basischen Doppelsalze, III. Über Magnesium-Aluminiumdoppelhydroxyd.
Helv. Chim. Acta, 1942, 25, 131.
6. Feitknecht, W. The formation of double hydroxide between bivalent and
trivalent metals. Helv. Chim. Acta, 1942, 25, 555.
7. Allmann, R. The crystal structure of pyroaurite. Acta Crystallogr., 1968,
B24, 972.
8. Taylor, H. F. W. Segregation and cation-ordering in sjögrenite and pyroaurite.
Miner. Mag., 1969, 37, 338.
9. Allmann, R.; Jepsen, H. P. Die Struktur des Hydrotalkits. Neues. Jahrb.
Miner. Monatsh., 1969, 12, 545.
17
10. Allmann, R. Doppelschichstrukturen mit brucitähnlichen Schichtionen
[Me(II)
1-x
Me(III)
x
(OH)
2
]
x+
. Chimia, 1970, 24, 99.
11. Taylor, H. F. W. Crystal structure of some double hydroxide minerals. Miner.
Mag., 1973, 39, 377.
12. Reichle, W. T. Synthesis of anionic minerals (mixed metal hydroxides,
hydrotalcite). Solid State Ionics, 1986, 22, 135.
13. Cavini, F.; Trifiro, E.; Vaccari, A. Hydrotalcite-type anionic clays –
preparation, properties and applications. Cat. Today, 1991, 11, 173.
14. Allmann, R., Refinement of the hybrid layer structure hexahydroxoalumino-
dicalcium hemisulfate trihydrate [Ca
2
Al(OH)
6
]
+
.[1/2·SO
4
.3H
2
O]
-
Neues.
Jahrb. Miner. Monatsh., 1977, 3, 136.
15. Drits, V. A.; Sokolova, T. N.; Sokolova, G. V.; Cherkashin, V. I. New
members of the hydrotalcite-manasseite group. Clays and Clay Minerals,
1987, 35, 401.
16. Tsuji, M.; Mao, G.; Yoshida, T.; Tamaura, Y. Hydrotalcites with an extended
Al
3+
substitution: Synthesis, simultaneous TG-TDA-MS study, and their
CO
2
adsorption behaviors. J. Mater. Res., 1993, 8(5), 1137.
17. Brindley, G. W.; Kikkawa, S. A crystal-chemical study of magnesium,
aluminum and nickel, aluminum hydroxy-perchlorates and hydroxy-
carbonates. Amer. Min., 1979, 64, 836.
18. Miyata, S. Physicochemical properties of synthetic hydrotalcites in relation
to composition. Clays and Clay Minerals, 1980, 28, 50.
19. Sato, T.; Fujita, H.; Endo, T.; Shimada, M. Synthesis of hydrotalcite-like
compounds and their physico-chemical properties. React. of Solids, 1988, 5,
219.
18
20. Pausch, I.; Lohse, H. H.; Schurmann, K; Allmann, R. Synthesis of
disordered and Al-rich hydrotalcite-like compounds. Clays and Clay
Minerals, 1986, 34, 507.
21. Miyata, S. The syntheses of hydrotalcite-like compounds and their structures
and physico-chemical properties. I. The systems magnesium(2+)-
aluminum(3+)-nitrate(1-), -chloride(1-) and -perchlorate(1-), nickel(2+)-
aluminum(3+)-chloride(1-), and zinc(2+)-aluminum(3+)-chloride(1-). Clays
and Clay Minerals, 1975, 23, 369.
22. Miyata, S. Anion-exchange properties of hydrotalcite-like compounds.
Clays and Clay Minerals, 1983, 31, 305.
23. Miyata, S.; Okada, A. Synthesis of hydrotalcite-like compounds and their
physico-chemical properties-the systems Mg
2+
-Al
3+
-SO
4
2-
and Mg
2+
-Al
3+
-
CrO
4
2-
. Clays and Clay Minerals, 1977, 25, 14.
24. Miyata, S.; Kumura, T. Synthesis of new hydrotalcite-like compounds and
their physico-chemical properties. Chem. Lett., 1973, 843.
25. Carlino, S. The intercalation of carboxylic acids into layered double
hydroxide: a critical evaluation and review of the different methods. Solid
State Ionics, 1997, 98, 73.
26. Miyata, S.; Kumura, T.; Hattori, H.; Tanabe, K. Physicochemical properties
and structure of magnesia-alumina. Nippon Kagaku Zasshi, 1971, 92, 514.
27. Goswamee, R. L.; Sengupta, P.; Bhattacharyya, K. G.; Dutta, D. K.
Adsorption of Cr(VI) in layered double hydroxides. Appl. Clay Sci., 1998
13(1), 21.
19
28. Pavan, P. C.; Gomes, G. de A.; Valim, J. B. Adsorption of sodium dodecyl
sulfate on layered double hydroxides. Microporous and Mesoporous
Materials, 1998, 21(4-6), 659.
29. Shin, H. S.; Kim, M. J.; Nam, S. Y.; Moon, H. C. Phosphorus removal by
hydrotalcite-like compounds (HTLCs). Water Science and Technology, 1996,
34(1-2), 161.
30. Ding, Y .; Alpay, E. Equilibria and kinetics of CO
2
adsorption on hydrotalcite
adsorbent. Chem. Eng. Sci., 2000, 55, 3461.
31. Yong, Z.; Mata V.; Rodrigues, A. E. Adsorption of carbon dioxide onto
hydrotalcite-like compounds (HTLCs) at high temperatures. Ind. Eng. Chem.
Res., 2001, 40, 204.
32. Alejandre, A.; Medina, F.; Rodriguez, X.; Salagre, P.; Cesteros, Y .; Sueiras, J.
E. Cu/Ni/Al layered double hydroxides as precursors of catalysts for the
wet air oxidation of phenol aqueous solutions. Appl. Catal. A:
Environmental, 2001, 30(1-2), 195.
33. Bahranowski, K.; Bueno, G.; Corberan, V. C.; Kooli, F.; Serwicka, E.M.;
Valenzuela R. X.; Wcislo, K. Oxidative dehydrogenation of propane over
calcined vanadate-exchanged Mg, Al-layered double hydroxides. Appl.
Catal. A: General, 1999, 185(1), 65.
34. Pinnavaia, T. J.; Gardner, E. On the nature of selective olefin oxidation
catalysts derived from molybdate- and tungstate-intercalated layered double
hydroxides. Appl. Catal. A: General, 1998, 167(1), 65.
35. Auer, S. M.; Grunwaldt, J. D.; Koppel, R. A.; Baiker, A. Reduction of 4-
nitrotoluene over Fe-Mg-Al lamellar double hydroxides. J. Molecular Catal.
A: Chemical, 1999, 139(2-3), 305.
20
36. Malherbe, F.; Depege, C.; Forano, C.; Besse, J. P.; Atkins, M. P.; Sharma,
B.; Wade S. R. Alkoxylation reaction catalyzed by layered double
hydroxides. Appl. Clay Sci., 1998, 13(5-6), 451.
37. Santhanalakshmi, J.; Raja, T. Selective N-methylation of aniline by calcined
Mg
II
-Al
III
layered double hydroxides. Appl. Catal. A: General, 1996, 147(1),
69.
38. Tichit, D.; Bennani, M. N.; Figueras, F.; Tessier, R.; Kervennal, J. Aldol
condensation of acetone over layered double hydroxides of the meixnerite
type. Appl. Clay Sci., 1998, 13(5-6), 401.
39. Ding, Y.; Alpay, E. (2000). Adsorption-enhanced steam-methane reforming.
Chem. Eng. Sci., 2000, 55, 3929.
40. Hufton, J.R.; Mayorga, S.; Sirkar S. Sorption-Enhanced Reaction Process
for Hydrogen Production, AIChE J., 1999, 45, 248.
41. Kim, Y.; Yang, W.; Liu, P.K.T.; Sahimi, M.; Tsotsis, T. T. Layered double
hydroxide (LDH) materials for CO
2
adsorption and separation, presented at
A&WMA’s 96
th
Annual Conference & Exhibition, June 22-26, San Diego,
CA, 2003.
42. Hibino, T.; Yamashita, Y .; Kosuge, K.; Tsunashima, A. Decarbonation
behavior of Mg-Al-CO
3
hydrotalcite-like compounds during heat treatment.
Clays and Clay Min., 1995, 43(4), 427.
43. Perez-Ramirez, J.; Mul, G.; Kapteijn, F.; Moulijn, J. A. (2000). Comments
on “Infrared emission spectroscopic study of the thermal transformation of
Mg-, Ni-, and Co-hydrotalcite catalysts” [Appl. Catal A:General,
184(1999)61]. Appl. Catal, A: General, 2000, 204, 256.
21
44. Puttaswamy, N. S.; Kamath, P. V . Reversible thermal behavior of layered
double hydroxides: a thermogravimetric study. J. Mater. Chem., 1997, 7(9),
1941.
45. Kanezaki, E. Effect of atomic ratio Mg/Al in layers of Mg and Al layered
double hydroxide on thermal stability of hydrotalcite-like layered structure
by means of in-situ high temperature powder X-ray diffraction. Mat. Res.
Bull., 1998, 33(5), 773.
46. Kanezaki, E. Thermal behavior of the hydrotalcite-like layered structure of
Mg and Al-layered double hydroxides with interlayer carbonate by means
of in-situ powder HTXRD and DTA/TG. Sol. State Ion., 1998, 106(3-4), 279.
47. Kanezaki, E. Thermally induced metastable solid phase of Mg/Al layered
double hydroxides by means of in-situ high temperature powder X-ray
diffraction. J. Mater. Sci. Lett., 1998, 17(5), 371.
48. Kloprogge, J. T.; Frost, R. L. Fourier transform infrared and Raman
spectroscopic study of the local structure of Mg-, Ni-, and Co-hydrotalcites.
J. Solid State Chem., 1999, 146, 506.
49. Kloprogge, J. T.; Frost, R. L. Infrared emission spectroscopic study of the
dehydroxylation of synthetic Mg/Al and Mg/Zn/Al-hydrotalcites. Phys.
Chem. Chem. Phys., 1999, 1, 1641.
50. Kloprogge, J. T.; Frost, R. L. Infrared emission spectroscopic study of the
thermal transformation of Mg-, Ni-, and Co-hydrotalcite catalysts. Appl.
Catal A:General, 1999, 184, 61.
22
Chapter 2
In-situ Thermal Evolution Study of Mg-Al-CO
3
LDH
2.1 Introduction
In this chapter, the thermal evolution of the Mg-Al-CO
3
LDH structure is
studied using various in-situ techniques, including DRIFTS, TG/DTA, thermal
gravimetric/microbalance-mass spectrometer (TG/MB-MS), and HTXRD. DRIFTS
is a sensitive and powerful technique that has been utilized in this study to monitor
in-situ the changes of functional groups in the Mg-Al-CO
3
LDH samples as a
function of temperature and other experimental conditions. The changes can be
quantified using the Kubelka-Munk theory, which predicts a linear relationship
between spectral intensity and the concentration of a sample under conditions of a
constant scattering coefficient and infinite sample dilution in a non-absorbing
matrix. In this study only functional group changes in the same material are
monitored with increasing temperature, so the scattering coefficient is assumed
constant. As for the second requirement, the sample is diluted with a non-absorbing
KBr matrix.
23
In-situ TG/DTA techniques were used to monitor the weight and energetic
changes of the Mg-Al-CO
3
LDH as function of temperature. Combined with
DRIFTS, TG/TDA provides additional quantitative insight into functional group
changes at various temperatures. In-situ TG/MB-MS techniques were used to
monitor the weight changes of the Mg-Al-CO
3
LDH samples and the gaseous
products generated during their thermal evolution as a function of temperature and
other conditions. Combined with DRIFTS, TG/MB-MS provides additional
quantitative insight into functional group changes at various conditions.
In-situ HTXRD was used to detect the structural changes of Mg-Al-CO
3
LDH as a function of temperature. With the combination of various in-situ
techniques, a model of thermal evolution of the Mg-Al-CO
3
LDH structure was
proposed in conclusion.
2.2 Experimental
The LDH sample was provided by the Media and Process Technology, Inc.,
of Pittsburgh, PA. Its composition is Mg
0.71
Al
0.29
(OH)
2
(CO
3
)
0.15
.
0.46(H
2
O)
(hereinafter referred to as LDH1), as determined by ICP and TGA. DRIFTS spectra
24
were recorded in-situ using a Genesis II (Mattson, FT-IR) instrument equipped with
a DRIFTS COLLECTOR
TM
II chamber (SpectraTech, Inc.) capable of operating
under high temperatures (up to 900
o
C) and pressures (up to 1500 psi). The chamber
windows are made of ZnSe to withstand these conditions, and to allow for better
infrared transmission. A controller is used to control the chamber temperature
utilizing a ceramic heater and a thermocouple in intimate contact with the sample.
With this chamber, the temperature and the sample environment can be easily
controlled. The experimental operating conditions were a DRIFTS scan-range from
4000 cm
-1
to 500 cm
-1
, scan numbers 16, and a scan resolution of 2 cm
-1
. The
spectra were calibrated for background with KBr. To obtain a strong signal intensity
and better resolution for quantitative measurements, the sample was first ground to
2-10 μm, diluted with KBr to 5~10% wt., placed in the sample cup and leveled with
a spatula. Experiments were carried out in an inert gas atmosphere. The temperature
was raised at a rate of 0.5
o
C/sec, and spectra were recorded at different
temperatures about 20
o
C apart.
For the thermal evolution of LDH’s study, the thermal gravimetric anlysis
(TGA) curves were recorded on a Pyris 1 TGA HT instrument (PE Company) by
25
heating the sample from 50 to 600
o
C in an Ar atmosphere, at a rate of 5
o
C/min,
and an Ar flow rate of 20 ml/min. Differential thermal analysis (DTA) was
performed with a DTA 7 instrument (PE Company) at the same conditions as TGA.
For in-situ TG/MB-MS studies, TG curve was recorded on a Cahn TGA
121 instrument (thermo Cahn), and a custom-made TG/MB-MS instrument using a
MKS UTI 100C Precision Gas Analyzer. HTXRD experiments were carried out in
the temperature range 30 – 650
o
C under vacuum (10
-2
Torr) using a Siemens D-
5000 X-ray diffractometer equipped with a Buhler high-temperature chamber HDK
1.4. The chamber is made of stainless steel with Be-windows for X-ray transmission
and Ta thermal shields acting as the isothermal block. The sample and its
surroundings were heated by Pd heaters at a rate of 0.5
o
C/s; the spectra were
recorded at different temperatures, typically 20
o
C apart. The sample was
equilibrated at any given temperature for 30 min. The Cu K
α
line was used for the
X-ray source with a monochromator positioned in front of the detector. Scans were
performed over a 2θ range from 5
o
to 75
o
.
Figure 2.1 Comparison between the DRIFTS and FTIR results of Mg-Al-CO
3
LDH.
26
27
2.3 Results and Discussion
Fig. 2.1 compares the DRIFTS and FT-IR spectra of the LDH1 at room
temperature in Ar. It shows that all key bands in the DRIFTS spectra for the LDH1
are in the same position as those in the FT-IR spectra. The intensities of the
DRIFTS signals are sufficiently strong to clearly identify all the important
functional groups. In accordance with prior FT-IR studies
1,2
, the DRIFTS bands
have been assigned to the following groups:
(1) The DRIFTS signal at ~3470 cm
-1
is due to the OH
-
group vibration in the
Mg-Al-CO
3
LDH sample;
(2) the DRIFTS signal at ~3070 cm
-1
is due to hydrogen bonding between water
and the carbonate species in the interlayer space of the Mg-Al-CO
3
LDH
sample;
(3) the DRIFTS signal at ~1620 cm
-1
is due to the H
2
O bending vibration of
interlayer water in the Mg-Al-CO
3
LDH sample;
(4) The DRIFTS signals at ν
3
= 1370 cm
-1
,
ν
2
= 940 cm
-1
, and ν
4
= 680 cm
-1
at
room temperature are due to the CO
3
2-
group vibration bands in the Mg-Al-
CO
3
LDH sample; the CO
3
2-
group in the hydrotalcite behaves more like it
28
would in a water solution, in which the vibration bands of CO
3
2-
are
observed at ν
3
= 1415 cm
-1
, ν
2
= 880 cm
-1
, and ν
4
= 680 cm
-1
. No ν
1
mode
vibration at ~ 1080 cm-1 and no splitting of the ν
3
band are observed. The
splitting of the ν
3
band and the ν
1
mode vibration band are usually generated
from the symmetry degradation, which results from the interaction between
CO
3
2-
and Mg
2+
. This means that CO
3
2-
in the Mg-Al-CO
3
LDH sample at
room temperature has a very weak, if any, direct interaction with positive
ions, such as Mg
2+
and Al
3+
.
Based on the above band assignments, one can use the in-situ DRIFTS technique to
characterize the LDH structural evolution process, especially the changes of the
functional groups of the Mg-Al-CO
3
LDH sample as a function of temperature. The
in-situ DRIFTS results are shown in Fig. 2.2, for which the LDH1 was treated in Ar.
In these experiments, starting from room temperature, the sample temperature was
increased at a rate of 0.5
o
C/s. Every 20
o
C or so the temperature increase would stop,
and the DRIFTS spectra would be recorded after keeping the sample isothermal for
a period of ~2 min. The testing was continued until the temperature of the sample
had reached 580
o
C. From Fig. 2.2, one can draw the following conclusions:
560
o
C
500
o
C
480
o
C
460
o
C
440
o
C
420
o
C
400
o
C
380
o
C
360
o
C
340
o
C
320
o
C
300
o
C
280
o
C
260
o
C
240
o
C
220
o
C
200
o
C
190
o
C
170
o
C
150
o
C
120
o
C
100
o
C
90
o
C
70
o
C
50
o
C
RT
4500 4000 3500 3000 2500 2000 1500 1000 500 0
Wavenumber, cm
-1
3470cm
-1
1370cm
-1
3070cm
-1
1620cm
-1
1530cm
-1
1350cm
-1
1371cm
-1
1090cm
-1
Kubelk Munk Units
Figure 2.2 In-situ DRIFTS of Mg-Al-CO
3
LDH as a function of temperature.
29
30
(1) The intensities of the interlayer water bands at 3070 cm
-1
and 1620 cm
-1
gradually decrease with increasing temperature, and disappear around 190
o
C.
