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i
QUEENSLAND UNIVERSITY OF TECHNOLOGY
Inorganic Material Research Group School of Physical and Chemical Sciences
STABILITY OF HYDROTALCITES FORMED FROM BAYER REFINERY ENVIRONMENTAL
CONTROL PROCESSES
by
Sara Jane Palmer
B.A.Sc. (QUT)
This thesis is submitted in fulfilment of the requirements of the degree of Doctor of Philosophy.
June, 2010
i
i
STATEMENT OF ORIGINALITY
The work presented in this thesis has not, to the best of my knowledge, been
previously submitted for a degree or diploma at any other higher education
institution. To the best of my knowledge this thesis contains no material previously
published or written by another person except where due reference is made.
Sara Jane Palmer
17th June 2010
ii
ACKNOWLEDGEMENTS I would like to thank the following people and organisations without who this thesis could not have been completed. They include:
(i) My QUT supervisors: Prof. Ray L. Frost and Dr. Wayde N. Martens, and my QRDC supervisors: Dr. Matthew K. Smith, Mr. John Anderson, Dr. Lyndon Armstrong, and Dr. Steve Healy for providing a challenging research project, guidance, financial support, and assistance in the editing of papers that have been published on this research.
(ii) I would like to extend my further appreciation to Prof. Ray L. Frost and Dr. Lyndon Armstrong for their emotional support and encouragement throughout the past few years.
(iii) Mr. Bill Kwiecien, Ms. Wathsala Kumar, and Mr. Shane Russell for
their advice and technical assistance with the operation of the ICP-OES and analysis preparation.
(iv) Dr. Llew Rintoul for his assistance with the vibrational spectroscopy instruments.
(v) Mr. Anthony Raftery for advice and technical support with the XRD instruments and preparation methods.
(vi) Mr. Lambert Bekessy, Dr. Thor E. Bostrom, and Dr. Loc Duong for their advice and technical support with the operation of the electron microscope.
(vii) The entire Frost group, postgraduate, and staff of the School of Physical
and Chemical Sciences who provided much needed support. Finally, I would like to give special thanks to my family and friends for their love, support and encouragement, especially my parents (Anna M. Palmer and David A. Ashfield) and Mr Marc Couperthwaite.
This work is dedicated to Marc Couperthwaite.
iii
ABSTRACT
Bauxite refinery residues (red mud) are derived from the Bayer process by the
digestion of crushed bauxite in concentrated sodium hydroxide at elevated
temperatures and pressures. This slurry residue, if untreated, is unsuitable for
discharge directly into the environment and is usually stored in tailing dams. The
liquid portion has the potential for discharge, but requires pre-treatment before this
can occur. The seawater neutralisation treatment facilitates a significant reduction
in pH and dissolved metal concentrations, through the precipitation of hydrotalcite-
like compounds and some other Mg, Ca, and Al hydroxide and carbonate minerals.
The hydrotalcite-like compounds, precipitated during seawater neutralisation, also
remove a range of transition metals, oxy-anions and other anionic species through a
combination of intercalation and adsorption reactions: smaller anions are
intercalated into the hydrotalcite matrix, while larger molecules are adsorbed on the
particle surfaces. A phenomenon known as ‘reversion’ can occur if the seawater
neutralisation process is not properly controlled. Reversion causes an increase in the
pH and dissolved impurity levels of the neutralised effluent, rendering it unsuitable
for discharge. It is believed that slow dissolution of components of the red mud
residue and compounds formed during the neutralisation process are responsible for
reversion.
This investigation looked at characterising natural hydrotalcite
(Mg6Al2(OH)16(CO3)·4H2
O) and ‘Bayer’ hydrotalcite (synthesised using the
seawater neutralisation process) using a variety of techniques including X-ray
diffraction, infrared and Raman spectroscopy, and thermogravimetric analysis. This
investigation showed that Bayer hydrotalcite is comprised of a mixture of 3:1 and
4:1 hydrotalcite structures and exhibited similar chemical characteristic to the 4:1
synthetic hydrotalcite. Hydrotalcite formed from the seawater neutralisation of
Bauxite refinery residues has been found not to cause reversion. Other components
in red mud were investigated to determine the cause of reversion and this
investigation found three components that contributed to reversion: 1) tricalcium
aluminate, 2) hydrocalumite and 3) calcium hydroxide. Increasing the amount of
magnesium in the neutralisation process has been found to be successful in reducing
reversion.
iv
LIST OF PAPERS PRODUCED FROM THIS INVESTIGATION
Chapter 1
1.1 Sara J. Palmer, Ray L. Frost, and Tai Nguyen, Hydrotalcites and their role
in coordination of anions in Bayer liquors: Anion binding in layered
double hydroxides. Coordination Chemistry Reviews, 253 (2009) 250-267.
Chapter 3
3.1 Sara J. Palmer, Ray L. Frost, Godwin Ayoko, and Tai Nguyen, Synthesis
and Raman spectroscopic characterisation of hydrotalcite with CO32-
and (MoO4)2-
3.2 Sara J. Palmer, Aurore Soisonard, and Ray L. Frost, Effect of pH on the
uptake of arsenate and vanadate and the stability of these anions in
alkaline solution. Accepted by Journal of Raman Spectroscopy, (2009).
anions in the interlayer. Journal of Raman spectroscopy, 39
(2008) 395-401.
3.3 Veronika Vagvoelgyi, Sara J. Palmer, Janos Kristof, Ray L. Frost, and
Erzsebet Horvath, Mechanism for hydrotalcite decomposition: A
controlled rate thermal analysis study. Journal of Colloid and Interface
Science, 318 (2008) 302-308.
3.4 Sara J. Palmer, Ray L. Frost and Tai Nguyen, Thermal decomposition of
hydrotalcite with molybdate and vanadate anions in the interlayer.
Journal of Thermal Analysis and Calorimetry, 92 (2009) 879-886.
3.5 Sara J. Palmer and Ray L. Frost, Determination of the mechanism(s) for
the inclusion of arsenate, vanadate, or molybdate anions into
hydrotalcites with variable cationic ratio. Journal of Colloid and
Interface Science, 329 (2008) 404-409.
Chapter 4
4.1 Sara J. Palmer and Ray L. Frost, Bayer hydrotalcites formed during the
seawater neutralisation of bauxite refinery residues. Submitted to
Journal of Water Resource and Protection, (2009).
v
4.2 Sara J. Palmer and Ray L. Frost, The effect of synthesis temperature on
the formation of hydrotalcites in Bayer liquor: a vibrational
spectroscopic analysis. Applied Spectroscopy, 63 (2009) 748-752.
4.3 Sara J. Palmer and Ray L. Frost, Thermal decomposition of Bayer
precipitates formed at varying temperatures. Journal of Thermal
Analysis and Calorimetry, 100 (2010) 27-32.
Chapter 5
5.1 Sara J. Palmer, Matthew K. Smith, and Ray L. Frost, The effect of high concentrations of calcium hydroxide in neutralised synthetic supernatant liquor. Submitted to Journal of Industrial and Engineering Chemistry, (2009)
5.2 Sara J. Palmer, Matthew K. Smith, and Ray L. Frost, Minimisation of
reversion using seawater neutralisation and magnesium chloride for
tricalcium aluminate solutions. Submitted to Environment International,
(2009).
Chapter 6
6.1 Sara J. Palmer, Mitchell Nothling, Kathleen H. Bakon, and Ray L. Frost,
Thermally activated seawater neutralised red mud used for the removal
of arsenate, vanadate and molybdate from aqueous solutions. Journal of
Colloid and Interface Science, 342 (2010) 147-154.
vi
Papers presented at conferences:
1. A poster was prepared and presented at the 21st
International Conference on Raman Spectroscopy, London, UK, August 2008, on the synthesis and characterisation of hydrotalcites with vanadate and molybdate. The presentation was based on the following papers:
Sara J. Palmer, Ray L. Frost, Godwin Ayoko, and Tai Nguyen, Synthesis and Raman spectroscopic characterisation of hydrotalcite with CO3
2- and (MoO4)2-
anions in the interlayer. Journal of Raman spectroscopy, 39 (2008) 395-401.
and
Sara J. Palmer, Tai Nguyen, and Ray L. Frost, Synthesis and Raman spectroscopic characterisation of hydrotalcite with CO3
2- and VO3-
anions in the interlayer. Journal of Raman Spectroscopy, 38 (2007) 1602-1608.
2. A poster and a short oral presentation was prepared and presented at the 8th
International Alumina Quality Workshop, Darwin, Australia, September 2008, on the characterisation of bauxite residues before and after seawater neutralisation. The presentation was based on the following papers:
Sara J. Palmer and Ray L. Frost, Characterisation of bauxite and seawater neutralised bauxite residue using XRD and vibrational spectroscopic techniques. Journal of Materials Science, 44 (2009) 55-63
and Sara J. Palmer, B. Jagannadha, and Ray L. Frost, Characterisation of red
mud by UV-vis-NIR spectroscopy. Spectrochimica Acta, Part A: Molecular and Biomolecular Spectroscopy, 71A (2009) 1814-1818.
vii
Papers not presented in this thesis:
A.14 Sara J. Palmer and Ray L. Frost, Characterisation of bauxite and
seawater neutralised bauxite residue using XRD and vibrational
spectroscopic techniques. Journal of Materials Science, 44 (2009) 55-63.
A.15 Sara J. Palmer, B. Jagannadha, and Ray L. Frost, Characterisation of red
mud by UV-vis-NIR spectroscopy. Spectrochimica Acta, Part A:
Molecular and Biomolecular Spectroscopy, 71A (2009) 1814-1818.
A.16 Sara J. Palmer, Tai Nguyen, and Ray L. Frost, Synthesis and Raman
spectroscopic characterisation of hydrotalcite with CO32- and VO3
-
A.17 Sara J. Palmer, Ray L. Frost, and Henry J. Spratt, Synthesis and Raman
spectroscopic study of Mg/Al,Fe hydrotalcites with variable cationic
ratios. Journal of Raman Spectroscopy, 40 (2009) 1138-1143.
anions in the interlayer. Journal of Raman Spectroscopy, 38 (2007) 1602-
1608.
A.18 Sara J. Palmer, Henry J. Spratt, and Ray L. Frost, Infrared and near-
infrared spectroscopic study of synthetic hydrotalcites with variable
divalent/trivalent cationic ratios. Spectrochimica Acta, Part A: Molecular
and Biomolecular Spectroscopy, 72A (2009) 984-988.
A.19 Sara J. Palmer, Henry J. Spratt, and Ray L. Frost, Thermal decomposition
of hydrotalcites with variable cationic ratios. Journal of Thermal Analysis
and Calorimetry, 95 (2009) 123-129.