This means that increasing amounts of interlayer water in the LDH1 are
removed with increasing temperature. In the presence of Ar beyond 190
o
C
the water that remains in the sample is below the detection limit of the
DRIFTS instrument.
(2) The intensity of the OH
-
vibration band at 3470 cm
-1
begins to decrease at
190
o
C, and completely disappears at 440
o
C. This suggests that in the
presence of Ar dehydroxylation of LDH1 begins at 190
o
C; by the time the
temperature reaches 440
o
C the concentration of the remaining OH
-
groups
is below the detection limit of the DRIFTS instrument.
(3) The band at 1370 cm
-1
for the CO
3
2-
ν
3
vibration begins to decrease in size
as the temperature increases, and also shifts to ~1350 cm
-1
. Gradually a band
at 1530 cm
-1
begins to form at temperatures higher than 170
o
C. The band
size at the lower wave number (~1350 cm
-1
) decreases as the temperature
increases (and so are the peaks at 940 and 680 cm
-1
). This is to be expected
as the amount of interlayer water decreases and, as a result, the CO
3
2-
group
31
begins to interact more strongly with the backbone of the hydrotalcite itself.
For comparison purposes, Fig. 2.3 shows the DRIFTS spectrum of MgCO
3.
The behavior of the two bands at 1499 and 1425 cm
-1
qualitatively mirrors
that of the bands at 1530 and 1350 cm
-1
. On the other hand, the state of
CO
3
2-
in the hydrotalcite is distinctly different from that in MgCO
3
. Note
that bands corresponding to the ν
3
vibration of CO
3
2-
in MgCO
3
are at
different positions, namely 1499 cm
-1
and 1425 cm
-1
. In Fig. 3, one observes
the ν
1
vibration band of CO
3
2-
in MgCO
3
. This band is not present in Fig. 2.
At higher temperatures all peaks corresponding to CO
3
2-
species in Mg-Al-
CO
3
LDH disappear. This is consistent with the MS data (see the discussion
below), which show that by that temperature all CO
3
2-
has left the
hydrotalcite structure as CO
2
.
Fig. 2.4 summarizes the in-situ DRIFTS results of Fig. 2.2, indicating the fraction
of each particular species that is removed at a given temperature. The calculations
are based on the peak area of the corresponding species recorded at different
temperatures. The 3470 cm
-1
band represents the OH
-
group, and the 3070 cm
-1
and 1620 cm
-1
band represent the interlayer water. For CO
3
2-
, since the DRIFTS
4500 4 000 3500 3 000 2500 2000 1500 1000 5 00 0
Kubelka Munk Units
W avenum ber,c m
-1
1499cm
-1
1425cm
-1
1100cm
-1
560
o
C
520
o
C
500
o
C
480
o
C
460
o
C
440
o
C
420
o
C
400
o
C
300
o
C
200
o
C
100
o
C
RT
850cm
-1
700cm
-1
Figure 2.3 In-situ DRIFTS of MgCO
3
as a function of temperature.
32
0 100 200 300 400 500 600
0
20
40
60
80
100
Interlayer H
2
O
OH
-
Fraction of Removed Species, %
Temperatures,
o
C
CO
3
2-
Figure 2.4 Fraction of species removed from Mg-Al-CO
3
LDH as a function of
temperature.
33
34
band at 1370 cm
-1
at lower temperatures was split into two DRIFTS bands, i.e.,
1530 cm
-1
and 1350 cm
-1
, in the quantitative analysis, the fraction removed was
calculated only at temperatures higher than 180
o
C using the 1530 cm
-1
and
1350 cm
-1
bands to represent CO
3
2-
. 100% removal in Fig. 2.4 corresponds to the
point when the particular band is no longer detectable by DRIFTS. It can be seen in
Fig. 2.4 that the interlayer water starts to disappear at 70
o
C, with the band no
longer detectable at 190
o
C. For the OH
-
groups, there is a noticeable difference in
the slope of the removal rate between the temperature ranges 190 ~ 250
o
C and 250
~ 440
o
C, suggesting, perhaps, that in these two temperature ranges the OH
-
group
finds itself in two different environments in the Mg-Al-CO
3
LDH sample. Fig. 2.4
also indicates that most of the CO
3
2-
was removed in the temperature range 390 ~
580
o
C. These results were also confirmed by in-situ MS and TG/MS experiments.
Fig. 2.5, for example, shows the results of in-situ MS analysis. It indicates that H
2
O
is continuously removed from the Mg-Al-CO
3
LDH1 sample until 420
o
C, which is
consistent with the results of DRIFTS, shown in Fig. 2.4. Being discussed above,
in-situ DRIFTS is a powerful technique to monitor the changes of the functional
groups in Mg-Al-CO
3
LDH sample as the temperature changes.
0 100 200 300 400 500 600 700
0
20
40
60
80
100
Fraction of H2O Removed, %
Temperature,
o
C
Figure 2.5 In-situ MS of Mg-Al-CO
3
LDH as a function of temperature.
35
Figure 2.6 In-situ TG/MS of Mg-Al-CO
3
LDH as a function of temperature.
36
37
In-situ TG/MS and TG/DTA, as previously noted, provide complimentary
technical information. Fig. 2.6 shows, for example, the results of in-situ TG/MS
analysis. It indicates that a relatively small amount of CO
2
(~11.5%) from the
LDH1 was detected in the temperature range of 190 ~ 390
o
C, and most of the CO
2
(~88.5%) from the Mg-Al-CO
3
LDH was detected in the temperature range of 390
~ 580
o
C. These results are consistent with the results of DRIFTS, shown in Fig. 2.4.
Fig. 2.7 shows the corresponding results of the TG/DTA analysis of the Mg-Al-CO
3
LDH1 sample as a function of temperatures in an inert gas atmosphere. Combining
the in-situ DRIFTS results with the TG/MS observations, one can draw quantitative
conclusions concerning the behavior of the various functional groups. One can
conclude, for example, that in Fig. 6 the first weight loss, ~13.5% by weight, in the
temperature range of 70 ~ 190
o
C should be mostly due to the interlayer water in the
Mg-Al-CO
3
LDH, together with relatively smaller amounts of CO
2
and H
2
O
resulting from the desorption of the OH
-
group. The theoretical weight fraction of
the interlayer water in Mg-Al-CO
3
LDH1 with the reported composition of
Mg
0.71
Al
0.29
(OH)
2
(CO
3
)
0.15
.0.46H
2
O is calculated to be 10.84 % by weight. The
difference (2.66 % by weight) between the experimental value (13.5 % by weight)
38
and the theoretical value (10.84 % by weight) should be attributed to the
contribution of CO
2
and H
2
O from the OH
-
group, because there are, indeed, small
amounts of CO
2
and H
2
O from the OH
-
group that were removed in the temperature
range of 70
o
C to 190
o
C, as shown in Figs. 2.4 and 2.6. The DTA results (Fig. 2.7)
show no distinct peaks in the temperature range of 70 ~ 190
o
C, which is consistent
with the conclusion that the removal of interlayer water, which is physically
adsorbed in the nano-slits between the layers of Mg-Al-CO
3
LDH, is predominantly
responsible for the sample weight loss in this region.
The second distinct weight loss region of 22.17 wt.% (Fig. 2.7) is in the
temperature range 190 ~ 400
o
C. It is accompanied by the evolution of H
2
O and
CO
2
species (see Figs. 2.4 and 2.6), and various heat flows (Fig. 2.7). The first
small heat flow shoulder peak (endothermic) is centered at ~205
o
C, and seems to
be coincident with the change of the 1370 cm
-1
peak in Fig. 2.2. The mass
spectrometric data indicate the evolution of CO
2
and some H
2
O in this region. From
~190 to ~220
o
C the weight change is ~ 0.2 %, of which 0.08 % corresponds to CO
2
and 0.12 % to H
2
O. The small heat flows are indicative that the OH
-
and CO
3
2-
that
are exchanged in this region may be physically bound in the LDH sample. Between
0 50 100 150 200 250 300 350 400 450 500 550 600 650
60
70
80
90
100
Temperature,
o
C
Weight,%
-15
-10
-5
0
5
10
Heat Flow Endo Down,mW
DTA
TG
Figure 2.7 In-situ TG/DTA of Mg-Al-CO
3
LDH as a function of temperature.
39
0 50 100 150 200 250 300 350 400 450 500 550 600 650 700
60
70
80
90
100
Temperature,
o
C
Weight,%
-40
-30
-20
-10
0
10
20
Heat Flow Endo Down,mW
DTA
TG
Figure 2.8 In-situ TG/DTA of Al(OH)
3
as a function of temperature.
40
0 50 100 150 200 250 300 350 400 450 500 550 600 650 700
60
70
80
90
100
Temperature,
o
C
Weight,%
-40
-30
-20
-10
0
10
20
Heat Flow Endo Down,mW
DTA
TG
Figure 2.9 In-situ TG/DTA of Mg(OH)
2
as a function of temperature.
41
42
~220 and ~ 400
o
C, there is a 21.97 % weight change; there are also two
characteristic endothermic flows in this region, one centered at ~255
o
C and the
other at ~340
o
C. The DTA data seem to be consistent with the DRIFTS data, which
indicate potentially two types of OH
-
groups that are removed in the same
approximate region of temperatures. In order to clarify the properties of these two
types of OH
-
groups, TG/DTA experiments with pure Al(OH)
3
and pure Mg(OH)
2
samples were also performed; the results are shown in Figs. 2.8 and 2.9. Fig. 2.8
indicates that the OH
-
group associated with Al
3+
is lost in the temperature range
190 ~ 300
o
C, and two heat flow peaks (endothermic) are observed in the same
region. Fig. 2.9 shows that the OH
-
group associated with Mg
2+
is removed in the
temperature range 300 ~ 405
o
C; only one heat flow peak (endothermic) is observed
at this case. Comparing Fig. 2.7 with Figs. 2.8 and 2.9, one observes that the
thermal behavior of Mg-Al-CO
3
LDH1 in Fig. 2.7 is a composite of the thermal
behavior of Al(OH)
3
and Mg(OH)
2
. The weight loss in the temperature range 190 ~
300
o
C is ~8.93 %. Of this 0.19 % corresponds to CO
2
and 8.74 % to the removal of
H
2
O, which likely results from Al-(OH)-Mg OH
-
groups. Between 300
o
C and 405
o
C the weight loss is ~13.24 %. Of this 0.72 % corresponds to CO
2
and 12.52 % to
43
the removal of H
2
O, which for this case, more likely, results from Mg-(OH)-Mg
OH
-
groups. Assuming that Al
3+
associates only with OH
-
groups, the total weight
change one would expect due to the evolution of H
2
O from OH
-
is 10.37 %. This
compares favorably with the 8.74 % change due to water measured in the region of
190 ~ 300
o
C. The weight change one would expect due to the evolution of H
2
O
from the OH
-
groups associated with Mg
2+
is 13.20 %. The experimental value in
the region 300 ~ 405
o
C is 12.52 %, which compares favorably with the calculated
value.
The third distinct weight loss region is between 405 and 580
o
C, which the
total weight loss is 7.33 %. It can be attributed to the removal CO
2
from CO
3
2-
in
the Mg-Al-CO
3
LDH, because almost no water was detected in this temperature
range by in-situ MS. This brings the total weight loss due to CO
2
to 8.32 %, as
compared with the calculated value of 7.15%.
Based on the results of in-situ DRIFTS coupled with the in-situ TG/DTA
and TG/MS studies, a model is proposed for the evolution of the structure of the
Mg-Al-CO
3
LDH sample under an inert atmosphere (Fig. 2.10). In Fig. 2.10 five
distinct stages were identified during the thermal evolution of this particular Mg-Al-
44
CO
3
LDH sample. The original Mg-Al-CO
3
LDH sample is referred to as Stage A;
Stage B develops from Stage A by the removal of the loosely held interlayer water
in the temperature range 70 ~ 190
o
C; Stage C evolves from Stage B by the removal
of OH
-
groups, likely bonded in a bridge Al-(OH)-Mg configuration, in the
temperature range 190 ~ 300
o
C; Stage D was achieved from Stage C by the
removal of OH
-
groups, likely bonded mostly with Mg
2+
(Mg-(OH)-Mg) in the
temperature range 300 ~ 405
o
C; Stage E is obtained by the decarbonation of Stage
D in the temperature range 405 ~ 580
o
C.
It is also interesting to investigate how the changes in the number and type
of functional groups with temperature manifest themselves in changes in the crystal
structure of the Mg-Al-CO
3
LDH1 sample. In order to study this, in-situ HTXRD
has been utilized. Fig. 2.11 shows the HTXRD patterns of the Mg-Al-CO
3
LDH
treated at different temperatures. There are five temperature regions that can be
identified based on the HTXRD patterns, i.e., (1) 30
o
C ≤ T ≤ 140
o
C, (2) 140
o
C ≤
T ≤ 180
o
C, (3) 180
o
C ≤ T ≤ 280
o
C, (4) 280
o
C ≤ T ≤ 360
o
C, and (5) 360
o
C ≤ T ≤
650
o
C.
C
E
Mg
0.71
Al
0.29
(OH)
2
(CO
3
2-
)
0.15
0.46H
2
O
Interlayer Water
-0.44H
2
O from OH
-
Group:
Al-OH
-
: 190
o
C-280
o
C
Mg
0.71
Al
0.29
O
0.44
(OH)
1.12
(CO
3
2-
)
0.15
-CO
2
405
o
C-580
o
C
CO
3
2-
Solid Solution
MgO + Al
2
O
3
Amorphous Phase
Brucite-like Layers
A
B
D
Mg
0.71
Al
0.29
(OH)
2
(CO
3
2-
)
0.15
-0.46H
2
O
70
o
C-190
o
C
OH
-
Group in
Mg-OH
-
-0.56H
2
O
280
o
C - 405
o
C
Figure 2.10 The thermal evolution of Mg-Al-CO
3
LDH as a function of
temperature.
45
10 20 30 40 50 60 70
30
o
C
70
o
C
140
o
C
160
o
C
120
o
C
100
o
C
200
o
C
240
o
C
180
o
C
220
o
C
260
o
C
280
o
C
300
o
C
320
o
C
340
o
C
360
o
C
2θ/Deg
Figure 2.11 In-situ HTXRD of Mg-Al-CO
3
LDH as a function of temperature.
46
47
For the first temperature region 30 ~ 140
o
C, all the HTXRD patterns are
attributed to the hydrotalcite-like structure, as reported by Kanezaki
3
. The
intensities of HTXRD lines are strong, and stay almost unchanged with increasing
temperature, meaning that the LDH structure is well preserved in this temperature
region. More detailed analysis of HTXRD data has shown that the basal spacing of
LDH has decreased from 7.63 Å to 7.5 Å, as the temperature increased from 30 to
140
o
C (Table 2.1).
In the second temperature region 140 ~ 180
o
C, the strong X-ray diffraction
line at 11.6 starts shifting towards ~12
o
, and a new line appears around 13.4
o
, which
intensifies with increasing temperature. This X-ray diffraction line was indexed as
001 by Kanezaki
3
. Based on the X-ray diffraction spectrum one can identify two
different co-existing crystal phases of Mg-Al-CO
3
LDH: Phase I with a basal
spacing ranging from 7.5 Å to 7.3 Å, and Phase II with basal spacing of ~ 6.6 Å.
The basal spacing of Phase I is approximately equal to the sum of the thickness of
one layer of Mg-Al-CO
3
LDH (4.8 Å) and the interlayer distance of Mg-Al-CO
3
LDH (~3.0 Å), as reported, for example, by Cavini, Trifiro, & Vaccari
4
. By
assuming the same thickness of one layer of Mg-Al-CO
3
LDH (4.8 Å) for both
48
phases the interlayer distance of Phase II of Mg-Al-CO
3
LDH is calculated to be
1.79 Å, a decrease from the 2.83 Å interlayer distance of Phase I. The decrease of
the interlayer distance can be attributed to the shrinkage of the layers of Mg-Al-CO
3
LDH due to the removal of interlayer H
2
O, since only interlayer H
2
O was removed
(DRIFTS and TG/DTA results) in the temperature range 70 ~ 190
o
C.
Table 2.1 The changes in the basal spacing of Mg-Al-CO
3
LDH with temperature
calculated from the HTXRD patterns.
T,
o
C 30 50 70 100 120 140 160 180
2θ 11.6 11.6 11.6 11.69 11.69 11.79 13.4 12.06 13.4 12.06 13.4
d, Å 7.63 7.63 7.63 7.56 7.56 7.50 6.59 7.33 6.59 7.33 6.59
Table 2.2 The changes in the basal spacing of Mg-Al-CO
3
LDH with temperature
calculated from the HTXRD patterns.
T,
o
C 200 220 240 260 280 300 320 340 360
2θ 13.5 13.6 13.8 13.9 13.9 14.0 14.0 14.0 14.0
d, Å 6.54 6.50 6.40 6.36 6.36 6.31 6.31 6.31 6.31
49
In the third temperature region 180 ~ 280
o
C, the X-ray diffraction line at
~12
o
gradually disappears, and the line at ~13.4
o
becomes stronger and shifts towards
13.9
o
simultaneously. This means that, in this region, Phase I transforms to Phase II.
With increasing temperature, the basal spacing of Phase II decreases from 6.59 Å to
6.36 Å. Simultaneously, the other X-ray diffraction lines of hydrotalcite begin to
weaken in this temperature region, but the hydrotalcite structure stays intact though
the interlayer spacing has decreased. This is consistent with the DRIFTS and
TG/TDA data, which indicate removal of Al-(OH)-Mg OH
-
groups, while the OH
-
/CO
3
2-
groups likely bonded with Mg
2+
still remain intact; the LDH, thus, retains the
brucite-like structure, as shown in Fig. 2.10 (Stage C). The decrease of the basal
spacing is attributed to the decrease in the thickness of the Mg-Al-CO
3
LDH layer
due to the removal of Al-(OH)-Mg OH
-
groups.
In the fourth temperature region 280 ~ 360
o
C, the intensity of the
diffraction line at ~14
o
decreases, the line disappearing when the temperature
exceeds 360
o
C. This means that the layered structure of Mg-Al-CO
3
LDH falls
apart above 360
o
C, consistent with the observation that the OH
-
/CO
3
2-
groups
(likely bonded with Mg
2+
) begin to leave at 280
o
C and are completely removed at
50
405
o
C (Stage D), as shown in Fig. 2.10. Removing these groups results in a
collapse of the brucite-like layers and the Mg-Al-CO
3
LDH layered structure.