viii
KEYWORDS
• Arsenate
• Bauxite refinery residue
• Bayer liquor
• Calcium hydroxide
• Hydrocalumite
• Hydrotalcite
• Inductively couple plasma optical emission spectroscopy
• Infrared spectroscopy
• Layered double hydroxides
• Magnesium chloride
• Molybdate
• Raman spectroscopy
• Red mud
• Reversion
• Seawater neutralised
• Thermal activation
• Thermogravimetric analysis
• Tricalcium aluminate hexahydrate
• Vanadate
• X-ray diffraction
ix
LIST OF ABBREVIATIONS
AsO43-
BHT Bayer hydrotalcite
Arsenate
Bppt. Bayer precipitate
CO32-
DTG Differential thermalgravimetric
Carbonate
EDX Energy dispersive X-ray analysis
FTIR Fourier transform infrared spectroscopy
HT Hydrotalcite
ICP-OES Inductively couple plasma optical emission spectroscopy
LDHs Layered double hydroxides
MoO42-
OH
Molybdate -
RM Red mud
Hydroxide ions
RML Red mud liquor
RMS Red mud slurry
SEL Strong evaporation liquor
SEM Scanning electron microscopy
SNL Supernatant liquor
SWN Seawater neutralised
TA Thermally activated
TCA Tricalcium aluminate hexahydrate
TGA Thermogravimetric analysis
VO43-
XRD X-ray diffraction
Vanadate
QUT Queensland University of Technology
QRDC RioTintoAlcan Queensland Research and Development Centre
x
TABLE OF CONTENTS
STATEMENT OF ORIGINALITY … i
ACKNOWLEDGEMENTS … ii
ABSTRACT … iii
LIST OF PAPERS PRODUCED FROM THIS INVESTIGATION … iv
KEYWORDS … viii
LIST OF ABBREVIATIONS … ix
LIST OF FIGURES … xviii
LIST OF TABLES … xxv
CHAPTER 1
Introduction
1. Bauxite refinery residues (red mud) … 2
1.1. Bayer process – Origin of red mud … 2
1.2. Components of red mud … 5
1.2.1. Iron oxides
… 5
1.2.2. Silica minerals
… 8
1.2.3. CaO and Ca(OH)2
1.2.3.1. Causticisation … 13
… 11
1.2.3.2. Tricalcium aluminate hexahydrate (TCA) … 14
1.3. Surface chemistry … 14
1.4. Removal of trace metals from solution … 17
2. Seawater neutralised bauxite refinery residues … 18
2.1. Introduction … 18
2.2. Reaction mechanism … 19
2.3. Formation of hydrotalcite … 20
2.4. Adsorption of anions on the surface of neutralised red mud … 21
3. Layered double hydroxides – LDHs … 23
3.1. Introduction … 23
3.2. Preparation of LDHs … 26
3.3. Anionic exchange … 27
xi
3.4. Thermal activation of hydrotalcite materials … 30
3.5. Characterisation of LDHs … 31
3.5.1. Vibrational spectroscopy – infrared and
Raman spectroscopy
3.5.1.1. Hydroxyl stretching and bending vibrations … 31
… 31
3.5.1.2. Carbonate stretching vibrations … 36
3.5.1.3. Lattice translational modes … 37
3.5.2. Thermal analysis – TGA/DTG
… 39
3.5.3. X-ray diffraction – XRD
3.6. LDHs in the alumina industry … 41
… 40
4. Chapter summary … 43
5. References … 44
CHAPTER 2
Experimental methods and analysis techniques
1. Introduction … 60
2. Experimental methods … 61
2.1. Synthesis of hydrotalcite with different oxy-anions … 61
2.2. Synthesis of Bayer precipitate … 61
2.3. Synthesis of synthetic Bayer precipitate … 63
2.3.1. Synthetic seawater (SW)
… 63
2.3.2. Synthetic supernatant liquor (SNL)
2.4. Seawater neutralisation of red mud … 64
… 63
2.5. Trigger experiments … 65
2.5.1. Trigger materials
2.5.1.1. Synthesis of hydrocalumite –
… 65
Ca2Al(OH)6Cl·2H2
2.5.1.2. Synthesis of whewellite – CaC
O … 65
2O4·H2
2.6. Thermal activation and treatment of aqueous solutions … 67
O … 67
3. Characterisation … 69
3.1. Inductively coupled plasma optical emission spectrometry … 69
3.2. X-ray diffraction … 70
xii
3.3. Spectroscopy … 70
3.3.1. Fourier-transform infrared spectroscopy
… 70
3.3.2. Fourier-transform Raman spectroscopy
… 70
3.3.3. Raman microspectroscopy
… 70
3.3.4. Band component analysis
3.4. Thermal analysis … 71
… 71
3.4.1. Thermogravimetric analysis
… 71
3.4.2. Dynamic experiment
… 72
3.4.3. Controlled rate thermal analysis experiment
3.5. Electron dispersive X-ray spectroscopy … 72
… 72
3.6. Potentiometric titration … 73
4. References … 73
CHAPTER 3
Synthesis and characterisation of synthetic hydrotalcites
1. Introduction … 75
2. Infrared and Raman spectroscopy … 77
2.1. Hydroxyl stretching region … 77
2.2. Carbonate vibrations … 79
2.3. Water OH deformation vibrations … 81
2.4. Vibrations associated with arsenate … 83
2.5. Vibrations associated with vanadate … 85
2.6. Vibrations associated with molybdate … 87
2.7. Cation deformation vibrations … 87
3. Effect of pH and Mg,Al hydrotalcite ratios on the removal of
oxy-anions from aqueous solutions … 89
3.1. Effect of synthesis pH … 89
3.2. Chemical stability of hydrotalcites synthesised over a 2, 24, and
48 hour period … 91
3.2.1. pH 10 … 93
xiii
3.2.2. pH 14
3.3. Raman spectra of hydrotalcite synthesised … 95
… 95
3.3.1. Carbonate vibrational region (1200-600 cm-1
) … 95
4. X-ray diffraction - XRD … 101
5. Controlled rate thermal analysis of carbonate hydrotalcite … 103
6. Thermal analysis and mass spectroscopy – TGA/DTG and MS … 106
6.1. Effect of different oxy-anions on the thermal analysis patterns
of 3:1 hydrotalcite … 106
6.1.1. Carbonate hydrotalcite HT(CO32-
) … 109
6.1.2. Carbonate and arsenate hydrotalcite
HT(CO32-, AsO4
3-) … 109
6.1.3. Arsenate hydrotalcite HT(AsO43-) … 111
6.1.4. Carbonate and vanadate hydrotalcite
HT(CO32-, VO4
3-) … 113
6.1.5. Vanadate hydrotalcite HT(VO43-) … 113
6.1.6. Carbonate and molybdate hydrotalcite
HT(CO32-, MoO4
2-) … 117
6.1.7. Molybdate hydrotalcite HT( MoO42-
) … 117
7. Mechanism of anion inclusion (intercalation and/or adsorption) … 119
7.1. Effect of cationic ratio on the thermal stability of hydrotalcites
with different interlayer anions … 119
7.1.1. Arsenate hydrotalcites
… 119
7.1.2. Vanadate hydrotalcites
… 123
7.1.3. Molybdate hydrotalcites
… 123
8. Chapter summary … 125
9. References … 127
xiv
CHAPTER 4
Synthesis and characterisation of Bayer hydrotalcites
1. Introduction … 129
2. Identification of hydrotalcite formation in seawater neutralised
red mud … 131
2.1. X-ray diffraction … 131
2.2. Thermal analysis … 133
3. Bayer hydrotalcites formed during the seawater neutralisation of
bauxite refinery residues … 135
3.1. X-ray diffraction … 135
3.2. EDX analysis … 136
3.3. ICP-OES analysis … 136
3.4. Raman and infrared spectroscopy … 139
3.5. Thermogravimetric Analysis … 143
4. The effect of synthesis temperature on the formation of
hydrotalcites in Bayer liquor … 145
4.1. X-Ray Diffraction … 145
4.2. Vibrational spectroscopy … 149
4.2.1. Hydroxyl stretching and bending vibrations
… 149
4.2.2. Carbonate vibrational region
… 153
4.2.3. Cation OH deformation modes
4.3. Thermal analysis –TG and DTG … 157
… 155
4.3.1. Decomposition between 30 – 230 °C
… 157
4.3.2. Decomposition between 250 – 400 °C
… 159
4.3.3. Decomposition between 400 – 650 °C
… 161
5. Chapter summary … 163
6. References … 165
xv
CHAPTER 5
Reversion
1. Introduction … 168
1.1. pH reversion … 169
1.1.1. Effect of volumetric seawater to RMS ratio
… 171
1.1.2. Effect of temperature
1.2. Reversion of dissolved metals … 175
… 171
1.3. Identification of the source of reversion … 175
1.4. Identification of triggers causing reversion … 177
2. Triggers causing reversion … 177
2.1. Seawater neutralised SNL - blank … 177
2.1.1. pH
… 177
2.1.2. ICP-OES
2.2. Synthetic SNL with calcium hydroxide (Ca(OH)
… 179
2
) … 183
2.2.1. pH
… 183
2.2.2. ICP-OES
… 189
2.2.3. XRD
… 189
2.2.4. TGA
2.3. Synthetic SNL with hydrocalumite (Ca
… 191
2Al(OH)6.Cl·2H2
O) … 193
2.3.1. pH
… 193
2.3.2. ICP-OES
2.4. Synthetic SNL with tricalcium aluminate hexahydrate … 197
… 195
2.4.1. pH
… 197
2.4.2. ICP-OES
… 199
2.4.3. Mechanism for TCA reversion
… 199
3. Triggers NOT causing reversion … 203
3.1. Synthetic SNL with Bayer precipitate … 203
3.1.1. pH
… 203
3.1.2. ICP-OES
3.2. Synthetic SNL with whewellite (CaC
… 205
2O4·H2
O) … 205
3.2.1. pH
… 205
3.2.2. ICP-OES … 207
xvi
3.3. Synthetic SNL with sodalite (Na8(AlSiO4)6
Cl) … 207
3.3.1. pH
… 207
3.3.2. ICP-OES
… 207
3.3.3. EDX
3.4. Synthetic SNL with Na
… 207
2CO3
… 210
3.4.1. pH
… 210
4. Minimising reversion … 211
4.1. Neutralisation ratio … 211
4.2. Addition of MgCl2·6H2
4.3. Addition of MgCl
O to synthetic supernatant liquor … 213
2·6H2
4.4. Confirmation of hydrotalcite formation … 217
O to red mud slurry … 215
5. Chapter summary … 218
6. References … 221
CHAPTER 6
Thermally activated seawater neutralised red mud used for the
removal of arsenate, vanadate and molybdate from aqueous
solutions
1. Introduction … 223
2. Effect of Mg:Al cationic ratio on anion removal for mixed
anion solutions … 225
2.1. 2:1 synthetic hydrotalcite … 228
2.2. 3:1 synthetic hydrotalcite … 228
2.3. 4:1 synthetic hydrotalcite … 229
2.4. Bayer hydrotalcite … 229
3. Red mud and seawater neutralised red mud … 231
4. Chapter summary … 232
5. References … 233
xvii
CHAPTER 7
Conclusions and recommendations for future work
1. Conclusions … 235
2. Recommendations … 239
APPENDIX
A.1 Calculation of water in the carbonate hydrotalcite – Chapter 3 … 241
xviii
LIST OF FIGURES
CHAPTER 1 Figure 1.1: Solubilities of goethite and hematite as a function of pH. [24]
Figure 1.2: Singly, doubly, triply coordinated and germinal surface hydroxyl
groups on iron oxides. [24]
Figure 1.3: Modes of ligand coordination to the iron oxide surface. [24]
Figure 1.4: The SiO2
Figure 1.5: Aluminosilicate solubility in a synthetic Bayer solution as a
function of Na
equilibrium solubility of sodalite and cancrinite formed
under different conditions. [59]
2CO3
Figure 1.6: CaO-Na
concentration at 90 ºC. [66]
2O-CO2-Al2O3-H2
Figure 1.7: Titration curves of red mud slurry (dotted line) and caustic solution
(solid line). [3]
O phase diagram. [23]
Figure 1.8: Schematic representation of the hydroxide layers in the
hydrotalcite.
Figure 1.9: Schematic representation of the hydrotalcite structure.
Figure 1.10: Water, hydroxyl and carbonate vibrations in the interlayer of Co
and Ni hydrotalcites. [204]
Figure 1.11: Infrared bands of adsorbed CO2
surface species on calcined
hydrotalcite. [220]
CHAPTER 2 Figure 2.1: XRD pattern of synthesised hydrocalumite and the corresponding
reference patterns.
Figure 2.2: XRD pattern of synthesised whewellite and the corresponding
reference pattern.
xix
CHAPTER 3 Figure 3.1: Raman and infrared spectra of carbonate hydrotalcite in the
hydroxyl stretching vibrational region.
Figure 3.2: Infrared spectra of the synthesised hydrotalcites, containing
arsenate, in the carbonate vibrational region.
Figure 3.3: Infrared spectra of the synthesised hydrotalcites, containing
vanadate, in the carbonate vibrational region.
Figure 3.4: Infrared spectra of synthesised hydrotalcites, containing
molybdate, in the carbonate vibrational region.
Figure 3.5: Raman spectra of the synthesised hydrotalcites, with arsenate, in
the carbonate vibrational region.
Figure 3.6: Raman spectra of the synthesised hydrotalcites, with vanadate, in
the carbonate vibrational region.
Figure 3.7: Raman spectra of synthesised hydrotalcites, containing molybdate,
in the hydroxyl stretching region.
Figure 3.8: Raman spectra of the cation deformation modes of arsenate
containing hydrotalcites.
Figure 3.9: Raman spectra of the cation deformation modes of vanadate
containing hydrotalcites.
Figure 3.10: Raman spectra of the cation deformation modes of molybdate
containing hydrotalcites.
Figure 3.11: Percentage of anions removed from solution during the synthesis
process for different hydrotalcites at varying reaction pH.
Figure 3.12: Molecular shape of the vanadate anion in the pH range 7-14.
Figure 3.13: Molecular shape of the arsenate anion in the pH range 7-14.
Figure 3.14: Raman spectrum in the anionic stretching region, 1200-600 cm-1
Figure 3.15: Raman spectrum in the anionic stretching region, 1200-600 cm
,
for hydrotalcites prepared for 2 hours at pH 8. -1
Figure 3.16: Raman spectrum in the anionic stretching region, 1200-600 cm
,
for hydrotalcites prepared for 48 hours at pH 8. -1
Figure 3.17: XRD patterns and references for the synthesised hydrotalcites with
molybdate and vanadate in the interlayer.