For the fifth temperature region 360 - 650
o
C, the HTXRD results for
temperatures higher than 360
o
C are not shown in Fig. 2.11, because all these
patterns are similar to those shown for 360
o
C. No HTXRD peaks are observed in
this temperature region. At this stage (Fig. 2.10) a solid solution of MgO and Al
2
O
3
is obtained by decarbonation of the Stage D compound in the temperature range 410
~ 580
o
C.
2.4 Conclusions
In-situ DRIFTS, DTA, TG/MS and HTXRD have been used in order to
investigate the evolution of the structure of Mg-Al-CO
3
LDH under an inert
atmosphere as a function of temperature. Based on these studies, the following
conclusions may be drawn:
(1) In the temperature range 70 ~ 190
o
C, loosely held interlayer water is lost,
and there are two different co-existing crystal phases of Mg-Al-CO
3
LDH present, i.e., Phase I with a basal spacing ranging from 7.5 Å to
51
7.3Å, and Phase II with basal spacing of ~ 6.6Å; the LDH structure,
itself, still remains intact.
(2) In the temperature range 190 ~ 280
o
C, the OH
-
group bonded with Al
3+
begins to disappear at 190
o
C, and is completely lost at 280
o
C. In this
temperature region, Phase I is transformed into Phase II.
(3) In the temperature range 280 ~ 405
o
C, the OH
-
group bonded with Mg
2+
begins to disappear at 280
o
C, and is completely lost at 405
o
C; a
degradation of the LDH structure is also observed in the same region.
(4) Finally, in the temperature range 405 - 580
o
C, CO
3
2-
loss begins and is
completed at 580
o
C. In this temperature range the material becomes a
metastable mixed solid oxide solution.
52
Chapter 2 References
1. Hibino, T.; Yamashita, Y .; Kosuge, K.; Tsunashima, A. Decarbonation
behavior of Mg-Al-CO
3
hydrotalcite-like compounds during heat treatment.
Clays and Clay Min., 1995, 43(4), 427.
2. Perez-Ramirez, J.; Mul, G.; Kapteijn, F.; Moulijn, J. A. (2000). Comments
on “Infrared emission spectroscopic study of the thermal transformation of
Mg-, Ni-, and Co-hydrotalcite catalysts” [Appl. Catal A:General, 1999, 18,
61]. Appl. Catal, A: General, 2000, 204, 256.
3. Kanezaki, E. Thermal behavior of the hydrotalcite-like layered structure of
Mg and Al-layered double hydroxides with interlayer carbonate by means
of in-situ powder HTXRD and DTA/TG. Sol. State Ion., 1998, 106(3-4), 279.
4. Cavini, F.; Trifiro, E.; Vaccari, A. Hydrotalcite-type anionic clays –
preparation, properties and applications. Cat. Today, 1991, 11, 173.
53
Chapter 3
A Study of the Sorption Properties of Mg-Al-CO3 LDH
3.1 Introduction
The use of a combination of several in-situ techniques to investigate the
thermal evolution of the structure of LDH1 under an inert atmosphere was reported
in the previous chapter, and a model was proposed to describe the structural
evolution of the Mg-Al-CO
3
LDH. In this chapter, the thermal evolution of the Mg-
Al-CO
3
LDH structure is studied further, and the focus in this study is on validating
the proposed structural model under more realistic reactive environments, in which
the LDH materials may find eventual applications. In addition, the sorption
characteristics and thermal reversibility of these materials under both inert and
reactive atmospheres are studied as well.
3.2 Experimental
The two LDH samples were provided by Media and Process Technology,
Inc., of Pittsburgh, PA. The composition of the one is Mg
0.71
Al
0.29
(OH)
2
(CO
3
)
0.15
.
54
0.46(H
2
O) (LDH1; the same LDH sample utilized in previous chapter), and the
other is Mg
0.645
Al
0.355
(OH)
2
(CO
3
)
0.178
.
0.105(H
2
O) (hereinafter referred to as LDH2)
determined by ICP and TGA. DRIFTS spectra were recorded in situ using a Genesis
II (Mattson, FT-IR) instrument equipped with a DRIFTS COLLECTOR
TM
II
chamber (SpectraTech, Inc.) capable of operating under high temperatures (up to
900
o
C) and pressures (up to 1500 psi). The experimental operating conditions were
a DRIFTS scan-range from 4000 cm
-1
to 500 cm
-1
, scan numbers 16, and a scan
resolution of 2 cm
-1
. The TGA curve was recorded on a Cahn TGA 121 instrument.
The MB-MS instrument is custom-made, using a MKS UTI 100C Precision Gas
Analyzer. SEM images were taken by a Cambridge Stereoscan 360, and TEM
images were taken by a Philips EM420 instrument. XRD data were generated by a
Rigaku X-ray diffractometer.
3.3 Results and Discussion
The weight losses for the two LDHs were studied by TG-MS, and the
results were shown to be consistent with the ICP results, as indicated in Table 3.1.
The LDH2 sample has a higher Al/Mg ratio (the value of x being near the higher
55
end of the typical LDH range), and contains less interlayer water. XRD
characterization of these two LDH samples indicates that the materials have the
typical LDH XRD spectra as shown in Fig. 3.1
1-3
, but the XRD spectrum of LDH2
is noisier than LDH1, and also not as strong as the spectrum of LDH1 (compare Fig.
3.1a with Fig 3.1b).
Table 3.1 Weight-loss from the TG/MB-MS studies, and calculated weight-loss
based on the ICP data for the samples. (a) LDH1 and (b) LDH2.
(a)
Weight-Loss H
2
O
OH
-
from Al
OH
-
from Mg
CO
2
from CO
3
2-
Total Weight-
Loss %
ICP Value 10.8% 10.4% 13.2% 7.2% 41.6%
Experimental
from TG/MS
12.93% 8.93% 12.65% 7.72% 42.23%
(b)
Weight-Loss H
2
O
OH
-
from Al
OH
-
from Mg
CO
2
from
CO
3
2-
Total Weight-
Loss %
ICP Value 2% 11.0% 16.1% 10.9% 40%
Experimental
from TG/MS
2% 12.08% 15.94% 11.06% 41.08%
(a)
10 20 30 40 50 60 70
0
1000
2000
3000
4000
5000
6000
7000
8000
9000
10000
(b)
10 20 30 40 50 60 70
0
500
1000
1500
2000
2500
3000
3500
4000
Figure 3.1 The XRD spectra of the LDH samples (a) LDH1 sample; (b) LDH2
sample.
56
(a)
(b)
Figure 3.2 (a) TEM picture of LDH1; (b) TEM picture of LDH2.
57
(a)
100 200 300 400 500 600 700
60
70
80
90
100
LDH 1
LDH 2
Temperature,
o
C
Weight %, %
MS Signal for CO
2
(m/e = 44)
58
0 100 200 300 400 500 600 700
0
20
40
60
80
0
(b)
10
Fraction of H2O Removed, %
Temperature,
o
C
LDH1
LDH2
Figure 3.3 The TGA spectra and CO
2
MS signal for the two LDH samples
generated with a scan rate of 5
o
C/min; (b) cumulative amount of H
2
O evolved.
59
The differences in the XRD spectra of these two LDH samples of similar
composition are accompanied by differences in the crystallite size of the materials
as shown in Fig. 3.2. The average crystallite size for LDH1 is 0.31 μm, while that
for LDH2 is 0.13 μm. The TGA spectra for the two LDHs generated with a scan
rate of 5
o
C/min, and the corresponding in-situ MS signals of CO
2
are shown in Fig.
3.3a. The cumulative amounts of H
2
O evolved are shown in Fig. 3.3b. Though
differences in the TGA and mass evolution spectra exist, the sequence of events
during the structural evolution of the two LDH remain the same, only the range of
temperatures where the different phenomena occur differs. Hibino et al. also
observed differences in the TGA spectra of Mg-Al-CO
3
-LDH with a different
Al:Mg ratios, but even between LDH with the same Al:Mg ratio and different
crystallite sizes
4
. For the LDH2, the loosely held interlayer water is lost in the
temperature range of 80 ~ 190
o
C, the OH
-
group begins to disappear at 190
o
C and
is completely lost around 520
o
C, and while some CO
3
2-
loss is observed at lower
temperatures, its substantial loss begins at 450
o
C, and is completed at 720
o
C. The
fractions in terms of the total of H
2
O and CO
2
evolved in different temperature
ranges for both LDH1 and LDH2 are shown in Table 3.2.
60
Table 3.2 The fractions of H
2
O and CO
2
(as % of the total sample weight) that are
evolved in different temperature ranges for both LDH1 and LDH2.
LDH1 LDH2
Weight-
Loss, %
H
2
O CO
2
Weight-
Loss, %
H
2
O CO
2
RT ~ 100
o
C 2.61 2.61 - 0.59 0.59 -
100 ~ 200
o
C 11.39 11.31 0.08 1.48 1.46 0.02
200 ~ 300
o
C 8.03 7.94 0.09 7.41 7.33 0.08
300 ~ 400
o
C 12.64 11.95 0.69 8.31 8.04 0.27
400 ~ 500
o
C 4.8 0.7 4.1 10.42 9.91 0.51
500 ~ 600
o
C 2.76 - 2.76 11.38 2.69 8.69
600 ~ 750
o
C - - - 1.49 - 1.49
Total 42.23 34.51 7.72 41.08 30.02 11.06
One of the issues of interest about the use of the LDH materials in the
preparation of CO
2
-permselective membranes and adsorbents pertains to their
ability to function stably under thermal cycling and other temperature changes in a
variety of gaseous atmospheres. To study the structural stability of these materials,
typically the temperature was raised in a linear fashion while monitoring weight-
loss and other structural characteristics in situ. For example, the temperature of
61
LDH1 sample was raised linearly under vacuum and in an inert Ar atmosphere,
while the sample was studied in situ by a variety of techniques, as previously
outlined.
For the heating rates utilized for the LDH1 sample, the observations made
up to 250
o
C were of equilibrium nature (no effect of the heating rate), but for some
of the heating rates utilized kinetic effects were apparent above this temperature.
Similar observations were previously made by Costantino and Pinnavaia
5
, and most
recently by Rhee and Kang
6
. To investigate the effect of heating rate further, the
weight-loss characteristics of the LDH2 sample have been studied for 4 different
heating rates, namely, 1, 3, 5 and 10
o
C/min, and in three different atmospheres. The
weight-loss results in an inert Ar atmosphere and the corresponding MS signals of
CO
2
are shown in Fig. 3.4a. For the results in Fig. 3.4, a fresh LDH2 sample with a
mass of about 110 ~ 120 mg was used in every experiment, and dry inert Ar (ultra
high purity grade) was utilized as a purge gas at a flow rate of 20 mL/min. For
heating rates below 5
o
C/min the weight-loss curves and the MS signals coincide,
indicating that the structural changes (loss of interlayer water, hydroxyl, and CO
3
2-
losses) occur rapidly enough, so that the LDH structure equilibrates within the time
(a)
100 200 300 400 500 600 700
60
70
80
90
100
1
o
C/min,
3
o
C/min,
& 5
o
C/min
Weight %, %
Temperature,
o
C
10
o
C/min
0.0
2.0x10
-7
4.0x10
-7
6.0x10
-7
8.0x10
-7
1.0x10
-6
1.2x10
-6
1.4x10
-6
1.6x10
-6
MS Signal
(b)
100 200 300 400 500 600 700
60
70
80
90
100
Ar
CO
2
Humidified CO
2
Weight %, %
Temperature,
o
C
Figure 3.4 The effect of varying (a) the heating rate; and (b) of using different
purging gases on the weight-loss for the LDH2 sample.
62
63
frame allotted by the changing temperature. However, the results with the 10
o
C/min heating rate contain kinetic artifacts. Similar observations were also made in
the presence of dry and humid CO
2
atmospheres, with the results showing absence
of kinetic effects for heating rates below 5
o
C/min.
The effect of varying the heating rate on weight-loss for the LDH2 sample
was also studied in the presence of a reactive atmosphere. For the experiments in
Fig. 3.4b, in addition to the weight-loss of the LDH2 sample in the presence of inert
Ar, we also show the weight-loss curve for the case in which dry CO
2
, instead of Ar,
was utilized as a purge gas atmosphere at a flow rate of 20 mL/min and a heating
rate of 5
o
C/min. The weight-loss results in the presence of a humidified CO
2
atmosphere are also shown in the same figure. For these experiments, the other
experimental conditions were the same as those with the other two weight-loss
curves shown in the same figure, with the exception that the CO
2
stream in this case
was humidified by bubbling it through a beaker containing distilled water.
Measurements of the water concentration of the gas exiting the beaker indicate that
the relative humidity (RH) of the CO
2
stream was ~ 70%. The results in Fig. 3.3b
indicate little effect of the gaseous atmosphere on the weight-loss curve in the first
64
region of temperatures associated with the evolution of interlayer water.
Differences exist, however, in the other regions. In the region where mostly CO
2
evolves, and DRIFTS indicates that all hydroxyls in the LDH structure have already
disappeared, the weight-loss curves for the humidified and dry CO
2
atmospheres
coincide, but are still different from the weight-loss curve under inert conditions;
the presence of CO
2
in the purge atmosphere appears to slow-down somewhat the
rate of CO
2
evolution. The dry and humidified CO
2
weight-loss curves are different
in the region associated with hydroxyl evolution, particularly in the range of
temperatures associated with loss of hydroxyls in a Mg-(OH)-Mg configuration.
Previously, Ding and Alpay
7
, who studied CO
2
adsorption on a K-promoted
commercial Mg-Al hydrotalcite at 400
o
C, noted a small (~ 10%) beneficial effect
of the presence of water on CO
2
adsorption. They also noted, however, that the
actual partial water pressure did not really matter, with even traces of water vapor
being capable of providing the same beneficial effect. Ding and Alpay
7
attribute this
beneficial effect to the ability of water vapor to either maintain the hydroxyl
concentration on the surface or to prevent the sites from poisoning through
carbonate or coke deposition.
65
The ability of the LDH to reversibly adsorb CO
2
and H
2
O is of significance
in the use of these materials as adsorbents and membranes. In the former case, the
ability to reversibly adsorb CO
2
is critical from the standpoint of being able to
regenerate the adsorbent; in the latter case the presence of a relatively mobile CO
2
phase within the LDH structure is important in determining the permeation rate
through the membrane. We investigated, therefore, the ability of the LDH materials
to reversibly adsorb CO
2
under a broad range of experimental conditions. To initiate
these studies, firstly the effect of utilized condition was investigated during the
experiments. For each series of experiments, 100 ~ 120 mg of a fresh LDH sample
was utilized. During the adsorption part of the cycle, 30 mL/min of CO
2
was
bubbled through a beaker containing distilled water (the CO
2
stream’s RH being ~
70%) and the sample was exposed to this humidified CO
2
stream for varying
periods of time. Subsequently to adsorption, the flow of CO
2
was shut down and the
desorption part of the cycle was initiated. To find optimal experimental conditions
for sorption studies, the effect of varying the duration of adsorption was
investigated. And it was observed that increasing the time of adsorption from 1 to 2
hr increased the total amount adsorbed by about 5%, but a subsequent increase from
66
2 hr to 3 hr had no additional significant effect. Therefore, for the remainder of the
study, an adsorption step time of 3 hr was utilized. Also two different methods to
carry out the desorption step was investigated. In the first method, upon termination
of the CO
2
flow, the sample was exposed to flowing UHP dry Ar at a rate of 30
mL/min. Typically, after 30 min the weight change of the sample ceased.
Subsequently, Ar was flowed to continue for a total desorption period of 1 hr. In the
second method, the chamber was evacuated for a period of 1 hr at a pressure below
40 mTorr. Evacuation was shown to be a more effective means for carrying out the
desorption step (~ a 10% increase in weight gain upon subsequent readsorption),
and was utilized in the remainder of the study.
Upon completion of the preliminary runs, the effect of temperature was
studied on the adsorption/desorption behavior of the LDH2. The results are shown
in Fig. 3.5. Fig 3.5a shows the total weight gain (as percent fraction of the original
weight of the LDH sample) during the adsorption part of the cycle, Fig. 3.5b shows
the corresponding total weight-loss during the desorption part of the cycle, and Fig.
3.5c shows the weight change corresponding to either the H
2
O or CO
2
lost during
the desorption part in the cycle. For each experiment, at any given temperature, a
(a) (b)
012 3456
1.00
1.20
1.40
1.60
1.80
2.00
150
o
C
170
o
C
180
o
C
190
o
C
200
o
C
250
o
C
300
o
C
350
o
C
Weight Loss, %
Cycle Numbers
Desorption
67
(c)
Figure 3.5 Weight-gain or loss. (a) weight-gain during adsorption for various
temperatures as a function of the cycle number; (b) weight-loss during desorption
for various temperatures as a function of the cycle number; (c) weight-change due
to loss of water or CO
2
during desorption as a function of temperature.
01 2 34 56
1.00
1.20
150
o
C
170
o
C
1.40
1.60
1.80
2.00
180
o
C
190
o
C
200
o
C
250
o
C
300
o
C
350
o
C
W Gain, %
Cycle Numbers
Adsorption
eight
150 200 250 300 350
0.0
0.5
1.0
1.5
2.0
2.5
3.0
1st cycle
2nd cycle
3rd cycle
4th cycle
5th cycle
Temperature,
o
C
H
2
0 Weight %
0.0
0.5
1.0
1.5
2.0
2.5
3.0
CO
2
Weight %
68
fresh LDH2 sample (100 ~ 120 mg) was used. The temperature was increased from
room temperature at a rate of 5
o
C/min in an Ar atmosphere (flow rate of 30
mL/min) to the preset point (e.g., 200
o
C, 250
o
C, etc.), and kept at this temperature
until the sample weight became constant, typically for 20~30 min. Subsequently,
the cyclic adsorption/desorption experiments were initiated. Three distinct regions
of different behavior can be distinguished in Fig. 3.5. The first region is for
temperatures below 190
o
C, in which during the cyclic adsorption/desorption
studies it was observed that the LDH2 sample reversibly adsorbed only water and
slight amounts of CO
2
. The results of the cyclic adsorption/desorption experiments
in this region of temperatures are consistent with the observations of the previous
chapter with the LDH1 sample and current studies with the LDH2, which indicate
that, in the same region of temperatures, the interlayer water is removed. The cyclic
adsorption/desorption experiments also indicate that the exchange of interlayer
water is a fairly reversible process.