,
for hydrotalcites prepared for 2 hours at pH 13.
xx
Figure 3.18: The dynamic thermogravimetric and differential
thermogravimetric analysis of carbonate intercalated Mg-Al
hydrotalcite.
Figure 3.19: The controlled rate thermal analysis of carbonate intercalated Mg-
Al hydrotalcite.
Figure 3.20: The thermogravimetric and differential thermogravimetric analysis
of HT(CO32-
Figure 3.21: The ion current curves for selected evolved gases in the thermal
decomposition of HT(CO
).
32-
Figure 3.22: The thermogravimetric and differential thermogravimetric analysis
of HT(CO
).
32-,AsO4
3-
Figure 3.23: The ion current curves for selected evolved gases in the thermal
decomposition of HT(CO
).
32-,AsO4
3-
Figure 3.24: The thermogravimetric and differential thermogravimetric analysis
of HT(AsO
).
43-
Figure 3.25: The ion current curves for selected evolved gases in the thermal
decomposition of HT(AsO
).
43-
Figure 3.26: The thermogravimetric and differential thermogravimetric analysis
of HT(CO
).
3,VO43-
Figure 3.27: The ion current curves for selected evolved gases in the thermal
decomposition of HT(CO
).
3,VO43-
Figure 3.28: The thermogravimetric and differential thermogravimetric analysis
of HT(VO
)
43-
Figure 3.29: The ion current curves for selected evolved gases in the thermal
decomposition of HT(VO
).
43-
Figure 3.30: The thermogravimetric and differential thermogravimetric analysis
of HT(CO
).
3,MoO42-
Figure 3.31: The ion current curves for selected evolved gases in the thermal
decomposition of HT(CO
).
3,MoO42-
Figure 3.32: The thermogravimetric and differential thermogravimetric analysis
of HT(MoO
).
42-
Figure 3.33: The ion current curves for selected evolved gases in the thermal
decomposition of HT(MoO
).
42-
Figure 3.34: Raman spectra of the synthesised hydrotalcites with variable
cationic ratio.
).
xxi
Figure 3.35: DTG curves of the synthesised hydrotalcites with variable cationic
ratios.
CHAPTER 4 Figure 4.1: Comparison of red mud and seawater neutralised red mud XRD
patterns.
Figure 4.2: Thermal analysis of an Australian red mud.
Figure 4.3: Thermal analysis of seawater neutralised red mud.
Figure 4.4: XRD pattern of precipitate formed during the SWN of Bayer
liquor.
Figure 4.5: Infrared and Raman spectra of the Bayer precipitate in the
hydroxyl stretching region.
Figure 4.6: Infrared spectrum of Bayer hydrotalcite in the carbonate
vibrational region.
Figure 4.7: Raman spectrum of Bayer precipitate in the 1150 to 950 cm-1
Figure 4.8: Raman spectrum of Bayer precipitate in the 800 to 200 cm
region. -1
Figure 4.9: DTG curves of Bayer precipitate, hydrotalcite, calcium carbonate,
and seawater.
region.
Figure 4.10: TG/DTG curve of the Bayer precipitate.
Figure 4.11: XRD patterns of Bayer precipitates synthesised at different
temperatures via the SWN process.
Figure 4.12: Raman and infrared spectra of Bayer precipitates in the
hydroxyl stretching region.
Figure 4.13: Infrared spectra of Bayer precipitates in the 1800 - 1200 cm-1
Figure 4.14: Raman spectra of Bayer precipitates in the 1200 - 900 cm
region. -1
Figure 4.15: Raman spectra of Bayer precipitates in the 900 - 200 cm
region. -1
Figure 4.16: Thermal analysis of Bayer precipitates formed at 0, 25, 55, and
75 °C.
region.
Figure 4.17: Stacked DTG curves of the Bayer precipitates in the
dehydroxylation/decarbonation region.
xxii
CHAPTER 5 Figure 5.1: Seawater neutralisation curve of a red mud slurry obtained from a
Gove refinery in 2008.
Figure 5.2: Effect of the volumetric seawater neutralisation ratio on pH
reversion.
Figure 5.3: Effect of temperature on pH reversion.
Figure 5.4: pH plot for the SWN of synthetic SW and SNL.
Figure 5.5: Aluminium concentration after neutralisation, determined by
ICP-OES.
Figure 5.6: Magnesium concentration after neutralisation, determined by
ICP-OES.
Figure 5.7: Calcium concentration after neutralisation, determined by
ICP-OES.
Figure 5.8: TG analysis of Bayer precipitate.
Figure 5.9: Sulfate concentration after neutralisation, determined by
ICP-OES.
Figure 5.10: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of Ca(OH)2
Figure 5.11: Concentration of magnesium cations in solution for varying
concentrations of Ca(OH)
.
2
Figure 5.12: XRD patterns of calcium aluminate species tested as triggers and
the corresponding reference patterns.
in SWN-SNL over 2 hours.
Figure 5.13: Concentration of calcium cations in solution for varying
concentrations of Ca(OH)2
Figure 5.14: DTG curves of 1.00M Ca(OH)
in SWN-SNL over 2 hours.
2
Figure 5.15: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of hydrocalumite.
before and after SWN.
Figure 5.16: Aluminium concentration in solution after the SWN of SNL with
varying concentrations of hydrocalumite.
Figure 5.17: Magnesium concentration in solution after the SWN of SNL with
varying concentrations of hydrocalumite.
Figure 5.18: Calcium concentration in solution after the SWN of SNL with
varying concentrations of hydrocalumite.
Figure 5.19: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of TCA.
xxiii
Figure 5.20: Concentration of aluminium in solution after the SWN process,
using ICP-OES.
Figure 5.21: Concentration of magnesium in solution after the SWN process,
using ICP-OES.
Figure 5.22: Flow chart of the reactions involved in the dissolution of TCA.
Figure 5.23: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of Bppt.
Figure 5.24: Aluminium concentration in solution after the SWN of SNL with
varying concentrations of Bppt.
Figure 5.25: Magnesium concentration in solution after the SWN of SNL with
varying concentrations of Bppt.
Figure 5.26: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of whewellite.
Figure 5.27: Aluminium concentration in solution after the SWN of SNL with
varying concentrations of whewellite.
Figure 5.28: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of sodalite.
Figure 5.29: Aluminium concentration in solution after the SWN of SNL with
varying concentrations of sodalite.
Figure 5.30: Combined pH plots for the SWN of synthetic SW and SNL with
varying concentrations of Na2CO3
Figure 5.31: pH curves for the addition of MgCl
.
2·6H2
Figure 5.32: Aluminium concentration for SWN synthetic SNL containing
0.10M TCA with additional MgCl
O to seawater and the
SWN of synthetic SNL containing 0.10M TCA.
2·6H2
Figure 5.33: Magnesium concentration for SWN synthetic SNL containing
0.10M TCA with additional MgCl
O added to seawater.
2·6H2
Figure 5.34: pH curves for the addition of MgCl
O added to seawater.
2·6H2
Figure 5.35: Thermal analysis of seawater neutralised red mud slurry with an
additional 1000 ppm of magnesium chloride.
O to seawater and the
SWN of red mud slurry.
xxiv
CHAPTER 6 Figure 6.1: Mixed solution removal capacity of thermally activated 2:1
synthetic hydrotalcite.
Figure 6.2: Mixed solution removal capacity of thermally activated 3:1
synthetic hydrotalcite.
Figure 6.3: Mixed solution removal capacity of thermally activated 4:1
synthetic hydrotalcite.
Figure 6.4: Mixed solution removal capacity of thermally activated Bayer
hydrotalcite.
Figure 6.5: Comparison of the removal abilities of thermally activated red mud
and seawater neutralised red mud for the removal of arsenate,
vanadate, and molybdate.
xxv
LIST OF TABLES
CHAPTER 1 Table 1.1: Mineralogy of some typical bauxites [19-23].
Table 1.2: Compositions, crystallographic parameters and symmetries for
some natural LDHs.
Table 1.3: Wavenumber (cm-1) and assignments of the hydroxide layer modes
of the types M-OH and M-O in the infrared spectra of
Mg,Al-layered double hydroxides in comparison to brucite
Mg(OH)2
Table 1.4: Infrared water bending vibrational positions of Mg,Al
hydrotalcites as a function of the interlayer anion, as reported in
literature. [136]
.
Table 1.5: Wavenumber (cm-1
Table 1.6: FT-IR interlayer carbonate vibrational modes. [209,211].
) and assignments of the hydroxide layer modes
of the type M-OH and M-OH in the Raman spectrum of
Mg,Al-layered double hydroxides.
CHAPTER 2 Table 2.1: Concentrations of Na2CO3, Na2HAsO4·7H2O, NaVO3, and
Na2MoO4
Table 2.2: Concentration and masses used to synthesis 2:1, 3:1, and 4:1
synthetic hydrotalcites.
used to synthesise hydrotalcites with different oxy-
anions.
Table 2.3: Composition of Bayer liquors, determined by Potentiometric
titration.
Table 2.4: Salts used to prepare synthetic seawater and relative
concentrations.
Table 2.5: Alumina, caustic and carbonate concentration of synthetic SNL
and real SNL, determined by Potentiometric titration.
Table 2.6: Concentration and mass of each trigger in 60 mL of synthetic SNL.
xxvi
CHAPTER 3 Table 3.1: CO3
2-
Table 3.2:
bands. [10]
VO43- and AsO4
3-
Table 3.3: Fifteen 3:1 hydrotalcites prepared at pH 8 and aged for 2, 24, and
48 hours.
bands from different sources. [10]
Table 3.4: Percentage dissolution of hydrotalcites formed over varying
synthesis periods in NaOH at pH 10 and pH 14. Note the results
for the mixed anion hydrotalcite are for the anion in bold.
Table 3.5: Thermal decomposition of carbonate intercalated hydrotalcite
under dynamic conditions.
Table 3.6: Decomposition stages under CRTA conditions.
Table 3.7: Summary of the TG analysis spectrum of the synthesised
hydrotalcites.
CHAPTER 4 Table 4.1: Quantitative XRD analysis of red mud.
Table 4.2: EDX analysis of the molar ratio of the three Bayer precipitates.
Table 4.3: Percentage removals of ions during the SWN of Bayer liquors.
Table 4.4: EDX results of the molar ratio of Bayer precipitates synthesised
at 0, 25, 55, and 75 °C.
CHAPTER 5 Table 5.1: Summary of pH during the SWN-RMS at 5, 25, 55, and 75 °C.
Table 5.2: Percentage increase of aluminium, arsenate, vanadate, and
molybdate 60 minutes after neutralisation.
Table 5.3: Initial and final pH of solution and the percentage increased over a
2 hour period.
Table 5.4: Concentration of Ca(OH)2 in g/L and the concentration of solid
Ca(OH)2
Table 5.5: Summary of pH results for hydrocalumite concentrations that
showed pH reversion.
left in SNL and SWN-SNL, if no reactions with the
dissolution products occur.
Table 5.6: Comparison of pH and the concentration of whewellite in SNL.
Table 5.7: Elemental ratio of sodalite scale.
1
CHAPTER 1
INTRODUCTION:
- Bauxite refinery residue
- Seawater neutralised bauxite refinery residue
- Hydrotalcite
2
1. Bauxite refinery residues (red mud)
1.1. Bayer process – Origin of red mud
Bauxite refinery residues are derived from the Bayer process, summarised in
Eq. 1, 2, and 3, [1] by the digestion of crushed bauxite in concentrated caustic
(NaOH) at elevated temperatures. Digestion temperatures are dependent on the
quantity of gibbsite (γ - Al(OH)3), boehmite (γ - Al(O)OH), and diaspore
(α - Al(O)OH) present in the bauxite ore. Bauxites containing predominantly
gibbsite require lower digestion temperatures (145 - 175 ºC), while those with
high boehmite and diaspore require stronger caustic concentrations and
temperatures (245 - 275 ºC). [2] The process results in the dissolution of gibbsite
(Al(OH)3
) and boehmite as sodium aluminate, while the remaining insoluble
residue (45% liquor and 55% solid mud), known widely as red mud, is removed
by means of flocculation and decantation. [1, 3] The exact composition of the fine
textured residue depends on the initial type of bauxite. [4] Roughly 1.0 to 1.5
tonnes of red mud residue is produced for every tonne of alumina produced, [5]
therefore millions of tonnes of red mud is produced annually. The liquor is
strongly alkaline (pH ranging from 10 to 13) [6-8] and requires neutralisation to a
pH below 9, with an optimum pH value of 8.5 to 8.9, [9, 10] thus reducing the
potential for environment impact. The liquor also contains relatively high
concentrations of aluminium and a variety of anionic species including oxy-anions
of transition metals. Many of these species can be detrimental to the environment
and therefore must be removed prior to disposal. Bayer process red mud and their
environmental applications have received substantial research. [11-15]
1. Extraction:
Al(OH)3(s) + NaOH (aq) → Na+ Al(OH)4-(aq)
AlO(OH)
(Gibbsitic bauxite)
(s) + NaOH(aq) + H2O → Na+ Al(OH)4-(aq)
Insoluble residue is removed
(Boehmitic
bauxite)
2. Precipitation:
Na+ Al(OH)4-(aq) → Al(OH)3(s) + NaOH
3. Calcination: (aq)
3
2Al(OH)3(s) → Al2O3(s) + 3H2O
Red mud varies in physical, chemical, and mineralogical properties due to
differing bauxite ores and refining processes. [16-18] Table 1.1 demonstrates the
variability of bauxites mined in different locations. Generally, red mud is
composed of iron oxides, primarily hematite (Fe
(g)
2O3), and goethite (FeOOH),
boehmite (AlOOH), other aluminium hydroxides, calcium oxides, titanium oxides
(anatase and rutile), and aluminosilicate minerals (sodalite). [3, 10, 16, 18]
Charged lime species may also be present in the form of calcium carbonate
(CaCO3), 3CaO·Al2O3·6H2O, various forms of calcium phosphate (carbonate or
hydroxyapatite), as well as the formation of perovskite (CaTiO3) and/or kassite
(CaTi2O4(OH)2
) at high bauxite digestion temperatures. [3] These minerals are
the chemically stable end products of bauxite formation and refining, and are the
components responsible for the high surface reactivity of red mud. [1, 3, 9, 10, 16]
Table 1.1: Mineralogy of some typical bauxites. [19-23]
Minerals Weipa Darling Range India Greece
[23] [22] [21] [19]
Gibbsite, Al(OH) 58.3 3 51.1 59.2 0
Boehmite, AlO(OH) 12.5 0.4 7.8 3.0
Diaspore, AlO(OH) 0.2 0.5 1.2 60.0
Kaolin, A12O3.2SiO2·2H2 10.3 O 6.5 5.6 3.0
Quartz, SiO Trace 2 17.4 1.4 0
Hematite, Fe2O 10.6 3 7.2 10.7 21.0
Goethite, FeO(OH) 3.9 9.5 6.2 4.0
Anatase, TiO 2.0 2 1.0 6.0 3.0
Rutile, TiO 0.7 2 0 0.5 0
P2O 0.1 5 NR NR NR
CaO 0.1 0 NR 0.7
98.7 93.7 98.6 94.7
NR = not reported.
4
Figure 1.1: Solubilities of goethite and hematite as a function of pH. [24]
5
1.2. Components of red mud
1.2.1. Iron oxides
In general, the solubility of FeIII oxides is low, while FeII oxides are sparingly
soluble. In the pH range 4 to 10 the level of total Fe in solution is < 10-6
M. [24]
Iron oxides dissolve slowly over a wide pH range. The solubility diagram
(Fig. 1.1) of hematite and goethite indicates that the iron oxides appear to have
minimum solubility around pH 7-8, which is around the point of zero charge
(PZC). As iron oxides are amphoteric, they dissolve in acid media to form
cationic hydroxo species and in basic media to form anionic hydroxo species. [24]
The solubility of the iron oxides rises at pH values greater and lower than 7-8.
The particle size of the solid will affect solubility, where crystals < 1μm may
increase solubility due to the high surface area. This occurs because of surface
properties, especially the surface free energy, rather than the properties of the bulk
solution, that govern the dissolution behaviour. Surface free energies of iron
oxides are relatively high, therefore particle sizes will have a noticeable effect on
the solubility of the compound.
The surface hydroxyl groups (whether they arise from the adsorption of water or
from structural OH) are the chemically reactive entities at the surface of the solid
in an aqueous environment. They possess a double pair of electrons together with
a dissociable hydrogen atom which enables them to react with both acids and
bases, therefore making iron oxides amphoteric (Eq. 4 and 5).
4. ≡ FeOH2+ FeOH + H+
5. ≡
where ≡ denotes the surface
FeOH FeO- + H
+
The surface groups can be replaced by silane groups, [25] or by titanate groups,
[26] (Eq. 6).
6. ROTi (-OR’)3 + ≡ OH → ≡ OTi (OR’)3
where R and R’ are alkyl groups and ≡ represents the oxide surface
+ ROH
6
Figure 1.2: Singly, doubly, triply coordinated and geminal surface
hydroxyl groups on iron oxides. [24]
Figure 1.3: Modes of ligand coordination to the iron oxide surface. [24]
7
Crystallographic considerations indicate that the surface hydroxyl groups may be
coordinated to one (singly), two (doubly), or three (triply) underlying Fe atoms
(Fig. 1.2). [24] The overall density of these groups depends on both the crystal
structure and the extent of development of different crystal faces. Therefore, the
density of the hydroxyl groups depends on the oxide and its crystal morphology.
The most reactive groups are singly coordinated, with total hydroxyl densities
between 8 and 16OH nm-2. [27] Due to the differences in the number of
underlying Fe atoms that are coordinated to the surface functional groups, the
acidity and hence, the reactivity of the different types of hydroxyl groups should
vary. Adsorption studies appear to indicate doubly coordinated surface hydroxyls
on goethite and hematite are inert over a wide pH range. [28-31] Adsorption of
ions on iron oxides is considered to involve only singly coordinated surface
groups. The density of surface functional groups on various iron oxide has been
measured by such techniques as acid/base titration, [32, 33] BET treatment of
water vapour isotherms, [34] D2
O or titanium exchange, [35] and by reactions
with the adsorbing species such as fluoride, phosphate or oxalate. [33, 36]
The adsorption process involves the interaction of the adsorbing species, the
adsorbate, with the surface hydroxyl groups on the iron oxide, the absorbent. The
oxygen donor atom of the surface hydroxyl group can interact with protons,
whereas the underlying metal ion acts as a Lewis acid and exchanges the OH
group for other ligands to form surface complexes. Adsorption of simple
inorganic anions, oxy-anions and organic ions on iron oxides has been widely
investigated. [37-45] Anions are ligands, i.e. they possess one or more atoms with
a lone pair of electrons and can therefore function as the donor in a coordinate
bond. Adsorption of anions on iron oxides can occur either specifically or non-
specifically. Specific adsorption involves the replacement of the surface hydroxyl
group by the adsorbing ligand, L (Eq. 7 and 8). It involves the direct coordination
of the adsorbing species to the surface metal atom of the solid (Fig. 1.3). It is also
termed chemisorption, inner sphere adsorption, and in the case of ligands, ligand
exchange. Specifically adsorbing ions modify the surface charge on the oxide and
hence, cause a shift in the PZC (discussed in surface chemistry of red mud). They
are usually tightly bound and are not easily displaced. Anions that adsorb
specifically on iron oxides include phosphate, silicate, selenate, arsenate, chloride,
fluoride, citrate, and oxalate. [24]
8
7. ≡ FeOH + L- → FeL + OH-
8. ≡
where ≡ denotes the surface
(FeOH)2 + L- → Fe2L+ + 2OH
-
Anion adsorption at any pH increases with increasing concentration of the
adsorbing species. Adsorption is at a maximum at low pH and decreases with
increasing pH except for silicate. [46] The decrease in adsorption at increasing pH
is a result of a decrease in the number of FeOH2+
groups present.
Adsorption of cations on iron oxides are also specific and non specific, where the
trivalent cations (Al3+ 9) appear to adsorb as surface hydroxo species (Eq. ).
9. ≡ FeOH + Al3+ + H2O ≡ Fe-O-AlOH+ + 2H+
Cation adsorption on iron oxides is initially rapid, but adsorption of trace metals
can continue to increase over days with long reaction times being needed to reach
equilibrium. Adsorption of Ni, Zn, and Cd on goethite rose as the reaction time
was extended from 2 hr to 42 days. [47] Adsorption of aluminium on iron oxides
is of interest due to the environmental effect of high levels of Al in fresh water
and soils. Adsorption on goethite takes place over the pH range 3 to 8.5. [48, 49]
The data fit the adsorption model with two monodentate surface hydroxo-
complexes (Fe-OAlOH+ and FeOAl(OH)2
) which formed successively as the pH
increased. Desorption of aluminium from goethite is extremely slow and
adsorption is only partly reversible.
1.2.2. Silica minerals
The main impurities in bauxites are compounds of silicon, iron, and titanium.
Silica is present as kaolinite (Al2O3·2SiO2·2H2O) and halloysite
(Al2O3·2SiO2·3H2O). [2] Silica, in the form of quartz, is not perceptibly attacked
during low temperature digestion in the Bayer process, but silica contained in clay
(reactive silica) readily dissolves in caustic soda. However, quartz can be attacked
during high temperature digestion. Dissolved silica then reacts with hydroxide and
alumina and rapidly re-precipitates as sodium aluminosilicate (DSP) or sodalite
(Eq. 10). [23, 50, 51] The vast majority of DSP is discarded with red mud, but
9
some inadvertently remains dissolved in solution, and this is the primary source of
scale throughout all alumina refineries. [52] The exact nature and location of scale
varies widely,[53] but its control and removal is one of the primary maintenance
costs for all major refineries. Bayer sodalite has the general composition:
(3(Na2O·Al2O3·2SiO2·nH2O)·Na2X) where n ranges from 0 to 2 and X represents
CO32-, SO4
2-, 2OH-, 2Cl-
, or a mixture of all, depending on liquor impurities. [3,
16] The process of formation is described by Eq. 11.
10. Al2O3·2SiO2 + NaOH → Na2SiO3
11. 6SiO
32- + 6Al(OH)4
- + 6Na+
→ Na
+ 2NaX
8(AlSiO4)6X2·nH2O(s) + (6-n)H2O + 12OH-
where X can be ½CO
(sodium aluminosilicate)
32-, ½SO4
2-, 2OH-, 2Cl
-
Cancrinite and sodalite are common sodium aluminosilicate compounds that form
in strongly caustic alkaline aqueous solutions. Cancrinite is defined as belonging
to the hexagonal crystal system with ABAB layer type packing. Large channels as
well as a series of smaller cages run parallel to the z-axis. [54] Cancrinite has
characteristic infrared bands at 1095, 1035, and 1000 cm-1 attributed to the
antisymmetric stretch, ν(Al-O-Si) of the aluminosilicate framework, [55] and at
690, 630, and 560 cm-1 for the symmetric stretch of the aluminosilicate
framework. [56] Sodalite has a general cubic crystal system with ABC layer
packing creating a network of large cages rather than channels, like that of
cancrinite. Sodalite has characteristic infrared spectra at 1000 cm-1 attributed to
the antisymmetric stretch of the Al-O-Si framework, with symmetric vibrational
bands located at 737, 713, and 668 cm-1
. [56]
Silica solubility has been shown to increase with increasing sodium hydroxide and
alumina concentrations. [57] There is some disagreement as to the effect of
temperature on silica solubility. Ostap, [57] Breuer et al., [58], and Barnes et al.,
[59], reported increasing solubility with increasing temperature for Bayer
solutions. Barnes et al, [59], investigated the solubility of sodalite and cancrinite
(expressed in terms of SiO2 concentration) with increased temperature (Fig. 1.4).
10
Figure 1.4: The SiO2
cancrinite formed under different conditions. [59]
equilibrium solubility of sodalite and
Figure 1.5: Aluminosilicate solubility in a synthetic Bayer
solution as a function of Na2CO3 concentration at 90 ºC. [66]
11
Sodalite had increased equilibrium SiO2 solubility compared to cancrinite at all
temperatures. [59]. However, Oku and Yamada, [60], reported no temperature
dependence, up to 150 ºC, in the desilication rate of Bayer liquors. Ni et al., [61],
have measured the solubility of sodium aluminosilicate solutions, where no
sodium carbonate (Na2CO3) has been added, and found that sodium
aluminosilicate solubility increases with increasing alkali concentration (Fig. 1.5).
The presence of Na2CO3
in solution decreases the solubility of both cancrinite
and sodalite.