In the second temperature range, from 190
o
C to 280
o
C, the studies in
previous chapter with LDH1 and current studies with LDH2 indicate that the water
that leaves the sample is from the hydroxyl groups that are bonded with Al cations.
69
In addition, some CO
2
is also emitted in this region. The cyclic
adsorption/desorption studies show that the same two species are also emitted
during the period in which the sample’s temperature is raised to the desired level.
Upon initiation of the adsorption/desorption cycle, however, only CO
2
appears to be
reversibly adsorbed in this region, with very little H
2
O emitted; the sample weight
change can be fully attributed to the reversibly adsorbed CO
2
.
In the temperature range 280 ~ 440
o
C, the previous studies, under inert
conditions, have indicated that the OH
-
group bonded with Mg
2+
begins to disappear
at 280
o
C, and is completely lost at 440
o
C (for the LDH2 the upper temperature
extends higher; see Table 3.2); a degradation of the hydrotalcite structure is also
observed in the same region. The cyclic adsorption/desorption experiments indicate
that the same two species are also emitted during the period in which the sample
temperature is raised to the desired level. Upon initiation of the
adsorption/desorption cycle, however, only CO
2
appears to be reversibly adsorbed
in this region, with very little H
2
O emitted; hence, the sample weight change, once
more, can be mostly attributed to the reversibly adsorbed CO
2
. As can be seen in
Fig 3.5c, the amount of CO
2
that is reversibly adsorbed in this region decreases as
70
the temperature increases, consistent with the observations that the crystallinity of
the hydrotalcite material also decreases, and its structure begins to fall apart in this
region.
To further validate the cyclic adsorption/desorption behavior observed
using the TG/MB-MS, in situ cyclic flow adsorption/desorption experiments were
carried out using the DRIFTS system, following the same experimental protocol as
with TG/MB-MS experiments described above. During these studies a number of
distinct peaks corresponding to various functional groups in the hydrotalcite were
monitored, and the various DRIFTS peaks were assigned similarly as in the
previous chapter. Fig. 3.6b shows the integrated peak area reflecting the CO
3
2-
ν
3
vibration during the cyclic adsorption/desorption experiments, which is expressed
as percent fraction of the peak area at the beginning of the experiment at room
temperature. Fig. 3.6c shows the 3470 cm
-1
band corresponding to the OH
-
vibration,
and the combined integrated peak areas for the interlayer water peaks (3070 cm
-1
and 1620 cm
-1
), again expressed as percent fraction of the same peak areas at the
beginning of the experiment at room temperature. In these figures the experimental
data at the end of the first and second adsorption and desorption cycles for various
4500 4000 3500 3000 2500 2000 1500 1000 500 0
680cm
-1
940cm
-1
1371cm
-1
1620cm
-1
3070cm
-1
1350cm
-1
1370cm
-1
1530cm
-1
Wavenumber,cm
-1
Kubeika Munk Units
3470cm
-1
750
o
C
700
o
C
650
o
C
600
o
C
550
o
C
500
o
C
450
o
C
400
o
C
350
o
C
300
o
C
250
o
C
200
o
C
150
o
C
100
o
C
50
o
C
RT
(a)
(b)
100 150 200 250 300 350 400
10
20
30
40
50
60
70
80
90
100
OH
-
1st Adsorption
OH
-
1st Desorption
OH
-
2nd Adsorption
OH
-
2nd Desorption
H
2
O 1st Adsorption
H
2
O 1st Desorption
H
2
O 2nd Adsorption
H
2
O 2nd Desorption
% of area
Temperature,
o
C
71
100 150 200 250 300 350 400
95
96
97
98
99
100
CO
3
2-
1st Adsorption
CO
3
2-
1st Desorption
CO
3
2-
2nd Adsorption
CO
3
2-
2nd Desorption
% of area
Temperature,
o
C
Figure 3.6 (a) In-situ DRIFTS of LDH2 as a function of temperature; (b) change in
the CO
3
2-
integrated peak area (left), and change in the OH
-
and H
2
O integrated
peak areas as a per cent fraction of the original peak area (right) during the
sorption/desorption cycles.
72
temperatures are shown. In Fig. 3.6b and 3.6c it is observed that at 150
o
C only the
combined integrated peak area for the interlayer water peak changed in a reversible
manner. No substantial changes are observed during the adsorption/desorption
cycles in the integrated peak areas reflecting the CO
3
2-
ν
3
vibration or the 3470 cm
-1
band corresponding to the OH
-
vibration. The hydrotalcite during the cyclic
adsorption/desorption experiments simply exchanges reversibly only interlayer
water. Above the temperature of 190
o
C, the interlayer water disappears during the
heating period in Ar before reaching the preset temperature. During the cyclic
adsorption/desorption process the hydrotalcite exchanges reversibly only CO
2
to
any substantial extent. Furthermore, the amount of CO
2
that reversibly adsorbs and
desorbs decreases as a function of temperature. Clearly, the hydrotalcite under the
conditions of the cyclic experiments described here is not capable of reversibly
exchanging the OH
-
, after a certain amount of OH
-
is emitted during the heating
period to reach the preset temperature. These observations are consistent with the
TG-MS experiments and the observations of the previous chapter.
Longer-term cyclic adsorption/desorption experiments were also carried out.
For the experiments with the LDH1, 10 mg of sample was utilized, and two
73
different temperatures 150
o
C, and 250
o
C were investigated (the same sample was
used for both experiments). For each experiment, the sample was first heated in
UHP dry Ar (20 mL/min) with the heating rate of 5
o
C/min, and the weight of the
sample and the gas composition were monitored. Upon reaching the desired
temperature, the feed was switched to humidified carbon dioxide (20 mL/min, 70%
RH), and kept there for 3 hr. Then the sample was evacuated for 1 hr, switched back
to humidified carbon dioxide for 3 hr, and so on. Fig. 3.7 shows the weight change
and the total weight (both as percent fractions of the original sample weight)
observed for a total of 14 cycles. Fig. 3.8 shows the corresponding MS signals
during the heating and evacuation parts of the cycle. Only water was detected
coming out of the sample at 150
o
C, and the observation is consistent with the
structural model for the LDH presented in the previous chapter. The experiments at
150
o
C indicate that the system reaches a steady-state, reversible behavior after the
11
th
cycle, with the corresponding weight change being 0.23%.
Upon completion of the 14 cycles at 150
o
C, the LDH1 sample was heated
in UHP dry Ar (20 mL/min; 5
o
C/min) until its temperature reached 250
o
C; the feed
was then switched to humidified CO
2
(20 mL/min, 70% RH), and kept there
02 46 8 10 12 14
0.0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
Adsorption
Desorption
Weight %, %
Number of cycles
LDH 1 @150
o
C
02 46 8 10 12 14
94.4
94.6
94.8
95.0
95.2
95.4
95.6
95.8
96.0
96.2
96.4
Overall Weight %, %
Number of Cycles
Adsorption
Desorption
LDH1 @150
o
C
Figure 3.7 Weight-gain or loss (top) and total sample weight (bottom) during the
sorption/desorption cycles.
74
0 500 1000 1500 2000 2500 3000
0.10
0.12
0.14
0.16
0.18
0.20
0.22
0.24
0.26
0.28
0.30
0.32
LDH1 @150
o
C
Time, min
MS Signal of H
2
O
Heating up
1st Cycle Desorption
-0.002
-0.001
0.000
0.001
0.002
0.003
0.004
0.005
0.006
0.007
0.008
0.009
0.010
MS Signal of CO
2
Figure 3.8 H
2
O and CO
2
MS signals during the heating, and the desorption part of
the cycles as a function of time.
75
0246 8 10 12 14
0.0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
2.0
2.2 LDH1 @250
o
C
Adsorption
Desorption
Weight %, %
Number of Cycles
02 46 8 10 12 14
87.6
87.8
88.0
88.2
88.4
88.6
88.8
89.0
89.2
89.4
89.6
89.8 LDH1 @250
o
C
Overall weight %, %
Number of Cycles
Adsorption
Desorption
Figure 3.9 Weight-gain or loss (top) and total sample weight (bottom) during the
sorption/desorption cycles.
76
0 500 1000 1500 2000 2500 3000 3500
0.20
0.25
0.30
0.35
0.40
0.45
0.50
0.55
0.60
Time, min
LDH1 @250
o
C
MS Signal of H
2
O
1st Cycle Desorption
Heating up
-0.14
-0.12
-0.10
-0.08
-0.06
-0.04
-0.02
0.00
0.02
0.04
0.06
MS Signal of CO
2
Figure 3.10 H
2
O and CO
2
MS signals during the heating, and the desorption parts
of the cycle as a function of time.
77
78
for 3 hr. Then, the sample was evacuated for 1 hr, switched back on to humidified
CO
2
for 3 hr, and so on. Fig. 3.9 shows the weight change observed for a total of 14
cycles for the LDH1 sample at 250
o
C. The sample reaches a steady state behavior
after the 9
th
cycle, with the corresponding reversible weight change being 0.31%.
Fig. 3.10 shows the corresponding MS signals during the heating and evacuation
parts of the cycle. During the heating part of the experiment, both water and CO
2
are emitted; however, subsequently to that only CO
2
is emitted during the cyclic
experiments, indicating that under these relatively low RH conditions the OH
-
are
not reversibly exchanged to any substantial extent. As noted previously, a similar
observation was also made for the LDH2 sample during the shorter term
reversibility experiments. The observations are, once again, consistent with the
structural model for the LDH proposed in the previous chapter. Fig. 3.11 describes a
long-term cycling experiment using 113 mg of LDH2 at 250
o
C, following
otherwise the same experimental conditions as with the experiments involving the
LDH1 sample. The sample reaches a steady state behavior after the 14
th
cycle, with
the corresponding reversible weight change being 0.32%.
0246 8 10 12 14 16
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
2.0
2.2
Weight %, %
Number of Cycles
Adsorption
Desorption
LDH 2 @ 250
o
C
02468 10 12 14 16
96.0
96.2
96.4
96.6
96.8
97.0
97.2
97.4
97.6
97.8
98.0
98.2
98.4 LDH 2 @ 250
o
C
Adsorption
Desorption
Overall weight %, %
Number of Cycle
Figure 3.11 Weight-gain or loss (top) and total sample weight (bottom) during the
sorption/desorption cycles.
79
0 20 40 60 80 100 120 140 160 180 200 220
86
88
90
92
94
96
98
100
20
40
60
80
100
120
140
160
180
200
220
Temperature,
o
C
150
o
C
200
o
C
Weight %, %
Time, min
LDH 1
Figure 3.12 Weight-loss/gain during the temperature cycling experiments. Solid
lines are the experiments from room temperature to 150
o
C; Dotted lines are
experiments from room temperature to 200
o
C.
80
81
After the cyclic adsorption/desorption experiments with the LDH samples
were completed, a number of experiments were initiated in which the weight-
gain/loss of the sample was monitored as its temperature was cycled from room
temperature to a preset temperature, and back down to room temperature. The same
LDH1 sample was used in all of temperature cycling experiments reported in this
section.
In the first experiment, the sample was heated in flowing Ar (20 mL/min)
with a heating rate of 3
o
C/min from room temperature to a temperature of 150
o
C;
subsequently the flowing Ar feed was substituted with a humidified CO
2
feed (20
mL/min) and was cooled down to room temperature with a 3
o
C/min cooling rate.
Upon reaching room temperature, the sample was kept at this temperature for an
additional 2 hr. The weight-loss/gain data are shown in Fig. 3.12. The total weight-
loss for the LDH1 sample was ~5.5% at 150
o
C; upon cooling in the humidified
CO
2
atmosphere, the sample recovered 98.4% of its original weight. The gaseous
components evolved during the heating step were monitored by MS (Fig. 3.13);
only water was detected during the experiment.
82
Upon termination of the experiment at 150
o
C, the humidified CO
2
atmosphere was switched back to flowing UHP dry argon, and the temperature was
slowly (3
o
C/min) increased to 200
o
C. Upon reaching this temperature, the same
experimental protocol was followed. The weight loss curve leveled off after 240
min. The sample, upon cooling, recovered 99.8% of its original weight (Fig. 3.12).
The gases evolved during the heating part of the cycle were also monitored (Fig.
3.14). Water was detected throughout the whole temperature range, and trace
amounts of carbon dioxide were detected from 195
o
C ~ 200
o
C. Most of the weight
loss was observed below 150
o
C; only ~0.2 % of the weight-loss was observed
between 180
o
C and 200
o
C.
Upon completion of the experiment at 200
o
C, the humidified CO
2
atmosphere was switched back to dry argon, and the temperature was slowly (1
o
C/min) increased to 250
o
C (in the 250
o
C and higher temperature cycling
experiments, the cooling/heating rates were changed from 3
o
C/min to 1
o
C/min).
Upon reaching this temperature, the flowing Ar feed was substituted with a
humidified CO
2
feed (20 mL/min), and the sample was cooled down to room
temperature with a 1
o
C/min cooling rate. Also the sample was allowed to
0 5 10 15 20 25 30 35 40 45
20
40
60
80
100
120
140
160
Time, min
Temperature,
o
C
0.22
0.24
0.26
0.28
0.30
0.32
0.34
0.36
MS signal of H
2
O
0 10 203040
20
40
60
80
100
120
140
160
Time, min
Temperature,
o
C
-0.10
-0.05
0.00
0.05
0.10
MS Signal of CO
2
Figure 3.13 MS signals for H
2
O (top) and CO
2
(bottom) during the temperature
cycling experiment from room temperature to 150
o
C.
83
0 1020 304050 60
0
20
40
60
80
100
120
140
160
180
200
220
Time, min
Temperature,
o
C
1.22
1.23
1.24
1.25
1.26
1.27
1.28
1.29
1.30
1.31
1.32
1.33
MS signal of H
2
O
0 102030 405060
0
20
40
60
80
100
120
140
160
180
200
220
Time, min
Temperature,
o
C
-0.02
-0.01
0.00
0.01
0.02
0.03
0.04
0.05
MS Signal of CO
2
Figure 3.14 MS signals for H
2
O (top) and CO
2
(bottom) during the temperature
cycling experiment from room temperature to 200
o
C.
84
85
equilibrate at room temperature for as long as necessary for the weight-gain curve
to level off. Upon cooling, 99.2% of original weight of the sample was recovered;
most of the weight-loss was again observed below 150
o
C (Fig. 3.15). The
components evolved during the heating part of the cycle were also monitored (Fig.
3.16). Water was detected throughout the whole temperature range, similar to the
previous experiments, and smaller amounts of carbon dioxide were detected in the
range from 195
o
C ~ 250
o
C. The experiment was repeated with the temperature
raised (1
o
C/min) to 300
o
C, cooled down in humidified CO
2
to room temperature
with a 1
o
C/min cooling rate, and left there for as long as necessary for the weight
gain curve to level off. As shown in Fig. 3.15, it took a much longer time than in the
previous experiments for the sample to regain its weight, which leveled off at 96%
of the original weight value. The composition of outlet gas showed water being
evolved through the whole range of temperatures; carbon dioxide was again
detected between 195
o
C and 300
o
C. The experimental results with the temperature
raised to 350
o
C are also shown in Fig. 3.15. Once more, it took a much longer time
for the sample to recover its weight, which leveled off at 93.8% of its original
weight; water was evolved through the whole region of temperatures and carbon
(a) (b)
Temperature,
o
C
0 200 400 600 800 1000 1200 1400
80
82
84
86
88
90
92
94
96
98
Time, min
Weight %, %
50
100
150
200
250
300
Temperature,
o
C
0 100 200 300 400 500 600 700 800
86
88
92
96
0
98
10
0
50
100
150
200
250
86
(c)
Figure 3.15 Weight-loss/gain during the temperature cycling experiments (a) from
room temperature to 250
o
C; (b) from room temperature to 300
o
C; (c) from room
temperature to 350
o
C.
84
90
94
Time, min
Weight %, %
0 400 800 1200 1600 2000
78
80
82
84
86
88
90
92
94
Time, min
Weight %, %
40
80
120
160
200
240
280
320
360
Temperature,
o
C
0 1020 3040506070
0
20
40
60
80
100
120
140
160
180
200
220
240
260
Time, min
Temperature,
o
C
1.59
1.60
1.61
1.62
1.63
1.64
1.65
1.66
1.67
1.68
MS Signal of H
2
O
0 1020 3040506070
0
20
40
60
80
100
120
140
160
180
200
220
240
260
Time, min
Temperature,
o
C
-0.005
0.000
0.005
0.010
0.015
0.020
0.025
0.030
0.035
0.040
MS Signal of CO
2
Figure 3.16 MS signals for H
2
O (top) and CO
2
(bottom) during the temperature
cycling experiment from room temperature to 250
o
C.
87
88
dioxide was detected between 195
o
C and 350
o
C. These observations are clearly
consistent with the structural model for the LDH developed by us in the previous
chapter, based on observations made under inert or vacuum conditions.
To further study the sorption reversibility behavior of the LDH, and in
order to establish a connection between the TG/MB-MS and DRIFTS results,
experiments were carried out at atmospheric pressure conditions and under
conditions one may encounter in the use of these materials in the water gas shift
(WGS) membrane reactor environment, similar experiments were carried out using
a moderate pressure adsorption flow system. The experimental system is equipped
with mass flow controllers and a flow control valve at the exit to maintain the
system pressure constant under flow conditions.
Two types of experiments were performed. In the first series of experiments
the flow system was first pressurized with flowing dry argon (50 mL/min) to 50
psig , then the temperature was increased to 250 °C, using a 5 °C /min heating rate
(it takes ~ 45 min), and the system was kept at 250 °C for 1 hr as a desorption step.
When the desorption step was over, the system was cooled down to 150°C (cooling
rate 5
o
C/min, ~ 20 min) in flowing dry argon (50 mL/min). Subsequently, the inlet
89
gas was changed to dry CO
2
(50 mL/min) from argon, while keeping the same
pressure of 50 psig for 3 hr as an adsorption step. During the adsorption step, the
outlet flow rate was monitored by a digital flow meter (while the reactor pressure
was maintained constant at 50 psig), and, from the flow rate change, the amount of
adsorption was calculated. Table 3.3 shows the weight gain during the sorption step
for the first 3 sorption/desorption cycles with the LDH2 sample.