Studies have shown that sodalite transforms to cancrinite
Na6Ca1.5Al6Si6O24(CO3)1.6) over time, [60, 62] however variations exist in the
literature in regards to the rate of transformation. The mechanisms and kinetics of
sodalite and cancrinite formation have been reported in literature. [63-65] The
presence of Na2CO3
in synthetic liquor causes a decrease in the rate of
transformation of sodalite to cancrinite. [63-65]
1.2.3. CaO and Ca(OH)
2
The addition of lime at various stages of the Bayer process provides numerous
benefits to the process, including:
• improving the dissolution of boehmite and diaspore during digestion
• helping to reduce liquor impurities (desilication and causticisation),
• assists in phosphate control in pregnant liquor,
• reduces soda losses in red mud. [23]
CaO (or burnt lime) is usually produced on an industrial scale by the calcination
of CaCO3. Water is then added to form Ca(OH)2
(slaked lime), which is stored as
a slurry called milk of lime. This slurry can then be added at different stages
throughout the refinery. The investigation by Giles et al. [67] reports that
hydration of CaO in water or caustic solutions proceeds as follows:
12. CaO + H2O → Ca(OH)2 → Ca2+ + 2OH- (surface) → Ca2+ + 2OH- (aq)
There are a number of factors that influence the rate of hydration including; 1)
CaO particle size, 2) hydration temperature (hydration rate increases with
increasing temperature, [68-71] and 3) impurities (chloride, sulfate, carbonate and
12
Figure 1.6: CaO-Na2O-CO2-Al2O3-H2O phase diagram. [23]
13
hydroxide). The CaO particle size is primarily controlled by the CaCO3
calcination temperature, where low-temperature calcination (900 °C) forms CaO
with a small particle size. It has been reported by Volzhenskii and Vinogradov
[70] that CaO hydration occurs rapidly for small CaO particles, where hydration
activity reduces as the CaO particle size is greater than 30 μm. Libby [72] also
reports an increase in the CaO slaking rate with decreasing surface area.
Konstantinov [73] observed a doubling in the rate of CaO hydration in the
presence of 1-3 % chloride and a decrease in hydration of a factor of 10 when 0.1
– 3 % of sulfate and carbonate are present. A similar observation has been
reported by Boynton. [68] Hydroxide ions have also shown a decrease in
hydration activity. [69]
The presence of impurities in Bayer liquor has a significant influence on the
solubility of Ca(OH)2. Hydroxide, carbonate, and sulfate all significantly reduce
the solubility of Ca(OH)2 in pure water. [23] However, an increased carbonate
concentration in synthetic liquor or plant liquor (143 °C) has been reported to
increase the calcium ion solubility slightly. [68, 70] The and Sivakumar [74]
report organic impurities with a number of hydroxyl groups are responsible for an
increase in the solubility of Ca(OH)2
attributed to the formation of an
organic/calcium ion complex. A decrease in calcium ion solubility is observed at
temperatures between 143 to 235 °C when organics are present.
1.2.3.1. Causticisation
Causticisation involves the removal of carbonate (Na2CO3) from Bayer liquors
through the addition of CaO or Ca(OH)2 to form CaCO3
:
13. Ca(OH)2 (or CaO ) + Na2CO3 CaCO3
14. 3CaO·A1
+ 2NaOH
2O3·6H2O + 3Na2CO 3 3CaCO 3 + 2NaAl(OH)4
+ 4NaOH
The CaO-Na2O-CO2-Al2O3-H2O phase diagram examining the NaOH/Na2CO3
equilibrium is shown in Fig. 1.6.
Solymar and Zoldi, [75] and Young [71] have all found that decreasing the caustic
concentration or increasing the reaction temperature thermodynamically favours
14
the formation of CaCO3
. Higher causticisation efficiencies have been found for
longer reaction times and increasing agitation. [76]
1.2.3.2. Tricalcium aluminate hexahydrate (TCA)
Tricalcium aluminate (TCA) readily forms when CaO or Ca(OH)2 reacts with
sodium aluminate solutions, and has a chemical composition of Ca3A12(OH)12.
The oxide composition (commonly used in the cement industry) of TCA is
3CaO·A12O3·6H2O, which is commonly shortened to C3AH6 where C=CaO,
A=Al2O3, and H=H2O. The primary use of TCA in the Bayer industry is as
filtration media for the isolation of precipitated gibbsite (product) from digested
liquor. However, TCA production consumes both caustic and aluminate and
reduces the overall yield of gibbsite (Al(OH)3), so it is desirable to efficiently
utilise this material. If conditions are not controlled properly during the reaction,
particularly temperature and residence time, TCA can “coat” Ca(OH)2 particles,
preventing full reaction and reducing lime utilisation efficiency. [23] It has also
been reported that TCA formation reduces the TiO2 content in gibbsite, [77] and
the soda content (hydrogarnet formation at high temperature, 250 °C, digestion) in
red mud residue. [23] Hydrogarnet is formed by the incorporation of silica into the
TCA structure giving the general formula Ca3Al2(SiO4)n(OH)(12-4n)
, where ‘n’
specifies the amount of silica present.
1.3. Surface chemistry
The surface chemistry of red mud is extremely complex due to the variable
composition of red mud particles. It is also difficult to determine the exact surface
chemical composition of these particles due to the thin surface layer thickness of
iron oxide, 50 Å to 1 μm. [3] However, it is well known that the majority of the
minerals and oxides of which red mud is composed demonstrate acid/base type
behaviour in aqueous solutions, [78, 79], so red mud particles should exhibit
similar behaviour. The acid/base properties of these particles are believed to be
due to the surface hydroxyl groups. [3] The specific surface area and adsorption
capacity for protons of acid treated red mud has been found to be 20.7 m2g-1 and
2.5 x 10-2 mol g-1, respectively. [80] Santona et al., [18] found the surface area of
red mud varies for non-treated and acid neutralised samples, 18.9 and 25.2 m2g-1.
15
The increase in surface area after acid neutralisation was attributed to the partial
dissolution of red mud species, possibly cancrinite which showed a 9 wt.%
decrease after neutralisation. [18]
Chevedov et al., [3], studied the surface properties of red mud by means of
potentiometric titration, [81, 82] and found that three zones (Fig. 1.7) existed due
to different mechanisms occurring at the red mud surface. Red mud particles can
consume H+ without a change in pH (Zone I) due to the presence of free
hydroxide ions (OH-) reacting with protons (Eq. 15) more readily than the ionised
surface hydroxyl groups. However, small amounts of surface hydroxyl groups
were found to protonate in this zone. The inflection point between Zones I and II
represents red mud particles in basic aqueous solutions carrying ionised surface
hydroxyl groups (S-O-
17
) consuming protons (Eq. 16). At neutral pH, Zone II, all
free hydroxide ions are consumed and protons added to the slurry are then
consumed by surface hydroxyl groups resulting in a constant pH (buffering). Once
these surface hydroxyl groups have been neutralised, accumulation of protons in
solution causes the second rapid drop in pH, Zone III (Eq. ).
15. Zone I: OH- + H+ → H2
S-O
O (primary reaction) - + H+
16. Zone II: S-O
→ S-OH - + H+
17. Zone III: S-OH + H
→ S-OH + → S-OH2
+
The amount of surface hydroxyl groups is roughly proportional to the reactive
silica content in the original bauxite. [83]. Sodalite is a zeolite-type compound
with a high surface area of exposed oxygen atoms that react with protons. [3].
Estimates for the number of surface hydroxyl groups on red mud obtained by
Chevedov et al., [3], were two orders of magnitude higher than the average values
obtained for metal oxides, [84], suggesting that a high level of sodalite on the
surface of red mud was present.
The surface charge of red mud can be derived from pH measurements and
determined by literature methods. [85-87] The point of zero charge (PZC) can be
used in the determination of the surface charge properties of materials, [85-87],
and is defined as the pH at which the net charge on the surface is zero. The PZC
16
Figure 1.7: Titration curves of red mud slurry (dotted line)
and caustic solution (solid line). [3]
17
provides an estimate of the acidity of the oxide surface. For most alumina and iron
oxides the PZC is approximately 7-8, [88, 89] with Fe2O3 and Al2O3
to 8.8 determined by potentiometric titration, [90] while goethite has a PZC of
around 8.9 to 9.5. [29, 91, 92] Some studies have shown that red mud can have a
PZC value of about 6.5, [3], while others have reported PZC values of around 8.3.
[78, 93, 94] Red mud with high silica content usually has PZC values of 6.3,
which suggests that the presence of these compounds reduces the PZC value. The
presence of different oxides in red mud, means that there are not only neutral
surface complexes and SOH sites at PZC, but also both positively charged (such
as FeOH
having PZC
values of 8.5 and 9.2 respectively. The PZC of hematite has been found to be 8.5
2+ and AlOH2
+) and negatively charged (such as TiO- and SiO-
) surface
complexes. [93] The shift in PZC to lower values is believed to be attributed to
the formation of differently charged oxide surface sites, and the release of free
hydroxide ions back into solution resulting in the increase in positive surface
charge.
1.4. Removal of trace metals from solution
Red mud has a strong binding capacity for heavy metals. [6, 11, 12, 14] Red mud
has the ability to adsorb trace metals from solution onto the very fine grained iron
oxides. These finely grained particles have high surface/volume and high
charge/mass ratios when the pH of the solution is above 5, [95, 96] which
increases the ability of red mud to remove trace metals. Increased adsorption
efficiency can also be achieved by ensuring the solution pH is greater than 5. [97,
98] High adsorption affinity of heavy metals on red mud is attributed to the
chemisorption reactions at the surface of the oxide components of red mud (e.g.
Fe2O3, Al2O3, and TiO2
), however, identification of the oxide with the highest
affinity for a given metal ion has not been determined. [17, 18, 99] The ability of
red mud to remove trace metals from solution increases over time (240 hours),
where 1 kg of aged dry red mud was able to remove approximately 1000 meq./kg
of trace metals from solution. [99]
18
Adsorption of heavy metals from solution increases with increased contact of the
solution with red mud, rendering heavy metal removal a time dependent process.
The metal concentrations retained in red mud can be calculated using Eq. 18.
Santona et al., [18], investigated red mud with high levels of cancrinite (zeolite-
like structure), and suggested that higher adsorption values obtained were due to
the presence of large quantities of cancrinite, which incorporated the heavy metal
cations in the cages and channels of its structure.
18. qe = (C0 – Ce
where q
)V / m
e is the sorbent phase (mg/g), C0 and Ce
are the initial and final
equilibrium concentrations of the metal ion in solution (mg/litre), V is the
solution volume (litres) and m is the mass of the sorbent (g).
The mechanism for the removal of dissolved metals using red mud has been
proposed to be comprised of four different processes:
i) co-precipitation of insoluble metal hydroxides that form successive layers
on the red mud surface,
ii) formation of kinetic intermediates [Fe2(OH)4]2+, [Fe3(OH)4]5+,
[Al4(OH)8]2+, and [Al8(OH)20]4+
iii) chemical adsorption which removes metal ions as uncharged hydroxides
condensed onto surface hydroxyl groups exposed on the red mud surface,
[100] and
, at the adsorbent surface,
iv) ion exchange.
The dominant mechanisms of removal are believed to be (i) and (iii). [101, 102]
2. Seawater neutralised bauxite refinery residues
2.1. Introduction
Bauxite refinery residues are characterised by relatively high concentrations of
sodium aluminate and sodium carbonate and a variety of anionic species. If left
untreated, these species will be detrimental to the environment. Therefore,
systems have been developed to remove these species prior to disposal. Several
groups have explored seawater neutralisation of bauxite refinery residues. [9, 10,
103, 104] A number of alumina refineries have implemented this process, and
19
found it provided a reduction in both pH and dissolved metal concentrations.
Glenister and Thornber, [9], concluded disposal of refinery residues at pH 8 was
optimal, since at this pH chemically adsorbed Na is released, neutralising alkaline
buffer minerals and rendering most of the dissolved metal species insoluble. This
coincides with the recommended pH value outlined by environmental
departments. [105] The addition of seawater to red mud residues reduces the
alkalinity of the slurry through the precipitation of Mg, Ca, and Al hydroxide and
carbonate minerals. [106] Some researchers have investigated the neutralisation of
red mud with strong acids, [79, 101, 106, 107] and have found that the initial
addition of acid results in a rapid decrease in pH, followed by the leaching of
alkaline solids from the red mud causing a slow rise in pH.
Implementation of seawater neutralisation of red mud at Queensland Alumina
Ltd. (QAL) initially began as an alternative to the use of freshwater, [103] and led
to the discovery of numerous benefits, including:
i) a decrease in freshwater use, [103]
ii) increased settling rates of ponds due to agglomerate consolidation, [108]
iii) decreased alkalinity and sodicity in the solid refinery residue and entrained
liquor, [103]
iv) increased acid neutralisation capacity, and
v) improved soil properties after rehabilitation.
2.2. Reaction Mechanism
The addition of seawater to un-neutralised red mud results in the formation of fine
mineral particles that flocculate into larger agglomerates. Multivalent exchange
cations, Ca and Mg, form electrostatic bridges, [109] which then act as nucleation
sites for the precipitation of magnesium and calcium hydroxides. Hanahan et al.,
[10], reported an increase in electrical conductivity indicating the increase in
soluble salt content. Formation of these hydroxides reduces the concentration of
hydroxide ions in solution, therefore reducing pH. [110] As the electrostatic
conditions of the surface changes, the agglomerates tighten, pH decreases, and
elements that exhibited colloidal behaviour initially at high pH lose stability.
[109] The further decrease in pH causes the precipitation of hydroxycarbonates of
20
aluminium, calcium, and magnesium, where the precipitation of hydrotalcite-like
compounds becomes favoured. [10]
Seawater neutralisation does not eliminate hydroxide from the system but
converts the readily soluble, strongly caustic refinery residue into less soluble,
weakly alkaline solids. The carbonate and bicarbonate alkalinity of the waste is
primarily removed through the precipitation of calcite and aragonite. [110]
McConchie et al., [99], described the seawater neutralisation process as the
precipitation of hydroxyl ions predominantly as brucite, but also as boehmite,
gibbsite, hydrocalumite, hydrotalcite, and p-aluminohydrocalcite. Most of these
species are already present in red mud, however, the reduction in pH after
seawater neutralisation influenced the continuation of crystal growth as
aluminium became less soluble. [99] Menzies et al., [7], reported the formation of
a white precipitate containing hydrotalcite, aragonite, and pyroaurite, determined
by XRD. The extensive characterisation of seawater neutralised red mud by
Hanahan et al., [10], revealed the complexity of the system, identifying 15
different mineral components (XRD). The major elemental components of
seawater neutralised red mud, determined by acid digestion and ICP-MS, were
Fe > Na > Al > Ca > Si > Mg. [10] Variations in reported values and components
of seawater neutralised red mud are due to the differences in physical, chemical,
and mineralogical properties of red mud.
2.3. Formation of hydrotalcite
The seawater neutralisation of aluminate liquor studies done by Smith et al., [111,
112], reported that the exact composition of the precipitate, including hydrotalcite,
calcite and aragonite, is dependent on the precipitation conditions. Smith et al.,
[111, 112], found that the composition of the hydrotalcite is dependent on the pH
at neutralisation: hydrotalcite formed at high pH (pH > 13) had a Mg:Al ratio of
2:1 (Eq. 19), while those precipitated at lower pH (below 9) had a Mg:Al ratio of
4:1 (Eq. 20). At high pH a more stable microcrystalline carbonate hydrotalcite
(Mg4Al2(CO3)(OH)12·xH2O) forms, due to the readily adsorbed CO2 from the
atmosphere producing a saturated carbonate solution. At lower pH (pH < 9.5) a
less well defined crystal structure forms. The decrease of available carbonate in
solution results in the intercalation of other anions into the hydrotalcite structure
21
(Mg8Al2Cl(CO3)0.5(OH)20·xH2O). The decrease in available carbonate is due to
the rapid decrease in hydroxide ions from solution resulting in a lower adsorption
of CO2
, and therefore a decrease in available carbonate anions for intercalation
(Eq. 21).
19. 4MgCl2(aq) + 2NaAl(OH)4(aq) + NaOH(aq) + Na2CO3(aq)
→ Mg
4Al2(CO3)(OH)12·xH2O(s) + 8NaCl
(s)
20. 8MgCl2(aq) + 2NaAl(OH)4(aq) + 12NaOH(aq) + ½Na2CO3(aq)
→ Mg
8Al2Cl(CO3)0.5(OH)20·xH2O(s) + 15NaCl
21. CO(s)
2(g) + 2Na2+(aq) + 2OH-
(aq) → 2Na2+(aq) + CO3
2-(aq) + H2O
(l)
Seawater neutralised red mud would consist of both the 2:1 and 4:1 hydrotalcite,
where a small quantity of the 2:1 hydrotalcite would precipitate initially before the
predominant 4:1 hydrotalcite forms at the reduced pH. The reduced level of
carbonate in solution allows for the inclusion of other anions, such as oxy-anions
of transition metals, vanadate and arsenate, into the hydrotalcite matrix. The rate
of adsorption of anions other than carbonate depends on the concentration of
carbonate in solution. Carbonate is the predominant anion intercalated into
hydrotalcite, therefore its presence hinders the intercalation of other anionic
species. Increased temperatures showed a slight increase in adsorption efficiency,
[111] attributed to the decrease in carbonate through the conversion of carbonate
to CO2
at higher temperatures.
2.4. Adsorption of anions on the surface of neutralised red mud
Removal of contaminates is not only limited to the intercalation of species in
hydrotalcite, but also through the adsorption of contaminants onto the surface of
neutralised red mud. Genc et al., [5], investigated the adsorption of arsenate from
water using neutralised red mud and found that adsorption of arsenate increased
with decreased pH. This agrees with the work by Smith et al., [111, 112], which
showed higher adsorbent concentrations and lower initial arsenate concentrations.
Seawater neutralised red mud consists of a complex mixture of fine grained iron
and aluminium hydroxides and hydroxycarbonates that exhibit a pH dependent
surface charge, [5, 94] and it was suggested that the pH dependence of arsenate
22
adsorption onto seawater neutralised red mud was through the exchange of an
aqueous ligand for a surface hydroxyl group (Eq. 22). The number of positively
charged surface sites available for adsorption is higher at pH 6.3, and decreases
with increased pH. [5, 94] Adsorption is believed to be facilitated by the
electrostatic and chemical attraction of arsenate for the positive surface charge. [5,
94] Adsorption increases when the pH of the solution is lower than the PZC of red
mud, due to the increase of positive charge on the red mud surface. At high pH
values, anions may be competing with hydroxide ions for the positively charged
sites on the red mud surface, which causes the decrease in adsorption.
22. ≡ S-OH + L- + H+ ≡ S-L + H2
where ≡
O
S represents the seawater neutralised red mud surface
The Langmuir isotherm (Eq. 23) is a commonly used adsorption isotherm for
assessing the potential uses of an adsorbent for particular applications. The
Langmuir isotherm has been used to study the adsorption capacity of seawater
neutralised red mud. [17, 18] To determine whether anion adsorption by seawater
neutralised red mud is a high-affinity adsorption, the dimensionless constant
separation term RL
can be calculated (Eq. 24).
23. qe = (Q0bCe)/(1 + bCe
where b is the adsorption constant related to the enthalpy of adsorption
(1 μmol
)
-1), Q0 is the adsorption capacity (μmol g-1), and Ce
24. R
is the
equilibrium concentration (μM).
L = 1 / (1 + bC0
where C
)
0
is the initial anion concentration (μM). [24, 113]
The parameter RL indicates the shape of the adsorption isotherm and 0 < RL
< 1
corresponds to high affinity adsorption. [5] Arsenate adsorption by seawater
neutralised red mud was found to be very efficient regardless of the pH or the
initial concentration. [5] Altundogan et el., [114], has reported adsorption follows
the chemisorption mechanism for heavy metal cations.
23
3. Layered double hydroxides – LDHs
3.1. Introduction
Layered double hydroxides (LDHs) have been extensively researched for many
years as host materials for a range of anionic exchange reactions, especially the
removal of anionic impurities from solution. [115-124] They are sometimes
referred to as anionic or hydrotalcite-like clays, and are based on the brucite
structure, Mg(OH)2. [125-127] LDH are represented by the general formula,
[M2+1-x M3+
x(OH)2]x+Am-x/m·nH2O, where M2+ is a divalent cation, M3+ is a
trivalent cation and A is an interlamellar anion with charge m-
i) boehmite (α-AlOOH) for x > 0.337,
. Pure LDH phases
exist for 0.2 ≤ x ≤ 0.33. Values outside the specified x range will form:
ii) hydromagnesite 4MgCO3·Mg(OH)2·4H2
iii) a mixture of hydromagnesite and Mg(OH)
O) for 0.105 < x < 0.201, and
2
for x < 0.105. [128-131]
Hydrotalcite is produced when M2+ = Mg2+ and M3+ = Al3+, giving the general
formula Mg6Al2(OH)16CO3·4H2O. LDHs consist of layers of metal cations (M2+
and M3+) of similar radii, which are randomly distributed in the octahedral
positions, which form brucite-like structures Mg(OH)2 (Fig. 1.8). The enthalpy of
bond formation within the layers is largely responsible for the thermodynamic
stability of these layered materials. [132] The brucite-type layers are stacked on
top of each other and are held together by weak hydrogen bonding interactions
(Fig. 1.9). [133] Substitution of divalent cations for trivalent ones gives rise to
positively charged layers, where a maximum of one in three trivalent sites are
substituted by a divalent cation. [129] The ratio of M2+ to M3+ cations determines
the degree to which the framework is positively charged, where a low M2+:M3+
ratio will result in highly positively charged layers. To maintain electroneutrality,
the interlamellar domain must be occupied by an adequate number of anions,
which are generally hydrated. [128, 134, 135] Charge neutrality is not confined to
the interlayer region, but also to the external surfaces of the LDH structure. The
resulting mineral has layers of ordered anions between hydroxyl sheets, giving
hydrotalcites the acronym LDH or ‘layered double hydroxides.’ As there is no
overall charge, hydrotalcites are quite stable.
24
Figure 1.8: Schematic representation of the hydroxide layers in the hydrotalcite.
Figure 1.9: Schematic representation of the hydrotalcite structure.
25
The interlayer region of LDHs are complex, consisting of anions, water
molecules, and other neutral or charged moieties. A large variety of anionic
species can be positioned between the hydroxide layers, including halides, oxy-
anions, oxy and polyoxy-metallates, anionic complexes, and organic anions. [136]
The interlayer interactions of LDHs are mediated by columbic forces between the
positively charged layers and the anions in the interlayer, and also hydrogen
bonding between the anions and interlayer water molecules. [136, 137] Water
molecules are connected through extensive hydrogen bonding to the hydroxyl ions
of the metal hydroxide layers and interlayer anions. [135, 138, 139] The quantity
of water present in the interlayer is governed by the nature of the interlayer
anions, water vapour pressure, and temperature. [140-144] Khan and O’Hare
found that water molecules are in a continuous state of flux, using NMR
techniques. [128] However, vibrational studies have shown that the hydrotalcite
interlayer has a highly structured yet mobile environment. [145-147]
Many types of hydrotalcites can be formed from different combinations of
divalent and trivalent cations and different interlayer anions. Some natural LDHs
are given in Table 1.2. The orientation of the ions in the interlayer is determined
by factors such as the charge of layers and the amount of interlayer water present.
The anion may be trivalent (phosphate), divalent, (carbonate, sulfate), or
monovalent, (hydroxide, chloride, or nitrate). [148-154]
An increase in anionic charge results in the electrostatic interactions between the
positively charged hydroxide layer and the anion to become stronger, therefore
rendering a more stable hydrotalcite, with a decrease in interlayer distances. This
means the formation of a hydrotalcite with a divalent anion is more favourable
over one containing monovalent anions. [155-157]
26
Table 1.2: Compositions, crystallographic parameters and symmetries for
some natural LDHs.
Name Chemical Composition
Unit Cell
Parameters Symmetry
a (nm) c (nm)
Hydrotalcite Mg6Al2(OH)16CO3·4H2 0.3054 O 2.281 3R
Manasseite Mg6Al2(OH)16CO3·4H2 0.3100 O 1.560 2H
Meixnerite Mg6Al2(OH)16(OH)2·4H2 0.3046 O 2.292 3R
Pyroaurite Mg6Fe 2(OH)16CO3·4.5H2 0.3109 O 2.341 3R
Sjögrenite Mg6Fe 2(OH)16CO3·4.5H2 0.3113 O 1.561 2H
Caolingite Mg10Fe 2(OH)24CO3·2H2 0.3120 O 3.750 3R
Iowaite Mg4.63Fe 1.32(OH)-
16Cl1.22·1.95H20.3119
O 2.425 3R
Stichtite Mg6Cr2(OH)16CO3·4H2 0.3100 O 2.340 3R
Barbertonite Mg6Cr2(OH)16CO3·4H2 0.3100 O 1.560 2H
Desautelsite Mg6Mn2(OH)16CO3·4H2 0.3114 O 2.339 3R
Takovite Ni6Al2(OH)16CO3·4H2 0.3025 O 2.259 3R
Reevesite Ni6Fe2(OH)16CO3·4H2 0.3081 O 2.305 3R
Where, 3R represents a rhombohedral stacking, while 2H represents a hexagonal stacking.