Table 3.3 Weight gain during the sorption step for the moderate pressure flow
experiments using dry CO
2
.
Weight-gain during the adsorption step (wt. %)
1
st
Cycle 2.621
2
nd
Cycle 2.542
3
rd
Cycle 2.476
In the second series of experiments the LDH2 sample (~ 14 g) was first heated to a
preset temperature in Ar (50 mL/min). Each cycle involved first evacuating the
sample for 1 hr as a desorption step. After the evacuation step, the flow system was
again pressurized to 50 psig in flowing Ar. When the outlet flow rate was stabilized
90
in flowing Ar at 50 psig, the inlet gas was then changed to either dry or humidified
CO
2
(50 mL/min)
for 3 hr while maintaining the same pressure of 50 psig.
Table 3.4 Weight gain during the adsorption step for the moderate pressure flow
experiments at various temperatures using dry CO
2
.
150
o
C 200
o
C 250
o
C
1
st
Cycle 2.02 1.84 1.72
2
nd
Cycle 1.97 1.79 1.69
3
rd
Cycle 1.87 1.71 1.68
4
th
Cycle 1.84 1.69 1.66
Upon completing the first adsorption/desorption cycle the procedure was
repeated for a number of additional cycles, and for a number of temperatures. The
cyclic adsorption/desorption experimental results for the various temperatures for
dry CO
2
are summarized in Table 3.4, while those for humidified CO
2
for one
temperature are shown in Table 3.5 (a fresh LDH2 sample was utilized for each set
experiments at every new temperature). During the humidified CO
2
experiments at
moderate pressures, a syringe pump was used (instead of a bubbler) to deliver a
predetermined amount of water into the CO
2
stream. For the experiments in Table
3.5, the CO
2
stream contains 2% mol of water, which corresponds approximately to
91
Table 3.5 Weight gain during the sorption step for the moderate pressure flow
experiments at various temperatures using humidified CO
2
.
200
o
C Weight gain at adsorption step (wt. %)
1
st
Cycle 1.85
2
nd
Cycle 1.79
3
rd
Cycle 1.73
4
th
Cycle 1.71
70% RH at the temperature and pressure of the experiment. For the run at 150
o
C
the weight lost during the desorption part corresponds to water with only traces of
CO
2
being lost. The opposite observation is true for the runs at 200
o
C and 250
o
C.
Comparing the run in humidified CO
2
(Table 3.5) with the corresponding run with
dry CO
2
, it is observed that there is little difference in the weight change, which is
consistent with the observations under atmospheric conditions that beyond the
initial heating step, the LDH does not exchange water reversibly under these
temperature conditions. The whole set of observations under moderate pressure
conditions are again in agreement with the TG/MB-MS and DRIFTS data under
atmospheric conditions, and the LDH structural model. This lends confidence in the
92
usefulness of these fundamental techniques in predicting the materials behavior
under realistic process conditions.
3.4 Conclusions
A test of the validity of the LDH thermal evolution model, proposed in
chapter 2 was provided, and the whole set of observations under atmospheric and
moderate pressure conditions are in agreement with the LDH structural model in the
previous chapter. The LDHs are shown capable of exchanging reversibly CO
2
for a
broad region of conditions. In addition to providing a test of the validity of the LDH
thermal evolution model, these experimental observations are of importance in their
own right for the use of the LDH materials in the preparation of CO
2
-permselective
membranes and adsorbents for high temperature membrane reactor applications.
One particular promising application is their use in reactive separations with the
water gas shift reaction. The ability of the LDH to reversibly adsorb CO
2
is of
significance in the use of these materials in such an application as adsorbents and
membranes. The ability to reversibly adsorb CO
2
is critical from the standpoint of
being able to regenerate the adsorbents, and the presence of a relatively mobile CO
2
93
phase within the LDH structure is important in determining the permeation rate
through the membrane.
94
Chapter 3 References
1. Kanezaki, E. Effect of atomic ratio Mg/Al in layers of Mg and Al layered
double hydroxide on thermal stability of hydrotalcite-like layered structure
by means of in-situ high temperature powder X-ray diffraction. Mater. Res.
Bull., 1998, 33(5), 773.
2. Kanezaki, E. Thermal behavior of the hydrotalcite-like layered structure of
Mg and Al-layered double hydroxides with interlayer carbonate by means
of in-situ powder HTXRD and DTA/TG. Solid State Ionics, 1998, 106(3-4),
279.
3. Cavini, F.; Trifiro, E.; Vaccari, A. Hydrotalcite-type anionic clays –
preparation, properties and applications. Catal. Today, 1991, 11, 173.
4. Hibino, T.; Yamashita, Y.; Kosuge, K.; Tsunashima, A. Decarbonation
behavior of Mg-Al-CO
3
hydrotalcite-like compounds during heat treatment.
Clays and Clay Minerals, 1995, 43(4), 427.
5. Constantino, V. R. L.; Pinnavaia, T. J. Basic properties of Mg
1-
X
(2+)Al
X
(3+) layered double hydroxides intercalated by carbonate,
hydroxide, chloride, and sulfate anions, Inorg. Chem., 1995, 34, 883.
6. Rhee, S. W.; Kang M. J. Kinetics on dehydration reaction during thermal
treatment of Al-CO3-LDHs, Kor. J. Chem. Eng., 2002, 19, 653.
7. Ding, Y.; Alpay, E. Equilibria and kinetics of CO
2
adsorption on
hydrotalcite adsorbent. Chem. Eng. Sci., 2000, 55, 3461.
95
Chapter 4
Measurements of Diffusivity and of the Adsorption
Isotherm of Carbon Dioxide
in Mg-Al-CO
3
LDH at Elevated Temperatures
4.1 Introduction
The transport characteristics and adsorption isotherms of carbon dioxide on
Mg-Al-CO
3
LDH were investigated using gravimetrically measured CO
2
uptake
data in the temperature range of 200 ~ 250
o
C. Since LDHs are intended for the
preparation of nanoporous membranes, which are utilized for the separation of CO
2
at high temperatures, it is desirable to carry out experiments at elevated
temperatures rather than ambient or sub-ambient temperatures.
The gravimetric method for uptake rate measurements is simple and straightforward,
and several research groups have reported the value of CO
2
diffusivity data
measured by this method at ambient temperature, as shown in Table 4.1. In the
experiments reported here, the transient CO
2
uptake data were measured
gravimetrically at each elevated temperature, and then the diffusion coefficient was
96
Table 4.1 Diffusivity data (D/r
2
) for CO
2
measured by the gravimetric method.
Adsorbent D/r
2
, s
-1
Temperature, K References
Coconut derived CMS 7 x 10
-2
273 6
Bergbau Forschung
CMS
5 x 10
-6
298 7
Zr-pillared clay 8 x 10
-2
298 8
Own zeolite crystals
(7.3-34 μm)
4.7 x 10
-5
500 9
LDH2 (r ≈ 100 μm) 3.6 x 10
-4
473 This work
LDH2 (r ≈ 100 μm) 8.9 x 10
-4
498 This work
LDH2 (r ≈ 100 μm) 1.3 x 10
-3
523 This work
estimated by fitting the acquired experimental data to the solution of the relevant
diffusion equation. The diffusivity observations are compared with the result of
molecular dynamic simulation, which were implemented as a parallel research
project. Adsorption isotherm data were also acquired by the gravimetric method by
uptake measurements, and the experimental data were fitted with the Langmuir
equation and various empirical adsorption isotherm equations.
97
For another approach for the measurement of the diffusivity constant of
CO
2
in LDH materials, solid state pulse field gradient (PFG) nuclear magnetic
resonance spectroscopy (NMR) experiments were conducted. In the literature, the
PFG NMR method has been shown to provide perhaps the most reliable means of
measuring self-diffusivities in adsorbent-adsorbate systems, since it is remarkably
sensitive to molecular displacements in the range of 10 nm ~ 100 μm without any
chemical interferences
1
. Also, PFG NMR provides a direct determination of the
mean square displacement in a given time interval, which can be regarded as a
direct measure of the self-diffusivity.
While, in a gravimetric method, the adsorbate is introduced into the adsorbent
during the experimental measurement, PFG NMR measurements are generally
implemented with a system in which the adsorbent has been equilibrated with the
adsorbate under known conditions prior to the experiment. Therefore, when NMR
measurements are made over a range of temperatures it is also important to
recognize that the distribution of molecules between adsorbed and vapor phases
may vary with temperature. However, this effect is not significant when the
98
equilibrium is sufficiently favorable so that almost all of the molecules are in the
adsorbed phase.
For the study of the self-diffusivity constant with PFG NMR experiments, three
LDH2 samples were prepared. The first sample was evacuated at room temperature
for 1 hr; the second sample was evacuated at 220
o
C for 1 hr; the third one was
evacuated at 220
o
C for 1 hr, followed by exposing it to
13
C enriched CO
2
. The
results from the first sample would provide the information about the mobility of
carbonate molecules in the presence of interlayer water; the results from the second
sample would provide the information in the absence of interlayer water, and the
results from the third sample would provide information about the diffusivity of
CO
2
molecules. Also, the results from the third sample may be able to provide proof
whether CO
2
molecules are transformed into the carbonate groups, and whether the
transport controlling mechanism is intra-crystalline or inter-crystalline diffusion.
During the course of diffusivity and adsorption isotherm experiments, it
was observed that there was a difference in the equilibrium uptake amounts for CO
2
in various samples of the same Mg-Al-CO
3
LDH but whose particle size and
particle size distribution varied. When the as prepared LDH2 was utilized for CO
2
99
uptake experiments without separating the LDH2 particles by size, the overall
equilibrium adsorbed amount was observed to be ~2 wt%. However, when sieved
LDH2 samples were utilized, whose average particle radius was 100 μm, the overall
uptake of CO
2
was observed to be only ~1.1 wt%. Since the only difference in the
experimental conditions between the above two LDH2 samples was the different
average particle size, it is obvious that the particle size of LDH2 affects the
transport and sorption properties. So far, the particle size effect on the adsorption
capacity for CO
2
has not been reported in the published LDH literature, but several
research groups have reported similar particle size effects in their adsorption
experiments with other adsorbents.
For example, Badruzzaman et al.
2
observed that there was a nonlinear relationship
between surface diffusivity and particle radius for arsenate adsorption onto granular
ferric hydroxide. On the other hand, Grande et al.
3
in their studies with their own
synthesized 4A zeolites reported that the adsorption capacity of propane and
propylene was independent of the particle size.
To investigate the particle size effect on CO
2
uptake in LDH, a Mg-Al-CO
3
LDH was synthesized and then sieved into six distinct samples, each with a
100
different average particle radius. For each LDH sample, the diffusivity and the
adsorption isotherm for CO
2
were investigated with gravimetric methods at 200
o
C.
The surface area and pore volume were measured by BET analysis.
4.2 Experimental
In addition to the Mg-Al-CO
3
LDH samples prepared by us, we also
measured the diffusivity and adsorption isotherm for a sample provided by Media
and Process Technology, Inc., of Pittsburgh, PA, with the composition
Mg
0.645
Al
0.355
(OH)
2
(CO
3
)
0.178
.
0.105(H
2
O), previously referred in Chapter 3 as the
LDH2 sample. Prior to the diffusivity and adsorption isotherm experiments the
LDH2 sample was sieved, and only particles with radii in the range 90 ~ 105 μm
were retained. For the CO
2
sorption measurements, 100 ~ 120 mg of a freshly
sieved LDH2 sample was utilized in order to measure the diffusivity and the
adsorption isotherm. For the experiments CO
2
/Ar mixtures were prepared using dry
UHP CO
2
and Ar using Brooks 5850E mass flow controllers.
Sorption data were recorded using a Cahn TGA 121 instrument. Before
initiating the experiment, the LDH sample was spread as thin layer on a bowl-
101
shaped quartz container in the TGA apparatus in order to minimize external mass
transfer resistance. Then the sample was heated to a preset temperature in UHP dry
Ar (at a flow rate of 30 mL/min) with a heating rate of 5
o
C/min. Ar was utilized as
the purge gas in order to minimize the buoyancy force effect when the purge gas
was switched to dry CO
2
. The temperature of experiments was selected between
200
o
C and 250
o
C since the interlayer water of LDH can be removed without
significant transformation of LDH structure in this range as outlined in chapter 2,
and based on the results of preliminary experiments at temperatures less than 190
o
C, for which there was no significant uptake of CO
2
in LDH2. When the
temperature reached the preset point, the sorption system was kept at the same
temperature for 70 min in order to stabilize the TGA microbalance. Only after the
microbalance had stabilized (less than 10 μg change in weight), the purge gas was
then switched to either CO
2
(for the diffusivity measurements) or to an Ar-CO
2
gas
mixture (for the adsorption isotherm experiments) with the flow rate of 30 mL/min.
The other Mg-Al-CO
3
LDH sample (LDH3) was synthesized by the typical
co-precipitation reaction from aqueous solution
4-5
in order to investigate the particle
size effect on adsorption; its chemical composition was Mg
0.743
Al
0.257
(OH)
2
102
(CO
3
)
0.129
.
0.098(H
2
O) as determined by ICP-MS and TGA. As noted above, the
LDH3 was fractionated into 6 different particle sizes using Fisher Scientific U.S.
Standard test sieves, with their mesh sieve designations being 325, 230, 200, 170,
120, 80, and 70 (corresponding nominal sieve opening is 43, 63, 75, 90, 125, 180,
and 212 μm respectively). For each sieved LDH3 sample, diffusivity and adsorption
isotherm experiments for CO
2
were conducted at 200
o
C, following the same
experimental protocol as described previously for the LDH2. The surface area of
each sieved LDH3 was measured by a Micrometrics ASAP 2010 BET instrument
by the BET method at liquid nitrogen temperature; the micropore volume of LDH3
was also determined by the same instrument using the Horvath-Kawazoe method.
For the PFG NMR experiments referred to above, the LDH2 sample, whose
particle radii were in the range of 90 ~ 105 μm was utilized. PFG NMR spectra
were obtained by the Bruker Advance 500 NMR spectrometer with programmable
pulse field generating auxiliary. 400+ MHz rated, J Young valve-equipped NMR
tubes from Wilmad were utilized as NMR sample tubes, and 99.9% (atomic %)
13
C
enriched CO
2
from Sigma-Aldrich was used as the adsorbate.
4.3 Theory and Results
To estimate the diffusion coefficient of CO
2
in the LDH, Crank’s model for
diffusion into a spherical particle was used
10
. According to this model, the
governing equation for the concentration C of a species diffusing into a
homogeneous spherical particle of radius r is given as:
∂C
∂t
= D
∂
2
C
∂r
2
+
2
r
∂C
∂r
⎛
⎝
⎜
⎞
⎠
⎟
(1)
where D is the diffusivity. The use of Eqn. 1 to describe transport in the LDH
sample implies the existence of only one type of species diffusing and adsorbing,
and that there are no separate gas and surface phases co-existing, which is
consistent the structural (XRD) data for these materials.
From the solution of Eqn. (1), the uptake M
t
at time t is given by:
M
t
M
∞
= 1−
6
π
2
1
n
2
exp −n
2
⋅π
2
⋅
D
r
2
⋅t
⎛
⎝
⎜
⎞
⎠
⎟
(2)
n=1
∞
∑
where M
∞
is the uptake at large times (equilibrium). For small times (M
t
/M
∞
< 0.25)
Eqn. (2) is approximated as:
M
t
M
∞
=
6 ⋅ D
1/2
⋅ t
1/2
π
1/2
⋅ r
(3)
103
(a)
01 234 5 6 789 10 11
0.00
0.02
0.04
0.06
0.08
0.10
0.12
0.14
0.16
0.18
0.20
0.22
0.24
0.26
M
t
/ M
oo
Time, sec
1/2
M
t
/ M
oo
< 0.25
(b)
0 50 100 150 200
-1.4
-1.3
-1.2
-1.1
-1.0
-0.9
-0.8
-0.7
-0.6
ln ( 1 - M
t
/ M
oo
)
Time, sec
Data
Linear Fit
95% Confidence Limit
0.5 < M
t
/ M
oo
< 0.75
Figure 4.1 A graph of (a) M
t
/M
∞
against t
1/2
, and (b) ln(1- M
t
/M
∞
) against t for the
uptake of carbon dioxide at 200
o
C.
104
(a)
01 2 3 45678 9 10 11
0.00
0.02
0.04
0.06
0.08
0.10
0.12
0.14
0.16
0.18
0.20
0.22
0.24
0.26
M
t
/ M
oo
Time, sec
1/2
M
t
/ M
oo
< 0.25
(b)
0 204060 80
-1.4
-1.3
-1.2
-1.1
-1.0
-0.9
-0.8
-0.7
0.5< M
t
/ M
oo
< 0.75
Time, Sec
ln ( 1 - M
t
/ M
oo
)
Data
Linear Fit
95% Confidence Limit
Figure 4.2 A graph of (a) M
t
/M
∞
against t
1/2
, and (b) ln(1- M
t
/M
∞
) against t for the
uptake of carbon dioxide at 225
o
C.
105
(a)
01 234 5 6 7 8 9
0.00
0.02
0.04
0.06
0.08
0.10
0.12
0.14
0.16
0.18
0.20
0.22
0.24
M
t
/ M
oo
Time, sec
1/2
M
t
/ M
oo
< 0.25
(b)
0 10203040 50
-1.3
-1.2
-1.1
-1.0
-0.9
-0.8
-0.7
Time, Sec
ln ( 1 - M
t
/ M
oo
)
Data
Linear Fit
95% Confidence Limit
0.5< M
t
/ M
oo
< 0.75
Figure 4.3 A graph of (a) M
t
/M
∞
against t
1/2
, and (b) ln(1- M
t
/M
∞
) against t for the
uptake of carbon dioxide at 250
o
C.
106
Therefore, if Eqn. (3) applies, the plot of M
t
/M
∞
versus t
1/2
should give a linear
relationship, with a slope of 6D
1/2
/π
1/2
r. However, as shown in Figs. 4.1a, 4.2a, and
4.3a, the experimental results do not follow linear plots. The short-time nonlinear
M
t
/M
∞
against t
1/2
behavior was also reported by Ruthven for a number of
crystalline materials, but no reasons for the behavior were provided
11
.
Since nonlinear behavior was observed in the short-time region, the long
time region was selected for the diffusivity constant calculations. For long times (M
t
/M
∞
> 0.5) the higher-order terms in Eqn. (2) become negligible so that the
expression simplifies to:
2
22
6
1 exp (4)
t
M Dt
Mr
π
π
∞
⎛⎞ −⋅ ⋅
=−
⎜⎟
⎝⎠
Therefore, a plot of ln(1- M
t
/M
∞
) versus t should be linear with a slope of -π
2
D/r
2
and intercept of ln(6/π
2
). In this region, fairly good linearity is observed with the
experimental data at all three temperatures as shown in Figs. 4.1b, 4.2b, and 4.3b.
The diffusion constants for CO
2
in LDH were estimated from the slopes of the plots
at long times at each elevated temperature, and the results are summarized in
Table 4.2. According to Ruthven, linearity in the long time region implies that the
107
1.90 1.95 2.00 2.05 2.10 2.15
2.0x10
-8
4.0x10
-8
6.0x10
-8
8.0x10
-8
1.0x10
-7
1.2x10
-7
1.4x10
-7
D, cm
2
/s
1000 / T, K
-1
Figure 4.4 Temperature dependence of diffusion coefficient for CO
2
in LDH2.
108
109
distribution of the crystal sizes of LDH is not significantly wide
11
. The temperature
dependence of the CO
2
diffusivity is shown in Fig. 4.4, and the activation energy of
diffusion was calculated as 52.86 kJ/mol (12.64 kcal/mol) by the Arrhenius
equation.
As a parallel study, diffusion coefficients were also calculated by molecular
dynamic simulations at each temperature
12
; it is obvious that the resulting values
from molecular dynamic simulation are in good qualitative agreement with
experiments, as shown in Table 4.2.
Table 4.2 Diffusivity constants measured by experiment and calculated by
molecular dynamic simulation.
Temperature,
o
C
by molecular dynamic simulation,
cm
2
/s
by experiment, cm
2
/s
200 3.23 x 10
-7
3.61 x 10
-8
225 4.84 x 10
-7
8.90 x 10
-8
250 5.78 x 10
-7
1.33 x 10
-7
The measurement of molecular diffusion using PFG NMR experiments is
based on observation of NMR signal attenuation (Ψ) after an appropriate sequence
of radiofrequency pulses and inhomogeneous magnetic field (field gradient) pulses
are applied. Spin-echo attenuations (Ψ) in a homogeneous system obey the relation
given as
13-15
:
Ψ= exp −γ
2
⋅δ
2
⋅ g
2
⋅ r
2
Δ−
1
3
δ
⎛
⎝
⎜
⎞
⎠
⎟
⎛
⎝
⎜
⎞
⎠
⎟
/6
⎡
⎣
⎢
⎤
⎦
⎥
where δ, g, and Δ denote the width, amplitude, and separation of the magnetic field
gradient pulses, respectively (see Fig. 4.5).
The Einstein relation relating the mean square displacement and the observation
time is given as:
r
2
(t) = 6 D t
And, with the Einstein relation, Ψ may be modified as:
Ψ= exp −γ
2
⋅ D ⋅δ
2
⋅ g
2
Δ−
1
3
δ
⎛
⎝
⎜
⎞
⎠
⎟
⎡
⎣
⎢
⎤
⎦
⎥
Therefore, the observation of molecular displacement by PFG NMR can provide the
diffusivity constant directly.
Prior to PFG NMR experiments with the LDH sample, NMR spectrum was
obtained with an empty NMR tube for background correction. And the acquired
PFG NMR spectra were subtracted with background to obtain real signals. It was
110
Figure 4.5 Schematic representation of the gradient-echo pulse sequence
13
. (a)
Radiofrequency pulse sequence. (b) Associated magnetic field gradient pulses
(shaded areas). The time between the start of the π/2 pulse and the center of the
primary echo, during which T
2
relaxation takes place, is denoted by 2τ; Δ is the
time between the leading edges of the field gradient pulses, g is the magnitude of
the magnetic field gradient pulses (mTm
-1
), and δ is their duration.
111
Figure 4.6 The PFG NMR spectrum of LDH, which the spectrum of a blank NMR
tube was subtracted.
112
observed that the CO
3
peaks in PFG NMR spectra were too weak to be
distinguished from the noise as shown in Fig. 4.6. The low intensity of CO
3
peak is
mainly due to the low concentration of
13
C in the LDH sample; to be able to utilize
the PFG NMR technique for further investigation of the diffusivity constant one
must, therefore, synthesize
13
C enriched LDH samples.
To analyze the adsorption isotherm data, first the experimental data were
fitted with the Langmuir equation
16
, described as
q
CO
2
=
m
CO
2
b
CO
2
P
CO
2
1+ b
CO
2
P
CO
2
where is the equilibrium adsorbed CO
2
concentration, is the saturation
coverage, is the Langmuir adsoprtion constant for CO
2
, and the partial
gas phase CO
2
pressure. The experimental data and the fitted curves are shown in
Fig. 4.7, and the values of parameters in the Langmuir model are summarized in
Table 4.3. The Langmuir constant represents how strongly the molecule is
adsorbed onto an adsorbent surface. And the temperature dependence of can
be written as
q
CO
2
m
CO
2
b
CO
2
P
CO
2
b
CO
2
b
CO
2
b
CO
2
= b
∞
exp(Q / R
g
T )
113
0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9
0.00
0.05
0.10
0.15
0.20
0.25
0.30
partial pressure of CO
2
, bar
CO
2
uptake, mmol/g
200
o
C; chi
2
= 0.0001
225
o
C; chi
2
= 0.00006
250
o
C; chi
2
= 0.00004
Langmuir model
0 5 10 15 20 25 30
3
4
5
6
7
8
9
10
11
1/ CO
2
uptake, [mmol/g]
-1
1/ CO
2
partial pressure, bar
-1
200
o
C; R
2
= 0.9946
225
o
C; R
2
= 0.9952
250
o
C; R
2
= 0.9962
Linearized Langmuir model
Figure 4.7 The experimental data and nonlinear curve fitting with the Langmuir
equation for adsorption isotherm of CO
2
in LDH2 (up), and the experimental data
and linear fitting with the linearied form of Langmuir equation (bottom).
114
Table 4.3 Langmuir adsorption parameters of CO
2
in LDH2.
Temperature,
o
C
m
CO
2
, mmol/g sample
b
CO
2
, bar
-1
200 0.27904 15.6125
225 0.26136 17.8243
250 0.24189 22.1608
where Q is the heat of adsorption (-ΔH
ad
), and R
g
is ideal gas constant (8.314
J/mole⋅K). The heat of adsorption (Q) of LDH2 was estimated as -2.8739 kJ/mol
from the slope of a linear plot of ln versus 1/T, which the negative value
meaning that the process is slightly endothermic. Normally, the adsorption constant
decreases with the temperature, since adsorption is usually exothermic, and it is
unclear why it is not so for adsorption on LDH. Potentially, the increase in the
adsorption constant may signify structural changes in the material. As it can be seen
in Table 4.3, the saturation coverage also changes with temperature, which is also
consistent with the idea that the material is undergoing structural changes.
b
CO
2
A more general equation for the adsorption isotherm has been suggested by
Sips
19
:
q
CO
2
= m
CO
2
(b
CO
2
P)
1/ n
1+ (b
CO
2
P)
1/ n
115
The Sips equation is similar to the Langmuir equation, but with an additional
parameter n. When the parameter n is unity, the equation becomes the classical
Langmuir model for single-site adsorption. The experimental data and fitted curves
with the Sips equation are shown in Fig. 4.8, and the values of the parameters in the
model are summarized in Table 4.4
Table 4.4 The Sips equation parameters for CO
2
in LDH2.
Temperature,
o
C
m
CO
2
, mmol/g sample b
CO
2
, bar
-1
n
200 0.25677 16.8464 0.69837
225 0.24339 18.6860 0.72111
250 0.23116 22.3253 0.79789
From Table 4.4, it is observed that the value of n and increase with
temperature. On the other hand the value of saturation coverage decreases with
temperature.
b
CO
2
Though the Sips equation provides some sense of the system’s non-ideality,
it does not possess the correct Henry law behavior. A third type of empirical
adsorption isotherm is due to Toth
20-22
. The Toth equation for the CO
2
adsorption
may be written as:
116
0.0 0.2 0.4 0.6 0.8
0.00
0.05
0.10
0.15
0.20
0.25
Langmuir-Freundlich model
Partial Pressure of CO
2
, bar
CO
2
uptake, mmol/g
200
o
C; chi
2
=0.00004
225
o
C; chi
2
=0.00002
250
o
C; chi
2
=0.00004
Figure 4.8 The experimental data and nonlinear curve fitting with the Sips
(Langmuir-Freundlich) equation for adsorption isotherm of CO
2
in LDH2.
117
q
CO
2
= m
CO
2
bP
[1+ (bP)
t
]
1/t
And it also contains an additional parameter (when t is equal to 1, the Toth isotherm
equation reduces to the Langmuir equation). The experimental data and fitted
curves are shown in Fig. 4.9, and the values of the parameters in the Toth model are
summarized in Table 4.5). It is observed that the values of the parameter t are
greater than unity, which again indicates a departure from the simple one-site
Langmuir-type behavior.
Table 4.5 The Toth equation parameters for CO
2
in LDH2.
Temperature,
o
C
m
CO
2
, mmol/g sample b
CO
2
, bar
-1
t
200 0.25218 10.19558 1.80226
225 0.24019 11.86695 1.64989
250 0.22934 15.97911 1.38246
The chi
2
values for nonlinear fittings of three isotherm equations were
summarized in Table 4.6. Based on chi
2
values, the experimental data were fitted
better by Toth model than either by Langmuir or Sips model.
118
0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9
0.00
0.05
0.10
0.15
0.20
0.25
Partial Pressure of CO
2
, bar
CO
2
uptake, mmol/g
Toth model
200
o
C; chi
2
=0.00003
225
o
C; chi
2
=0.00001
250
o
C;; chi
2
=0.00002
Figure 4.9 The experimental data and nonlinear curve fitting with the Toth equation
for adsorption isotherm of CO
2
in LDH2.
119
Table 4.6 The chi
2
values of nonlinear fittings for different models
Temperature,
o
C Langmuir model Sips model Toth model
200 10 x 10
-5
4 x 10
-5
3 x 10
-5
225 6 x 10
-5
2 x 10
-5
1 x 10
-5
250 4 x 10
-5
4 x 10
-5
2 x 10
-5
For the empirical equations described so far, the adsorption mechanism is
assumed to be surface layering (formation of successive layers). However, for
microporous solids, another important adsorption mechanism is pore filling, which
was originally developed by Dubinin
23-26
. Also, according to the suggestion of
Bering et al.
27, 28
, the adsorption in pores less than 15Å should follow the pore
filling mechanism rather than surface coverage. For the description of adsorption
isotherm in microporous solids with pore filling mechanism, a widely used semi-
empirical equation is the Dubinin-Radushkevich (DR) equation, which may be
written as:
C = C
s
exp −
1
E
2
R
g
T ln
P
P
0
⎛
⎝
⎜
⎞
⎠
⎟
2
⎡
⎣
⎢
⎢
⎤
⎦
⎥
⎥
120
121
Table 4.7 The characteristic energies for CO
2
in LDH2 with DR equation.
Temperature,
o
C E, kJ/mol
200 12.350
225 13.677
250 15.782
where C is the amount adsorbed, C
s
is maximum adsorption capacity, E is
characteristic energy, R
g
is ideal gas constant. The increase of characteristic energy
means the adsorption is stronger since the solid has stronger energy of interaction
with adsorbate. To find the values of characteristic energy E, the adsorption data of
different temperatures were plotted as the logarithm of the fractional loading versus
the square of logarithm of reduced pressure. From the slopes of the linear fit, the
values of E were calculated. The experimental data and linear fit are shown in Fig.
4.10, and the summary of results is shown in Table 4.7. It is observed that the
characteristic energy increased as the temperature increased, and, therefore, it
implies that the interaction between LDH and CO
2
also increases as temperature
increases. It may signify the structural changes in the material, which it is also
024 6 8 10 12
-1.1
-1.0
-0.9
-0.8
-0.7
-0.6
-0.5
-0.4
-0.3
-0.2
-0.1
0.0
ln ( C / C
s
)
[ ln ( P / P
0
) ]
2
200
o
C; R=0.99773
225
o
C; R=0.99719
250
o
C; R=0.99688
Figure 4.10 The experimental data and linear fitting with linearized DR equation
for adsorption isotherm of CO
2
in LDH2.
122
123
consistent with the idea that the material is undergoing structural changes with
temperature increase.
For the study of particle size effect, CO
2
uptake experiments, and
adsorption isotherms with sieved LDH3 particles were conducted following the
same experimental as with the LDH2 sample. The particle size distribution of each
sieved section was shown in Fig. 4.12. To obtain particle size distribution of LDH3,
several SEM images were taken for each fractionated LDH3 particles (some images
are shown in Fig. 4.11). And then the acquired images were processed with
ImageJ
®
software to count the numbers of different particle radii.
Each LDH3 sample (consisting of particles with different average pore
sizes) was heated to 200
o
C in dry UHP argon flow at the rate of 30 ml/min with a
heating rate of 5
o
C/min. When the temperature reached 200
o
C, the sample was
kept at this temperature for 70 min in order for the TGA microbalance to stabilize
(less than 10 μg change). Then the Ar purge gas was switched to CO
2
with a flow
rate of 30 ml/min. It was observed that at the temperature of 200
o
C the total
amount of CO
2
adsorbed at saturation on LDH3 (in terms of mmol CO
2
/g LDH)
decreases as the average particle radius increases (Fig. 4.13a). However, if the
(a)
(b)
(c)
(d)
(e)
(f)
Figure 4.11 SEM images of LDH3 particles for fractionated sections. The range of
particle diameter: (a) 43~63 μm (b) 63~75 μm (c) 75~90 μm (d) 90~125 μm (e)
125~180 μm (f) 180~215 μm.
124
(a) (b)
0
5
10
15
20
25
17.50 21.25 23.75 26.25 28.75 31.25 33.75 37.50
Radius, μm
Wt(%)
0
5
10
15
20
25
30
35
28.25 30.75 32.25 33.75 35.25 36.75 38.25 40.75
Radius, μm
W t(% )
(c) (d)
0
5
10
15
20
25
34.25 36.5 38.5 40.5 42.5 44.5 46.5 48.75
Radius, μm
Wt(%)
0
5
10
15
20
25
42.5 46.5 49.5 52.5 55.5 58.5 61.5 65.5
Radius, μm
Wt(%)
(e) (f)
0
5
10
15
20
25
30
60.5 64.5 67.5 70.5 73.5 76.5 79.5 83.5
Radius, μm
Wt(%)
0
5
10
15
20
25
30
82.5 87.5 92.5 97.5 102.5 107.5 112.5 117.5
Radius, μm
Wt(%)
Figure 4.12 Particle size distribution of each sieved section of LDH3. The range of
particle diameter: (a) 43~63 μm (b) 63~75 μm (c) 75~90 μm (d) 90~125 μm (e)
125~180 μm (f) 180~215 μm.
125
20 30 40 50 60 70 80 90 100
0.42
0.44
0.46
0.48
0.50
0.52
0.54
0.56
0.58
0.60
0.62
0.64
0.66
Data
polynomial fit
Equation:
y = A + BX + CX
2
+ DX
3
Chi
2
= 0.00003
R
2
= 0.99834
A 0.97432
B -0.01697
C 0.00019
D -7.2161E-7
Particle radius, μm
CO
2
uptake, mmol/g
20 30 40 50 60 70 80 90 100
5.0x10
-3
1.0x10
-2
1.5x10
-2
2.0x10
-2
2.5x10
-2
3.0x10
-2
3.5x10
-2
Particle radius, μm
CO
2
uptake, mmol/m
2
Figure 4.13 The uptake amount of CO
2
with different LDH3 particle sizes at 200
o
C
126
127
Table 4.8 The uptake amount and BET surface area of LDH3 particles.
Particle radius
[μm]
Uptake
[mmol/g]
BET surface area
[m
2
/g]
Normalized uptake
[mmol/m
2
]
26.5 0.64545 36.2732 0.01779
34.5 0.57727 32.9144 0.01754
43.75 0.53500 30.8792 0.01733
53.75 0.49318 30.0916 0.01639
76.25 0.45227 25.4140 0.01780
98.75 0.43636 21.8627 0.01996
adsorbed amount at saturation is presented in terms of (mmol CO
2
/m
2
BET surface)
as in Fig 4.13b then it is fairly insensitive to the average particle size. According to
Table 4.8, it is evident that the BET surface area of LDH3 particle is decreasing
when the particle radius is increasing, but it is not clear the reason why it is so.
Adsorption isotherm experiments were also carried out with the LDH3
samples of varying average particle size. For each point of the isotherm a fresh
LDH3 sample was loaded at room temperature in the TGA balance, and the sample
was heated to 200
o
C in UHP Ar flow (30 mL/min) with a heating rate of 5
o
C/min.
After waiting for the balance to stabilize, the Ar flow was changed to an Ar + CO
2
gas mixture of predetermined CO
2
concentration. The sorption system was allowed
to reach equilibrium. The system was then cooled down to room temperature in an
Ar flow, and then the LDH3 sample was exchanged with a fresh one (with the same
particle distribution) and the experiment was repeated but with an Ar +CO
2
mixture
with a different CO
2
partial pressure, the experiment repeated each time with a fresh
LDH3 sample till the whole adsorption isotherm was generated.
Table 4.9 The Langmuir adsorption parameters of CO
2
in LDH3 at 200
o
C with
different particle size.
Particle radius, μm
m
CO
2
, mmol/g sample
b
CO
2
, bar
-1
26.5 0.75129 10.5694
34.5 0.65995 11.6985
43.75 0.56710 10.6203
53.75 0.52253 11.0997
76.25 0.51523 13.2044
98.75 0.43705 13.6583
128
The experimental data and the fitted curves with the Langmuir equation for
LDH3 particles are shown in Fig. 4.14, and the values of parameters are
summarized in Table 4.9. According to Table 4.9 and Fig. 4.14(b), it is evident that
Langmuir parameter is relatively constant over the range of particle sizes, and
b
CO
2
the value is close to the one obtained from Langmuir equation with LDH2 at
200
o
C in Table 4.3. On the other hand the value of (mmol CO
2
/g LDH)
decreases as the average particle increases, but the value expressed in terms
of (mmol CO
2
/m
2
LDH) as shown in Table 4.8 stays fairly invariable with average
particle size.
b
CO
2
m
CO
2
m
CO
2
The experimental data were also studied with Sips. The nonlinear fitting
curves are shown in Fig. 4.15, and the values of parameters are summarized in
Table 4.10. As shown in Fig. 4.15 and Table 4.10, the values for and n are not
changing significantly as a function of particle size.
b
CO
2
Table 4.10 The Sips isotherm equation parameters of CO
2
in LDH3 at 200
o
C with
different particle size.