3.2. Preparation of LDHs
A variety of methods exist for LDH production such as co-precipitation, [158-
160] urea reduction, [161, 162] salt-oxide method, [163] hydrothermal, [160, 163]
electrochemical, [164, 165] and sol-gel. [166] The most frequently used methods
are co-precipitation and urea reduction, while electrochemical and sol-gel are the
least used methods. Co-precipitation is based on the slow addition of a mixed
solution of divalent and trivalent metal salts to an alkaline solution in a reactor,
which leads to the co-precipitation of the two metallic salts. Formation of the
LDH is based on the condensation of hexa-aqua complexes in solution that form
the brucite-like layers containing both metallic cations. [136] Interlamellar anions
either arise from the counter-anions of the metallic salts, or anions from the
alkaline solution. At high pH, hydroxyl ions are prevalent and therefore can be
intercalated, however if the alkaline solution is prepared with sodium carbonate
27
the intercalated anion is carbonate due to its higher affinity for the LDH
interlamellar region. [167]
In order to obtain well organised phases, the preparation conditions have to be
optimised for the desired product. For well ordered hydrotalcite-like structures to
form, a pH range between 7 and 10 is required. At lower pH values, an amorphous
compound is obtained, while at higher pH values Mg(OH)2
crystallises with the
LDH phase. [136] The study conducted by Crepaldi et al., [134], on the
comparison of constant and variable pH co-precipitation reactions demonstrated
that maintaining a constant pH throughout the reaction yielded LDHs with higher
crystallinity, smaller average particle sizes, higher average specific surface area,
and higher average pore diameters, in comparison to those produced with variable
pH. Scanning electron microscopy showed that variable pH also leads to
heterogeneous products, due to the different precipitates produced initially at high
pH, while those obtained at lower pH showed homogeneously aggregated
particles. [134]
3.3. Anionic exchange
Recent studies have focused on using LDHs to undergo anionic exchange
reactions across a wide range of applications, especially the removal of toxic
anions from aqueous systems. [115, 168-171] The interlayer region is less stable
than the brucite-like layers, and therefore readily undergoes anion exchange. The
interlayer interactions can be direct, [172], or mediated through other species
present in the interlayer region. [173] LDHs predominantly have mediated
interlayer interactions, making the mechanism for anion exchange complicated.
Uncertainty exists in the literature regarding the exact mechanism of LDH anion
exchange. [128, 157, 174, 175] The general assumption is a topotactic
mechanism, [155, 156] however other mechanisms have been proposed including:
i) a two-step process involving the dissolution of the LDH phase followed by the
re-precipitation of a new LDH with the desired anion (D-R mechanism), [176]
ii) first order kinetics, [177] or
iii) another two-step mechanism involving the adsorption of the incoming anion
followed by the desorption of the initial anion in the interlayer. [178, 179]
28
Anion exchange reactions are thought to take place topotactically, based on the
assumption that a close structural relationship between parent and product phases
exists. The only structural change brought about by anion exchange is a variation
in the interlayer distance, which is dependent on the size of the incoming anion.
However, observations have been noted in recent studies that suggest the anion
exchange reaction follow the D-R mechanism. The observations included a mass
loss during anion exchange, which can be attributed to bulk dissolution, [180-
182], and unitary salts formed as impurity phases during anion exchange
reactions. [181]
According to the topotaxy mechanism, [183] the lamellar structure of LDHs
allows for diffusion of anionic species in the interlayer regions for anions of
higher affinity (Eq. 25).
25. [MII-MIII-X] + Y → [MII-MIII
-Y] + X
where MII-MIII
X represents the anionic species in the interlayer
are positively charged hydroxide layers
Y represents an anionic species with a higher affinity for the
interlayer region which will replace X.
According to Eq. 25, the outgoing X anion is exchanged for the incoming Y anion
in a single step, where the host hydroxide layer essentially remains unperturbed. A
two-step topochemical reaction has also been proposed, [184], where the initial
step is the separation of the LDH lattice into its corresponding positively charged
hydroxide layers and free anions (Eq. 26) followed by restacking of the layers to
form the LDH with the new anionic species incorporated into the interlayer region
(Eq. 27).
26. [MII-MIII-X] → [MII-MIII]+
+ X
27. [MII-MIII]+ + Y → [MII-MIII
-Y]
It was surmised that under specific temperature, pH and anion concentration
conditions, the precursor LDH could dissolve (dissolution step, Eq. 28), followed
29
by re-precipitation with the incoming anions, (Eq. 29). Intercalation of the new
anionic species is based on 2 factors; (i) they have a higher affinity than the
original anionic species, and (ii) the formation of LDH has a greater
thermodynamic stability than the original LDH structure, reflected by a lower
solubility product. [132] Radha et al., [132], proposed the D-R mechanism, based
on the fact that no reliable estimates of the strengths of these interactions and how
they compare with the strength of interlayer bonding has been reported in the
literature.
28. [Mg2Al(OH)6]NO3·2H2O → 2Mg2+ + Al3+ + 6OH- + NO3- + 2H2
O
29. 2Mg2+ + Al3+ + 6OH- + 1/nXn- + 2H2O → [Mg2Al(OH)6](Xn-)1/n·2H2
O
Identifying which mechanism is responsible for anion exchange is difficult due to:
i) the high rate of anion exchange reactions, making kinetic studies difficult,
ii) intermediate phases are highly unstable and react quickly to form new LDH
phases,
iii) in the D-R mechanism, the dissolution of LDH takes place at the solid-liquid
interface. [132]
Extensive studies by Miyata et al., [131, 139, 185-187], exposed the anionic
exchange properties of a number of species, establishing a ranking of affinity for
intercalation. Hydrotalcite shows the greatest affinity for anions of high charge
density. [186, 188] The affinity of monovalent anions was determined to be
OH- > F- > Cl- > Br- > NO3- > I-, while the order for divalent anions was
CO32- > SO4
2-
. The carbonate anion has proven to be the predominate anion for
intercalation, and once intercalated is very difficult to exchange with other anions.
This high affinity prevents its use as an anion-exchange material in Mg,Al
hydrotalcites, unless precautionary steps such as a nitrogen atmosphere, carbonate
free solutions or calcination are used.
Theoretically, LDHs have an anion exchange capacity of 3.6 mequiv./g if all the
carbonate in the general formula was exchanged. [186] Experiments conducted by
Miyata et al., [186], showed that a hydrotalcite prepared under a nitrogen
atmosphere with carbonate free solutions could obtain an anion exchange capacity
30
of 3 mequiv./g. The theoretical capacity value cannot be obtained due to
hydroxide anions present in solution competing with the desired anion. [186]
Removal of carbonate from all sources is essential in exchange reactions, as any
carbonate present in the exchange solutions will be incorporated preferentially to
other anions. Anion exchange capacity values were determined by comparing the
anion concentrations of the initial and final solutions after the addition of a known
amount of hydrotalcite by atomic adsorption spectroscopy and the Dionex
method. [186]
3.4. Thermal activation of hydrotalcite materials
Recent studies have shown that LDHs can have a so-called ‘memory effect’
whereby a hydrotalcite material can be thermally treated to remove water,
hydroxyl, and carbonate units from its matrix, then re-hydrated in an aqueous
solution to reform the original structure. [161, 189] The restoration of the layered
structure in hydrotalcites is a ‘structural memory effect’. [190-192] This effect can
be used effectively to remove harmful anions, both organic, [120, 122] and
inorganic, [159, 161, 193, 194] from waste water solutions.
The calcination of hydrotalcite, from temperatures of 350ºC to 800ºC, removes
interlayer water, interlayer anions (carbonate anions), and hydroxyls. The result is
the formation of periclase-like Mg,Al oxides. This dehydration process makes the
hydrotalcite product chemically more reactive. Therefore, exposure of the
dehydrated structure with a solution of anions will immediately re-hydrate the
hydrotalcite removing anions within the solution. XRD studies have shown the
collapse of the crystalline hydrotalcite to an amorphous magnesium oxide with
dispersed aluminium ions as a solid solution. [159, 161, 194, 195] The carbonate
anions are decomposed to carbon dioxide (CO2) and O2-, leaving O2-
anions
between the layers. [131, 186, 196, 197] Re-hydrating the calcined product
regenerates the LDH, where water is absorbed to reform the hydroxyl layers, as
well as being absorbed into the interlayer along with the anions in solution. [122]
Anions that are reabsorbed are not necessarily the original anions, since any
available anion in the re-hydrating solution will be absorbed. For example, the re-
hydration of calcined hydrotalcites in carbonate free solutions will yield a
31
carbonate free hydrotalcite. Parker et al., [195], reported a 50 % decrease in
adsorption in the anion exchange capacity of LDHs due to a slight alteration in the
re-formed hydrotalcite. Heating to temperatures above 900 ºC produces spinel
(MgAl2O4
), totally degrading the hydrotalcite lattice and preventing any
reformation.
3.5. Characterisation of LDHs
3.5.1. Vibrational spectroscopy – infrared and Raman Spectroscopy
Spectroscopy has been a widely used technique in the industry for the structural
and compositional analysis of inorganic, organic, organometallic, metalorganic,
and polymeric materials. Vibrational spectroscopy involves the use of light to
probe the vibrational behaviour of molecular systems, usually via absorption,
emission, or light scattering experiments. Both infrared and Raman spectroscopy
give rise to a vibrational spectrum as a set of absorption or scattering peaks,
corresponding to the energies of transitions within the sample (wavenumber of
vibrational modes).
3.5.1.1. Hydroxyl stretching and bending vibrations
The vibrational spectra of hydrotalcites exhibit various forms of water hydroxyl-
stretching vibrations. These include water in the interlayer between the hydroxide
layers, which may or may not form bridging-type bonds with the exchangeable
anions, water adsorbed on the outer surface, and free water between layers. Water
hydroxyl-stretching vibrations are intense in an infrared spectrum, because of the
large change in dipole moment during the vibration of water, however this
vibration is not always observed in the Raman spectrum. Therefore, the
comparison of the two techniques allows for the identification of the bands
associated with water and those associated with hydroxyl stretching vibrations.
Water bending modes are situated around 1600-1700 cm-1 accompanied by OH-
stretching vibrations in the 3000-4000 cm-1
region. [151, 162, 198, 199]
The replacement of Mg2+ by Al3+, in hydrotalcites, results in stronger hydrogen
bonds between the hydroxide layers, when compared with brucite, due to Al3+
32
having a higher charge and smaller ionic radius. [200] This change in O-H bond
lengths can be detected in infrared spectra with shifts to higher wavenumber in the
bending region. Shifts to lower wavenumber in the stretching region are
33
Figure 1.10: Water, hydroxyl and carbonate vibrations in
the interlayer of Co and Ni hydrotalcites. [204]
34
associated with the strength of the hydrogen bonds. [136] A similar observation
can be seen for the lattice translation modes in the low frequency region of the
infrared spectra. [201] The OH-stretching vibration for brucite is situated around
3570-3555 cm-1, while for Mg,Al hydrotalcites the corresponding band is located
at around 3450 cm-1
. This shift is associated with the shorter O-H bonds existing
in hydrotalcite than in brucite, causing an increase in the electrostatic attraction
within the hydrotalcite layer. [201]
Extensive overlapping of bands exists in the OH-stretching region of LDHs
between metal-OH bands of the hydroxide layers and the OH-bands of water. For
water adsorbed on clay minerals the OH-stretching modes of weak hydrogen
bonds occur in the region between 3580 and 3500 cm-1, while strong hydrogen
bonds are observed below 3420 cm-1. Water coordinated to cations shows
stretching vibrations around 3220 cm-1. [136] Fourier Transform IR spectra
obtained by Jose dos Reis et al., [202], showed a broad band at
3400 cm-1 assigned to the ν(OH) mode ascribed to interlayer water and hydroxyl
groups in the hydroxide layers of hydrotalcite. Numerous studies conducted by
Kloprogge and Frost (Table 1.3) have reported the infrared hydroxyl modes of
Mg,Al hydrotalcites. [136] A broad band around 3300-3000 cm-1 with a shoulder,
sometimes visible, comprised of two or three overlapping bands are attributed to
the OH-stretching vibrations and a stretching vibration of interlayer water. The
shoulder at 3050 cm-1 was assigned to hydroxyl interactions with carbonate ions
in the interlayer, [136, 139, 202-205], and has been attributed to the bridging
mode H2O-CO32-
.
The corresponding H-O-H bending vibration of interlayer water interacting with
interlayer carbonate has been found to be located at around 1750 cm-1
. [212] The
high vibrational frequency is attributed to symmetry restrictions induced by the
hydrogen bonded carbonate ions to hydroxyl groups of the hydroxide sheets. [204,
212] Fig. 1.10 shows where water and hydroxyl group vibrational bands originate
from within the hydrotalcite structure of Ni,Al and Co,Al calcined hydrotalcite
samples. Slight shifts in these values are expected for Mg,Al hydrotalcite
structures.
35
A weak peak around 1630 cm-1 in the infrared spectrum is attributed to the δH2O
mode of interlayer water, [136, 151, 198, 199, 202, 213] while the OH-bending
vibrations are located at around 1040 cm-1
in the Raman spectrum. The interlayer
anion has been found to have an effect on the position of the OH-bending mode of
interlayer water (Table 1.4).