Particle radius,
μm
m
CO
2
, mmol/g sample b
CO
2
, bar
-1
n
26.5 0.68824 12.33558 0.76988
34.5 0.6578 11.77357 0.79127
43.75 0.53429 11.85758 0.74523
53.75 0.48755 12.55263 0.71473
76.25 0.47115 12.96453 0.72593
98.75 0.41754 13.05532 0.74152
129
(a)
0.00.1 0.20.3 0.40.5 0.60.7 0.8
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
CO
2
Partial pressure, bar
CO
2
uptake, mmol/g
Average particle radius
26.50 μm 34.50 μm
43.75 μm 53.75 μm
76.25 μm 98.75 μm
(b)
0 5 10 15 20 25 30 35
0
1
2
3
4
5
6
7
8
9
Average particle radius
26.50 μm 34.50 μm
43.75 μm 53.75 μm
76.25 μm 98.75 μm
1 / CO
2
uptake, [mmol/g]
-1
1 / partial pressure of CO
2
, bar
-1
Figure 4.14 (a) The experimental data and nonlinear curve fitting with the
Langmuir equation, and (b) the experimental data and linearized Langmuir equation
for adsorption isotherm of CO
2
in LDH3.
130
(a)
0.00.1 0.20.3 0.40.5 0.60.7 0.8
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
Average particle radius
26.50 μm 34.50 μm
43.75 μm 53.75 μm
76.25 μm 98.75 μm
CO
2
Partial pressure, bar
CO
2
uptake, mmol/g
(b)
20 30 40 50 60 70 80 90 100
2
4
6
8
10
12
14
16
18
20
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
n
b
CO
2
, bar
-1
Particle radius, μm
Figure 4.15 (a) The experimental data and nonlinear curve fitting with the
Langmuir-Freundlich equation, and (b) the parameter values of Langmuir-
Freundlich equation for adsorption isotherm of CO
2
in LDH3.
131
132
Table 4.11 The raw data of CO
2
uptake of sieved LDH3.
43~63 μm 63~75 μm 75~90 μm
Time, sec Weight, mg Time, sec Weight, mg Time, sec Weight, mg
0 123.976 0 126.213 0 119.491
10 124.0238 10 126.2504 10 119.5226
20 124.222 20 126.383 20 119.5571
30 124.515 30 126.6638 30 119.5895
40 125.6795 40 127.1651 40 119.7421
50 126.5217 50 128.435 50 120.2252
60 126.9543 60 128.9487 60 120.8812
70 127.2369 70 129.2346 70 121.3838
80 127.3609 80 129.366 80 121.7555
90 127.4775 90 129.4188 90 121.9105
100 127.4969 100 122.1381
120 122.2164
130 122.2849
140 122.3038
CO
2
diffusivity constants in sieved LDH3 were also measured by the
gravimetric method following the same experimental and analysis procedures in
long time region as previously described. The CO
2
uptake raw data of each
fractionated section were shown in Table 4.11.
The measured diffusivity values appear to be dependent on the particle size,
increasing with increasing particle size, as shown in Fig. 4.16. The observed trends
Table 4.11 The raw data of CO
2
uptake of sieved LDH3 (continued).
90~125 μm 125~180 μm 180~215 μm
Time, sec Weight, mg Time, sec Weight, mg Time, sec Weight, mg
0 123.159 0 118.652 0 116.007
10 123.1719 10 118.6634 10 116.0139
20 123.2105 20 118.6975 20 116.0236
30 123.3848 30 118.9223 30 116.0305
40 123.6434 40 119.0328 40 116.1061
50 124.1177 50 119.499 50 116.1549
60 124.5312 60 119.8643 60 116.3096
70 125.1615 70 120.4827 70 116.542
80 125.4926 80 120.7683 80 116.7914
90 125.6558 90 120.9001 90 117.4566
100 125.7413 100 120.961 100 117.8274
110 125.7852 110 120.9891 110 118.1274
120 125.8078 120 121.0021 120 118.1613
130 125.8316 130 121.0132 130 118.196
140 118.2177
150 118.2272
160 118.2343
are very similar to Badruzzaman et al’s findings, in their study of arsenic adsorption
granular ferric hydroxide
2
.
For better understanding of the nonlinear relation between diffusivity
coefficients and LDH3 particle radii, Crank’s solution was modified with
introduction of the particle size distribution function f(R). First, assuming that the
radius of LDH3 particle is uniformly distributed in the system, and if and
t
M
133
20 30 40 50 60 70 80 90 100
0
5
10
15
20
25
Particle radius, μm
D / R
2
x 10
3
, sec
-1
30 40 50 60 70 80 90 100
10
D x 10
7
, cm
2
/sec
Particle radius, μm
Y = - 2.39736 + 2.1463 X
D = 0.004 x R
2.1463
R
2
= 0.97985
Figure 4.16 Nonlinear relation between diffusion constants and particle radii of
LDH3.
134
M
∞
are divided by the weight of the particle
3
4
3
P
R π ρ
⎛⎞
⋅⋅ ⋅
⎜⎟
⎝⎠
, one gets
and
which they are the corresponding the amounts adsorbed per gram sample. Since
reflects equilibrium conditions, it is independent of radius and it only depends
on the particle type and prevailing conditions.
M
t
'
M
∞
'
M
∞
'
If a total weight of a given sample is defined as W
t
in the particle range of
[
], and if W
t
f(R)dR is the weight of particles with the radii between R and
R+dR (where dR is very small so that the uptake per gram is described by Eqn. (1)),
then (the total amount adsorbed on particles with radii between R and
R+dR) is given by the following equation :
R
e
, R
u
M
t
(R)dR
'22
22 2
1
61
() 1 exp ()
t t
n
D
M RdR M n t W f RdR
nR
π
π
∞
∞
=
⎧⎫ ⎛⎞ ⎡⎤
=⋅ − − ⋅ ⋅ ⋅ ⋅ ⋅
⎨⎬
⎜⎟
⎢⎥
⎣⎦ ⎝⎠ ⎩⎭
∑
M
t
T
which is defined as the amount adsorbed on the total sample, which consist of
particles with radii between R
e
and R
u
, is given as : M
t
T
()
'22
22 2
1
22
22 2
1
()
61
1exp (
61
1exp ()
u
e
u
e
u
e
R
T
tt
R
R
tt
n
R
R
T
n
R
MMRdR
D
) MWnt
nR
D
Mnt
nR
π
π
π
π
∞
=
∞
∞
=
=
⎧⎫ ⎛⎞ ⎡⎤
=⋅ ⋅ − ⋅ −⋅ ⋅ ⋅ ⋅
⎨⎬
⎜⎟
⎢⎥
⎣⎦ ⎝⎠ ⎩⎭
⎧⎫ ⎛⎞ ⎡⎤
=⋅ − ⋅ − ⋅ ⋅ ⋅ ⋅
⎨⎬
⎜⎟
⎢⎥
⎣⎦ ⎝⎠ ⎩⎭
∫
∑
∫
∑
∫
fRdR
fRdR
135
Table 4.12 Polynomial form of particle size distribution function, f(R)= Ax
6
+ Bx
5
+
Cx
4
+ Dx
3
+ Ex
2
+ Fx+ G .
A B C D E F G
43~63μm -0.057 1.4865 -14.998 73.716 -182.85 217.38 -91.892
63~75μm -0.0143 0.3893 -4.0321 19.773 -46.163 46.374 -9.7863
75~90μm 0.0034 -0.1006 1.1923 -7.0714 21.188 26.198 14.788
90~125μm -0.0207 0.53 -5.2103 24.904 -59.357 67.831 -24.098
125~180μm 0.0425 -1.2566 14.737 -86.627 264.58 389.72 224.45
180~215μm 0.0612 -1.5156 14.453 -67.255 159.93 179.41 74.839
Assuming the particle size distribution is uniform, and then f(R) may be written as:
f (R) =
1
R
u
− R
e
And, above equation can be written as:
22
22 2
1
16 1
1exp
u
e
R
T
t
T
n
ue R
M D
nt
MR R n R
π
π
∞
=
∞
dR
⎧ ⎫ ⎛⎞ ⎡⎤
=⋅ − ⋅ −⋅⋅⋅
⎨ ⎬
⎜⎟
⎢⎥
−
⎣⎦ ⎝⎠ ⎩⎭
∑
∫
As shown in Fig. 4.12, it is evident that particle size distribution is not uniform, and
therefore the assumption that f(R) is uniform is incorrect. To obtain more realistic
particle size distribution functions, 6
th
order (the highest order MS Excel
®
software
allows) polynomial equations were extracted by nonlinear fitting of each curve, and
136
137
those were summarized in Table 4.12. The extracted f(R) equations were intended to
be inserted into modified Crank’s equation above, and it is necessary to utilize
numerical methods to solve the equation.
4.4 Conclusions
The adsorption isotherms for carbon dioxide on Mg-Al-CO
3
were
investigated at three different temperatures between 200
o
C and 250
o
C by
gravimetric method. The experimental diffusivity were also obtained by
thermogravimetric experiments. The experimental values were compared with those
obtained by molecular dynamic simulations and are found to be in good qualitative
agreement.
The experimental adsorption isotherms for CO
2
on LDH have been
analyzed with the Langmuir isotherm equation and two additional empirical
adsorption isotherm equations. It is found that the experimental data are fitted the
best with the Toth equation, based on chi
2
values.
We also investigated the effect of particle size on the transport and sorption
of CO
2
in LDH. It is observed that the total amount of CO
2
adsorbed (expressed in
terms of mmol CO
2
/g adsorbent) increases as the particle size decreases; on the
other hand, the BET surface area increases as the particle size decreases. As a result,
when the amount adsorbed was normalized with respect to the BET surface area, it
was found that it remained fairly constant with respect to the particle size.
Adsorption isotherm data were also analyzed with respect to the particle
size. When studied with respect to the Sips isotherm equation it was observed that
the values of and n remained relatively constant for the whole range of particle
sizes. (mmol/g), on the other hand, changed with respect to the particle size.
However, when scaled with respect to the surface area it also remained fairly
constant with respect to the particle.
b
CO
2
m
CO
2
The diffusivities when calculated with a homogeneous diffusion model
were shown to decrease as a function of particle size. However, to understand the
observations, more investigation is needed by numerical methods.
138
139
Chapter 4 References
1. Callaghan, P. T. Principles of NMR Microscopy, Clarendon Press, Oxford,
1991.
2. Badruzzaman, M; Westerhoff, P; Knappe, D. R. U. Intraparticle diffusion
and adsorption of arsenate onto granular ferric hydroxide., Water Research,
2004, 38, 4002-4012.
3. Grande, C. A.; Basaldella, E.; Rodrigues, A. E. Crystal size effect in
vacuum pressure-swing adsorption for propane/propylene separation., Ind.
Eng. Chem. Res., 2004, 43, 7557-7565.
4. Miyata, S. The syntheses of hydrotalcite-like compounds and their
structures and physico-chemical properties I. The systems magnesium(2+)-
aluminum(3+)-nitrate(1-), -chloride(1-) and -perchlorate(1-), nickel(2+)-
aluminum(3+)-chloride(1-), and zinc(2+)-aluminum(3+)-chloride(1-).
Clays and Clay Minerals, 1975, 23, 369.
5. Prinetto, F.; Ghiotti, G.; Graffin, P.; Tichit, D. Synthesis and
characterization of sol-gel Mg/Al and Ni/Al layered double hydroxides and
comparison with co-precipitated samples. Microporous and Mesoporous
Materials, 2000, 39, 229-247.
6. Chagger, H. K.; Ndaji, F. E.; Sykes, M. L.; Thomas, K. M. Kinetics of
adsorption and diffusional characteristics of carbon molecular sieves,
Carbon, 1995, 33, 1405.
7. Kapoor, A.; Yang, T. Kinetic separation of methane-carbon dioxide mixture
by adsorption on molecular sieve carbon, Chem. Eng. Sci., 1989, 44, 1723.
140
8. Yang, R. T.; Baksh, M. S. A. Pillared clays as a new class of sorbents for
gas separation, AIChE J., 1991, 37, 679.
9. Yucel, H.; Ruthven, D. M. Diffusion of Carbon dioxide in 4A and 5A
zeolite crystals, J. Colloid Interface Sci., 1980, 74, 186.
10. Crank, J. The Mathematics of Diffusion, 2
nd
Edition, Clarendon Press,
Oxford, 1975.
11. Ruthven, D. M., in Adsorption, Science and Technology (Edited by
Rodriguez, A. E.), Kluwer Academic Publishers, Netherlands, 87-114, 1989.
12. Kim, N.; Kim, Y.; Tsotsis, T. T.; Sahimi, M., Atomic simulation of
nanoporous layered double hydroxide materials and their properties. I.
Structural modeling., J. Chem. Phys., 2005, 122, 214713.
13. Kärger, J.; Pfeifer, H.; Heink, W. Principles and application of self-
diffusion measurement by nuclear magnetic resonance, Adv. Mag. Res.
1988, 12, 1.
14. Abragam, A. Principles of Nuclear Magnetism, Oxford University Press,
New York, 1961.
15. Stejskal, E. O.; Tanner, J. E. Spin diffusion measurement: spin echoes in the
presence of a time-dependent field gradient. J. Chem. Phys. 1965, 42(1),
288.
16. Langmuir, I. The adsorption of gases on plane surface of glass, mica, and
platinum, J. Ame. Chem. Soc., 1918, 40, 1361.
17. Freundlich, H. Kinetics and energetics of gas separation, Trans. Farad. Soc.,
1932, 28, 195.
141
18. Rudzinski, W.; Everett, D. H., Adsorption of Gases on Heterogeneous
Surfaces, Academic Press, San Diego, 1992.
19. Sips, R. Structure of catalyst surface, J. Chem. Phy., 1948, 16, 490.
20. Do, D. D. Adsorption Analysis: Equilibria and Kinetics, Series on
Chemical Engineering V ol. 2, Imperial College Press, London, 1988.
21. Toth, J. Uniform interpretation of gas/solid adsorption, Adv. Colloid
Interface Sci., 1995, 55, 1.
22. Toth, J. Adsorption. Theory, Modeling, and Analysis, Dekker, New York,
2002.
23. Dubinin, M. M., Porous structure and adsorption properties of active
carbon, Chemistry and Physics of Carbon, 1966, 2, 51.
24. Dubinin, M. M., Adsorption in micropores, J. Colloid Interface Sci., 1967,
23, 487.
25. Dubinin, M. M., Fundamentals of the theory of physical adsorption of gases
and vapors in micropores, Adsorption – desorption Phenomena, Proc. Int.
Conf., 1972, 2nd, 3.
26. Dubinin, M. M., Physical adsorption of gases and vapors in micropores,
Progress in Surface and Membrane Science, 1975, 9, 1.
27. Bering, B. P.; Dubinin, M. M.; Serpinsky, V. V., Theory of volume filling
for vapor adsorption, J. Colloid Interface Sci. 1966, 21, 378.
28. Bering, B. P.; Dubinin, M. M.; Serpinsky, V. V., Thermodynamics in
micropores, J. Colloid Interface Sci. 1972, 38, 185.
142
Chapter 5
Preparation and Characterization of Mg-Al-CO
3
LDH membranes
5.1 Introduction and overview
LDHs appear to be suitable for the preparation of membranes to be used in
environmental applications, but significant progress in this area has not been
achieved, so far, due to the difficulty of making defect-free thin continuous LDH
films. Efforts to prepare thin films typically result in powdery translucent
aggregates that do not adhere well to the substrate. In one case, reported in the
literature LDH films as thin as 100 nm were synthesized by co-precipitation; the
particles prior to film formation were separated by centrifugation to select only
colloidal size particles
1
. In preparation of thin films this separation step adds
significant complexity and also results in significant loss of material. Gardner et al.
utilized a procedure which did not require a fractionation step. He utilized alkoxide
intercalated derivatives of LDH materials, to prepare thin films on supported glass
2
consisting of aggregates of LDH particles with a lamellar texture. These assemblies
143
of LDH materials on glass can be used as functional coatings for chemical sensors
3,
4
or as clay-modified electrodes
5, 6
; however, they are not suitable to be used as
membranes for the gas separation.
To prepare LDH membranes for use in gas separation, we have studied
several methods, including in situ co-precipitation under well-controlled reaction
conditions, and sol-gels techniques combined with an anion exchange step. Co-
precipitation is the conventional method to prepare LDHs
7
in a powder form; to the
best of our knowledge, the co-precipitation method has not been utilized in the
literature to prepare continuous LDH film. In our studies α or γ alumina porous
support were placed in the reagent solution, and during the co-precipitation LDH
particles were deposited on the support to generate a film with a lamella texture.
The transport properties of these LDH films were studied with the measurement of
single gas permeances at room temperature; no separation for the CO
2
/N
2
gas pair
was observed.
To prepare LDH membranes using the sol-gel method, first a procedure
proposed in a number of recent papers
8, 9
was utilized. However, following this
procedure did not lead in the formation of a gel, as described in these papers.
10 20 30 40 50 60 70
0
50
100
150
200
Figure 5.1 The XRD spectrum of LDH powder synthesized by sol-gel.
144
145
Instead, a suspension of particles resulted, and these particles had a weak and noisy
XRD spectrum that resembled that of LDH, as shown in Fig. 5.1. Since it is difficult
to prepare LDH membranes with the aid of suspended powders, the procedure was
replaced by ethoxide intercalated LDH synthesis followed by anion exchange with
carbonates (the ethoxide groups can be easily exchanged by carbonate groups
10-14
).
Like the co-precipitation method, α or γ alumina substrates were utilized as porous
support materials, and the LDH membranes were prepared by dip-coating.
Transport properties of the resulting LDH membranes were investigated by single-
gas permeance measurements at ambient or elevated temperature; for the CO
2
/N
2
gas pair, the best result was close to the Knudsen factor (~1.25 in favor of N
2
).
A different approach we tried, in order to prepare LDH membranes,
involved separating the nucleation and aging steps
15
. Small LDH crystals were
generated by rapid termination of the co-precipitation reaction at ambient or sub-
ambient temperatures, and then they were deposited on the pores of supports by dip
coating or suction from the bottom of support with the aid of a mechanical pump.
The deposited small LDH crystallites were dried at room temperature, and then
were subjected to hydrothermal aging with the goal of closing the pores by growing
146
up the LDH crystallites. Porous stainless steel disks and α-alumina disks were
utilized as a porous support, and the presence of LDH crystallites was confirmed by
Mg peak in the energy dispersive x-ray spectroscopy (EDS) spectrum of the cross-
section of the coated supports. The transport properties of LDH membranes were
investigated by the measurements of single gas permeances at ambient
temperatures; the best results were obtained with the α-alumina support, however,
the separation factor calculated were similar to the Knudsen factors.
In the attempts to prepare LDH membrane described above, the results
were not better than the Knudsen factor, leading to the conclusion that the
aggregation of LDH crystals generates mesoporous materials, since the crystals do
not appear to intergrow significantly. The focus of our studies to prepare LDH
membranes, as a result, shifted to finding a technique to intergrow the LDH crystals
in order to eliminate the mesoporous spaces in between the crystals. In a recent
Japanese patent, sulfates were reported to be good binder for ceramic materials
containing Al
16
. The binding mechanism of sulfate group may be explained by
water removal followed by sulfate addition, as shown Fig. 5.2. We studied
addition of sulfates to LDH powders in order to prepare thin LDH films. LDH
membranes were prepared on α-alumina or SiC disks by dip-coating. Transport
properties of LDH membranes were investigated by the measurements of single gas
permeance at ambient temperature; the resulting permeances were low, however,
the separation factors were close to the Knudsen factor.
Al OH
H
+
Al
-H
2
O
Al
+
Al O S
O
O
O Al OH
2
+
SO
4
2-
2
Figure 5.2 The mechanism of the binding reaction between LDH and sulfate groups.
5.2 Experimental
FTIR spectra of synthesized LDH were recorded using a Genesis II
(Mattson, FT-IR) instrument, and XRD spectra were generated by Rigaku X-ray
diffractometer, as described in a prior chapter. The transport properties of the LDH
membranes were investigated by the measurement of single gas permeances using a
custom-made apparatus equipped with a bubble flow meter for the measurement of
147
148
gas flow. The permeabilities of gas mixtures were also measured on occasion in
which case the compositions of determined by an HP 5790A Gas Chromatograph or
by a custom-made mass spectrometer equipped with a MKS 100C UTI precision
gas analyzer.
For the preparation of membrane films by the co-precipitation method,
ACS reagent-grade anhydrous Na
2
CO
3
, Mg(NO
3
)
2
·6H
2
O and Al(NO
3
)
3
·9H
2
O from
Aldrich was used. The mixture of 0.9 M Mg(NO
3
)
2
, 0.3 M Al(NO
3
)
3
,
and 0.02 M
Na
2
CO
3
aqueous solution was prepared at room temperature, and then an α-alumina
tube was placed in the mixture. To initiate the reaction, a separately prepared 1.5 M
NaOH aqueous solution was added drop-wise to the mixture under vigorous stirring
at room temperature. When the addition of the NaOH solution was completed, the
mother liquor was hydrothermally aged. After the hydrothermal aging step, the
LDH-coated α-alumina tube was repeatedly washed with distilled water, and then
was dried at 70
o
C for 24 hr. FT-IR and XRD analysis of the resulting LDH
materials were carried out and were shown to be identical to those of LDH
materials reported in the previous chapters, as shown in Fig. 5.3.
(a)
4000 3500 3000 2500 2000 1500 1000 500
1620cm
-1
1370cm
-1
940cm
-1
3070cm
-1
Wavenumbers, cm
-1
3470cm
-1
(b)
10 20 30 40 50 60 70
0
2000
4000
6000
8000
10000
12000
14000
2θ
Figure 5.3 (a) FT-IR spectrum and (b) XRD spectrum of LDH powder synthesized
by coprecipitation method.
149
150
For the preparation of membranes by the sol-gel method, a solution of 0.9
M Mg(NO
3
)
2
and 0.3 M Al(NO
3
)
3
in absolute ethanol was prepared. Separately 40 g
of NaOH were dissolved in 100 mL of ethanol at a temperature of 60
o
C, and the
solution was refluxed for 6 hr at the same temperature. The ethoxide solution thus
prepared was added drop-wise to the reagent solution at room temperature, and,
after the addition was complete, the mixture was refluxed again at 80
o
C. After 3
days of reaction time, the mixture was washed with absolute ethanol by centrifuging
in sealed containers to minimize exposure to atmospheric moisture and to CO
2
. This
washing process was repeated until the pH of decanted solution was below 8. The
ethoxide intercalated LDH resulted in viscous slurry in ethanol, and it became a
translucent stable colloidal solution after hydrolysis. These translucent solutions
were used to prepare LDH films on tubular α- (or γ-) alumina supports by dip-
coating followed by evaporating the ethanol at room temperature.
For the method which involved separate nucleation and aging steps, an
aqueous solution of 0.9 M Mg(NO
3
)
2
, 0.3 M Al(NO
3
)
3
,
and 0.02 M Na
2
CO
3
was
prepared. A separately prepared 1.5 M NaOH aqueous solution was rapidly added to
the solution with vigorous stirring at ambient or subambient temperatures. After 2
151
min of vigorous stirring at room temperature, the mixture was washed with distilled
water by centrifuging in sealed containers until the pH of decanted solution was
below 8. The mixture became translucent solution after washing step, and these
translucent solutions were utilized to prepare LDH films on porous stainless steel or
α-alumina supports.
For the procedure utilizing sulfuric acid as a binder, 99% ACS reagent
grade sulfuric acid from VWR was used as the sulfate source. Commercial synthetic
hydrotalcite from Aldrich was utilized as the LDH material. A 1.9 M H
2
SO
4
solution in absolute ethanol was prepared at room temperature, and the LDH
powders were also dispersed in absolute ethanol by ultrasonic treatment for 15
minutes. The 1.9 M H
2
SO
4
solution was added drop-wise to the solution containing
the dispersed LDH. The mixture of LDH particles and sulfuric acid was prepared at
the ratio of 0.1 g of LDH per 1 mL of solution. After sulfate group addition was
completed, the mixture was sonicated for 10 min at ambient temperature. The LDH
membranes were prepared on LDH or SiC disks by dip coating.
152
5.3 Results and discussion
5.3.1 The co-precipitation method
The effect of hydrothermal aging on the transport properties of LDH
membranes prepared by co-precipitation was investigated with the measurement of
N
2
or CO
2
permeances. First, in order to study the effects of aging time on LDH
membrane characteristics, the coated α-alumna tubes were hydrothermally treated
in an autoclave for 6, 12, and 18 hr at 80
o
C. It was observed that the permeance
was not much affected by aging time, as shown in Fig. 5.4; the separation factors
were also not affected significantly by the aging time. For further investigation on
the effect of hydrothermal aging, six different temperatures were utilized (80, 100,
120, 140, 160, and 180
o
C) for aging for a pre-determined 6 hr period. It was
observed that there was no significant effect until a temperature of 120
o
C; negative
effects were observed when the aging temperature was higher than 120
o
C, as
shown Fig. 5.5.
The measured ideal separation factors for N
2
/CO
2
were 0.97 (average) in
favor of CO
2
using pressure drops (ΔP) in the rage of 20 ~ 60 psi, which means that
practically the membrane does not separate these two gases. This observation was
(a)
0 2 4 6 8 1012 141618
1000
1200
1400
1600
1800
2000
2200
2400
2600
Flow rate, mL/min
Aging time, hr
ΔP= 40 psi
ΔP= 30 psi
(b)
0 2 4 6 8 1012 141618
0.86
0.88
0.90
0.92
0.94
0.96
0.98
1.00
1.02
1.04
ΔP= 40 psi
ΔP= 30 psi
Separation Factor, N
2
/ CO
2
Aging time, hr
Figure 5.4 The effect of aging time on (a) the flow rate of CO
2
via LDH film on
tubular α-alumina support (b) the separation factor for N
2
/CO
2
with the LDH
membrane prepared by coprecipitation method.
153
80 100 120 140 160 180 200
0.90
0.92
0.94
0.96
0.98
1.00
1.02
1.04
Separation Factor, N
2
/ CO
2
Temperature,
o
C
Figure 5.5 The effect of hydrothermal aging temperature on the separation factor
for N
2
/CO
2
with the LDH membrane prepared by coprecipitation method.
154
(a) (b)
(c)
Figure 5.6 SEM image of (a) cracked surface, and (b) cross section of LDH coated
-alumina support; (c) not covered surface of LDH coated α-alumina support by
oprecipitation method.
α
c
155
156
onsidered to be the result of defects at the LDH membrane, and there were indeed
isible cracks found on the top layer on the inside tube, as shown in Fig. 5.6(a).
ig. 5.6(b), the concentration of the reaction mixture
was diluted 5 times from the original composition, in order to reduce the thickness
of the resulting LDH film. However, even with the diluted concentration,
continuous LDH film were not observed on the top of the inside surface of α-
alumina tubes. Instead, a partially coated α-alumina tube was observed as shown in
Fig. 5.6(c). In order to find an optimal concentration for the preparation of
continuous LDH films, the concentration of reaction mixture was varied. However,
the attempts to prepare defect-free LDH membranes by modifying the reaction
concentration have not been successful, so far.
5.3.2 The sol-gel method
Transport properties of LDH films prepared by the sol-gel method were
also studied with the same procedures utilized with films made by the co-
precipitation method. The permeances of single gases were observed to be about
20% smaller than the values measured with the LDH films prepared by the co-
c
v
Since it was thought that the cracks may be caused by the relatively thick layer (~20
μm) deposited, as shown in F
157
solution was refluxed for 6, 12, and
bserved that the separation factor for CO
2
/N
2
was not
affected
precipitation method, but there were no significant enhancements observed in the
separation factors, as shown in Fig. 5.7. When a 2
nd
LDH layer was coated on the
alumina substrates, the value of the separation factor of CO
2
/N
2
was slightly
decreased, however a significant change was not observed. Upon completion of the
2
nd
LDH layer coating, a 3
rd
LDH layer was coated, and it was observed that the
value of separation factor increased, slowly approaching the Knudsen separation
factor (CO
2
/N
2
=
~0.8). However, when a 4
th
LDH layer was coated on the
membrane, the permeance highly increased, and it was found that the coated LDH
layer cracked.
To improve the separation factor, the effect of aging was investigated. To
study the effect of aging time, three-times coated alumina substrates were placed in
a 2 M Na
2
CO
3
aqueous solution, and then the
18 hr at 100
o
C. It was o
significantly by varying the aging time, as shown in Fig. 5.8(a). The effect
of aging temperature was also investigated with the same procedure as during the
co-precipitation method, and it was observed that, similar with the co-precipitation
method, there was no significant change on the separation factor for CO
2
/N
2
, until
158
0.86
0.90
0.94
0.96
1.00
1.04
1.06
20 30 40 50 60
0.88
0.92
0.98
1.02
1
st
coating
2
nd
coating
3
rd
coating
Sepa factor, CO
2
/ N
2
LDH layer by Dip coating
ration
ΔP
Figure 5.7 Separation factor for CO
2
/N
2
with LDH membrane prepared by sol-gel
method at room temperature.
159
(a)
0 2 4 6 8 1012 14 161820
0.80
0.85
0.90
0.95
1.00
1.05
Aging time, hrs
Separation Factor, CO
2
/ N
2
(b)
0.85
80 100 120 140 160 180 200
0.80
0.90
0.95
1.00
1.05
Separation Factor, CO
2
/ N
2
Temperature,
o
C
Figure 5.8 The effect of (a) aging time, and (b) aging temperature on the separation
factor for CO
2
/N
2
with the LDH membrane prepared by sol-gel method.
160
20 30 40 50 60
0.84
0.86
0.88
0.90
0.92
0.94
0.96
0.98
1.00
1.02
1.04
at room temperature
at 220
o
C
Separation factor, CO
2
/ N
2
ΔP
3rd LDH layer by Dip coating
Figure 5.9 Separation factor for CO
2
/N
2
with 3
rd
layer on tubular α-alumina support
at room temperature and at 220
o
C.
161
120 C, but negative effects were observed when the aging temperature was higher
than 120
o
C, as shown Fig. 5.8(b). The permeance was also measured at 220
o
C
with a three times coated LDH membrane on an α-alumina substrate. At 220
o
C, the
interlayer water molecules are removed without significant collapse of the LDH
structure, as discussed in Chapter 2, and the removal of interlayer water may allow
gases to permeate faster through the LDH layers. It was observed that the separation
factors for CO
2
/N
2
, compared to the values measured at room temperature, were
increased by only ~5% at 220
o
C as shown in Fig. 5.9. However, significant
improvement of separation factor was not observed.
.3.3 Membrane preparation by separate nucleation and aging steps
Transport properties of LDH films prepared by separating the nucleation
nd aging steps were also studied with the same procedures described above. It was
range between 20 ~ 60 psi. When porous stainless steel disks were utilized as
supports, and when their average pore size was less than 0.5 μm, the LDH
crystallites were aggregated and retained on the top of the disk. It was observed that
the films formed by the aggregated LDH crystallites cracked after drying at room
o
5
a
observed that the flux of single gas was small less than 50 mL/min, with ΔP in the
162
er than 120
o
C, as shown Fig.
or CO
2
/N
2
,
temperature. Therefore, to avoid cracks only stainless steel disks with average pore
size bigger than 0.5 μm were utilized. To study the effect of aging time, LDH
coated substrates were placed in a 2 M Na
2
CO
3
aqueous solution, and then the
solution was refluxed for 6, 12, and 18 hr at 100
o
C. It was observed that the
separation factor for CO
2
/N
2
decreased a little, but it was not significantly affected
as shown in Fig. 5.10(a). The effect of aging temperature was also investigated with
the same procedure as described above, and it was observed that there was no
significant change in separation factor for CO
2
/N
2
till 120
o
C, but negative effect
were observed when the aging temperature was high
5.10(b). Although the hydrothermal aging step improved the separation f
nevertheless the values that were observed were very close to the the Knudsen
factor.
5.3.4 Using sulfuric acid as a binder
The fluxes of N
2
and CO
2
through LDH membranes prepared by using
sulfuric acid were very small values, ~ 0.3 mL/min, but the separation factors
(~1.32 in favor of N
2
) were close to the expected Knudsen factor. When the feed
was changed to binary gas mixtures instead of single gas, it was observed that the
(a)
0.90
0.95
0 3 6 9 12 15 18
0.75
0.80
0.85
1.00
1.05
Separa Fact O
2
/ N
Aging Time, hrs
tion or, C
2
LDH layer preparation by separate nucleation and aging
on α -alumina
on porous stainless steel
(b)
0.80
85
0.90
1.00
1.10
163
80 100 120 140 160 180
0.70
0.75
0.
0.95
1.05
1.15
Separatio ctor / N
2
n Fa , CO
2
Temperature, C
LDH layer preparation by separate nucleation and aging
on α-alumina
on porous stainless steel
o
Figure 5.10 The effect of (a) aging time, and (b) aging temperature on the
separation factor for CO
2
/N
2
with the LDH membrane prepared by the separation of
nucleation and aging method.
164
actor was increased to 1.70 in favor of N
2
at room temperature. When
the system temperature was increased to 200
o
C, it was observed that the separation
factor was decreased to 1.55 in favor of N
2
.
The permeance tests were also conducted with binary gas mixtures
containing 5% moisture. It was observed that the separation factor was 1.38 in favor
of CO
2
at room temperature, and at 200
o
C it was 2.29 in favor of CO
2
again. It is
evident that the moisture affects the separation factor favorably, but it is not clear
t the mechanism is that causes this favorable type of behavior.
5.4 Conclusions
Various methods for the preparation of LDH membrane were studied,
including co-precipitation, sol-gel, using separate nucleation and aging steps, and
using sulfate groups as a binder. The transport properties of the prepared
membranes were studied by the measurement of single or binary gas mixtures
separation f
wha
through the LDH layers. Among all prepared LDH membranes, the one prepared
by the use of sulfuric acid as a binder was shown to have the best transport
165
properties. However, the separation factors measured were small, typically close to
the Knudsen factors. The presence of humidity seemed to have a beneficial effect.
166
C
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Abstract (if available)
Abstract
Several in situ techniques have been used to investigate the thermal evolution of the structure of Mg-Al-CO3 layered double hydroxide under inert atmosphere. Based on the results of the study, a model was proposed to describe the structural evolution of the Mg-Al-CO3 LDH. The sorption characteristics and thermal reversibility of Mg-Al-CO3 LDH were also investigated with in situ techniques under both inert and reactive atmospheres. The experimental observations are shown to be consistent with the structural model proposed. The structure, sorption characteristics, and thermal reversibility of Mg-Al-CO3 LDH materials are important in their use for the high temperature applications.
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Asset Metadata
Creator
Kim, Yongman
(author)
Core Title
In situ studies of the thermal evolution of the structure and sorption properties of Mg-Al-CO3 layered double hydroxide
School
Viterbi School of Engineering
Degree
Doctor of Philosophy
Degree Program
Chemical Engineering
Publication Date
02/28/2006
Defense Date
04/29/2005
Publisher
University of Southern California
(original),
University of Southern California. Libraries
(digital)
Tag
hydrotalcite,in situ charactrization,Materials,nano-structure,OAI-PMH Harvest,porous media,sorption reversibility
Language
English
Advisor
Tsotsis, Theodore T. (
committee chair
), Goo, Edward K. (
committee member
), Sahimi, Muhammad (
committee member
)
Creator Email
ymkim2@gmail.com
Permanent Link (DOI)
https://doi.org/10.25549/usctheses-m50
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UC1111941
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etd-Kim-20060928 (filename),usctheses-m40 (legacy collection record id),usctheses-c127-4107 (legacy record id),usctheses-m50 (legacy record id)
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etd-Kim-20060928.pdf
Dmrecord
4107
Document Type
Dissertation
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Kim, Yongman
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texts
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University of Southern California
(contributing entity),
University of Southern California Dissertations and Theses
(collection)
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Libraries, University of Southern California
Repository Location
Los Angeles, California
Repository Email
cisadmin@lib.usc.edu
Tags
hydrotalcite
in situ charactrization
nano-structure
porous media
sorption reversibility