The OH-stretching vibrational modes are weaker but sharper in the Raman
spectrum compared to the corresponding modes in the infrared spectrum. Raman
bands observed around 3600-3450 cm-1 are attributed to the stretching modes of
hydroxyl groups bonded to Al, Mg or a combination of both. Table 1.5 illustrates
some reported literature values and the corresponding assignments. Two bands
around 470 and 550 cm-1 have been assigned as hydroxyl groups associated with
Al or Mg. [204, 211] The band at 470 cm-1 is only Raman active, while the band
at 550 cm-1
has an equivalent mode in the infrared spectrum in the same location.
Table 1.3: Wavenumber (cm-1) and assignments of the hydroxide layer modes
of the types M-OH and M-O in the infrared spectra of
Mg,Al-layered double hydroxides in comparison to brucite
Mg(OH)2
Brucite
.
Reference source Assignment
[140] [207] [208] [209] [210] [211]
- CO CO3 NO
3 3
SO,
4Cl
, CO CO3 CO3 Interlayer
3 anion present
3570 2700- 4000
3597 A2uν1“Mg/Al”-OH
(OH-HOH) or
3470 3421 3392- 3422 3441 3467 A2uν2
or “Al”-OH (OH-HOH)
998 985 950 960-945 939 νsym ν
anion or def
Al-OH
865 799 853-830 874 850 870 Euν
(OH) or 2CO3
680
2-
670 651 668 671 663 635 Eu“Mg”-(OH)
(OH) or
Translation
455 Na na 616-584 556 555 553 A2u(T) “Al”-OH
or
Translation 365 Na na 426-419 451 451 Eu(T)
36
Table 1.4: Infrared water bending vibrational positions of Mg,Al hydrotalcites
as a function of the interlayer anion, as reported in literature. [136]
Interlayer anion Band position (cm-1
CO
)
3 1640 2-
CO3 1655 2-
CO3 1591 2-
CO3 1647 2-
OH 1628 -
OH 1625 -
NO3 1629 -
SO4 1642 2-
CrO4 1639 2-
V10O28 1653 6-
Table 1.5: Wavenumber (cm-1) and assignments of the hydroxide layer modes
of the type M-OH and M-OH (M represents Mg2+ or Al3+
[200]
) in the
Raman spectrum of Mg,Al layered double hydroxides.
[201] [214] Assignment
3560 3572 3580 A1g
3460
(OH) or “Mg/Al”-OH
3454 3454 A1g
(OH-HOH) or “Al”-OH
3358
1061 1053 Eg(R)
(OH)
979
695 694 Eg(R) or ν4(E’)CO
3
557 552 E
u(T)
483 476 A
1g(T)
393 388 A
1g(T)
307 303 Acoustic overtone
37
3.5.1.2. Carbonate stretching vibrations
When the carbonate species is present as a free ion, it will exhibit a planar triangle
with point symmetry D3h. Group theoretical analysis of the carbonate ion predicts
four normal modes: the ν1 symmetric stretch of A1 symmetry normally observed
at 1063 cm-1, the antisymmetric stretch of E’ symmetry observed at 1415 cm-1, the
ν2 out of plane bend at 879 cm-1 and the in-plane bend at 680 cm-1. [215, 216] All
modes are both Raman and infrared active except for the ν2
mode, which is IR
active only. Incorporation of the carbonate species into the hydrotalcite structure
will promote a shift towards lower wavenumbers, due to the interaction of
carbonate with interlayer water molecules and/or hydroxyl groups from the
hydrotalcite layer.
Hydrotalcites with carbonate incorporated into the interlayer typically show
infrared bands at around 1360-1400, 875 and 670 cm-1. Assignments for the
carbonate modes are outlined in Table 1.6. A strong peak at around 1360 cm-1
observed by Jose dos Reis et al., [202], attributed to the ν3 mode of the carbonate
species, agreed with literature values. An additional band at 1550-1500 cm-1 has
been reported, [217, 218], and attributed to the formation of a bicarbonate ion
upon dehydration (proton transfer from the hydroxide sheets to the carbonate ion).
The presence of this band indicates a change in the carbonate symmetry. In the
Raman spectrum the symmetrical stretching vibration ν1(A’1), the antisymmetric
stretching vibration ν3(E’), and the bending angular vibration ν4(E’) around 1063,
1415, and 680 cm-1, respectively, are observed for the free carbonate anion. [200,
215, 216] Weak ν3 and ν4 band modes have been observed at around 1053 and
1403 cm-1
. [200, 219]
Table 1.6: FT-IR interlayer carbonate vibrational modes. [209, 211]
Mode Mg,Al hydrotalcite (cm-1 Mg,Al hydrotalcite (cm) [211] -1
Δν
) [209]
36 3 36
ν 1401 3a 1400
ν 1365 3 1364
ν 1012 1 1060
ν 870 2 874
ν 667 4 671
38
Exposure of hydrotalcite leads to the adsorption of CO2 onto the hydrotalcite
structure and has been characterised by infrared spectroscopy. [204, 220-222]
These studies reported the three different carbonate species: (i) unidentate (ii)
bidentate, and (iii) bicarbonate (Fig. 1.11). These different carbonate species
reflect different types of surface basic sites and their relative strengths. Unidentate
carbonates were proposed to be bonded to high-strength basic sites, bidentate
carbonate to medium-strength basic sites, and bicarbonate to low-strength basic
sites. [221, 222] The same relationship was seen in a study by Di Cosimo et al.,
[220] Morterra et al., [221] and Philipp et al., [222], which reported that the
strength of the surface basic sites depended on the Al content of the adsorbing
species - an increase in Al content increases the basic site density. The increase in
site density is attributed to rearrangement of the MgO lattice by Al3+
cations,
forcing adjacent oxygen anions to become co-ordinately unsaturated.
Di Cosimo et al., [220], reported unidentate carbonate exhibited a symmetric
O-C-O stretching vibration at 1360-1400 cm-1 and an antisymmetric O-C-O
stretching vibration at 1510-1560 cm-1, while bidentate carbonate showed a
symmetric O-C-O stretching vibration at 1320-1340 cm-1 and an antisymmetric
O-C-O stretching vibration at 1610-1630 cm-1. The bicarbonate species involves
surface hydroxyl groups and showed a C-OH bending mode at 1220 cm-1 as well
as symmetric and antisymmetric O-C-O stretching modes at 1480 and 1650 cm-1
,
respectively. Vibrational modes reported by Pérez-Ramírez et al., [204], for Ni,
Al, and Co, Al calcined hydrotalcites were shifted to lower wavenumbers than
those reported by Di Cosimo et al. [220]
3.5.1.3. Lattice translational modes
The lower wavenumber region of the infrared spectrum, 1000-400 cm-1, are
complicated due to the presence of lattice translational modes (650 cm-1),
librational modes of hydroxyl and water molecules (1000-700 cm-1), Al-O bonds
(450 cm-1), [210] and the ν4(E’) carbonate band (680 cm-1). [200] A broad
complex band is also observed at around 650-600 cm-1 with [201, 209, 210, 219]
or without [200, 219] a separate band at 550 cm-1 due to Al-O and Mg-O bonds.
Interpretation of a band at 870 cm-1 appears to have some disagreement between
authors, with some ascribing the band to the ν2(A’’2) mode of the interlayer
39
Figure 1.11: Infrared bands of adsorbed CO2
species on calcined hydrotalcite. [220]
surface
40
carbonate, [200, 214, 219], while Kagunya, [201], ascribed the band to the
Eu(R)
(OH) mode for LDHs with not only carbonate, but also with nitrate and
hydroxyls as the interlayer anions. Both assignments are plausible due to the
broadness of the band, indicating that a possible overlap of both bands may exist.
3.5.2. Thermal analysis – TGA/DTG
The decomposition of the Mg,Al hydrotalcite structure occurs in three steps:
i) removal of adsorbed water (< 100 ºC),
ii) elimination of the interlayer structural water (100 – 200 ºC), and
iii) the simultaneous dehydroxylation and decarbonation of the hydrotalcite
framework (300 – 400 ºC). [131, 134, 185, 220, 223-225]
A fourth decomposition step may occur for the loss of either a volatile anion
species ( e.g. Cl-, NO3-, and CO3
2-) or a non-volatile species in which the anion is
included in the formation of a mixed metal oxide. [131, 185, 225] The
determination of the decomposition steps of hydrotalcite depends on the dryness
of the sample, stability of the interlamellar species, and possible guest-host
interactions mobilising the hydroxyl groups in the hydrotalcite lattice. [185] The
thermal decomposition of carbonate hydrotalcites consist of two decomposition
steps between 300 and 400 ˚C, attributed to the simultaneous dehydroxylation and
decarbonation of the hydrotalcite lattice. Water loss ascribed to dehydroxylation
occurs in two decomposition steps, where the first step is due to the partial
dehydroxylation of the lattice, while the second step is due to the loss of water
interacting with the interlayer anions. Dehydroxylation results in the collapse of
the hydrotalcite structure to that of its corresponding metal oxides, including
MgO, Al2O3, and MgAl2O4
(at temperatures over 900 ºC). [150, 223] The exact
decomposition product relies on the hydrotalcite and its counter balancing anions.
The rate of dehydroxylation has been used as a measure of the thermal stability of
the hydrotalcite structure, where a delay in dehydroxylation indicates a more
thermally stable hydrotalcite. [185, 224, 225] Hydrotalcite stability has been
found to be anion dependent, [185], suggesting that hydrotalcite stability can be
controlled by the incorporation of more stable, less reactive anions. The presence
of oxy-anions improves the stability of most hydrotalcites, and delays
41
dehydroxylation in comparison to carbonate hydrotalcites. [185] This is due to the
substantial number of hydroxyl groups interacting with an extensive network of
solvated hydrogen bonded anions. The antisymmetric shape of the DTG curve
obtained by Malherbe et al., [185], for hydrotalcite containing oxy-anions
indicated the presence of both free water molecules, and water molecules
hydrogen bonded to the anionic species. Existence of different interlamellar water
has been reported previously and was thought to be related to the charge density
of the hydroxylated brucite-like sheets. [226, 227]
3.5.3. X-ray Diffraction – XRD
X-ray diffraction techniques are traditionally used for the characterisation of
minerals. [228, 229] Identification of minerals by this technique is based on the
reflection of X-rays by the characteristic atomic lattice planes within the mineral
crystal. [230] The X-ray diffraction pattern is a measurement of the distance
between single planes of atoms in a crystal, providing a direct measure of the
height of layers as well as information about the bulk properties of the sample,
such as the crystalline phases present. [229, 230] Since different crystalline
materials have different cell parameters, space groups, and symmetry,
characteristic diffraction patterns are produced.
X-ray diffraction can distinguish between the two different stacking sequences of
the brucite-type sheets in LDHs, rhombohedral (3R) or hexagonal (2H). [209,
231] Table 1.2 gives the symmetry and cell parameters for a few different natural
LDH structures. Hydrotalcite normally crystallises with the rhombohedral 3R
stacking sequence, which is the three layer form. The parameters of the unit cell
are a and c=3c’, where c’ is the thickness of one layer (sheet + interlayer). [231]
The other stacking sequence, hexagonal (2H), usually forms manasseite, [133],
which is the two-layer form and is generally obtained at high temperatures. [232]
A third stacking sequence, (1H), has been reported for the most hydrated variety
of hydrotalcite compounds containing sulfate anions, however the symmetry of
this structure is unknown. [158]
Properties of anionic species, such as size, charge, orientation, and the interactions
of anionic species with the positively charged interlayer, contribute to the degree
42
of intercalation and the separation between layers. [128, 136] Anion exchange
reactions can be monitored by the shifts of the basal reflections 003 and 006.
[132] The typical d(003) spacing obtained for hydrotalcites is 7.9 Å. [211, 213,
233] Deviations in value of the d(003)
basal spacing are associated with the type of
anionic species intercalated into the interlayer region. Smaller basal spacings are
generally associated with ions of small ionic radii. Inorganic species are typically
smaller than organic species and as a result have smaller basal spacing values.
Miyata and Kumura, [187], showed that the separation of the layers, determined
by the (006) d-spacing, increased linearly with an increasing number of carbon
atoms of the anionic species. Kooli et al, [234], reported a high layer charge,
associated with low Mg:Al ratios, resulting in greater electrostatic repulsions
between the positively charged layers, and larger basal spacing.
3.6. LDHs in the alumina industry
LDHs have the potential to be used for the removal of a variety organic and
inorganic species in Bayer liquor. The proposed removal mechanisms are a
combination of intercalation and anionic adsorption onto the external surfaces -
smaller anions are intercalated while larger organic molecules are adsorbed. [185,
235-239]
Hydrotalcite has been examined as a method for removing humate material from
Bayer liquor. Schepers et al., [239], proposed the addition of magnesium
compounds to contaminated Bayer liquors, and found a brown precipitate formed
containing magnesium and aluminium hydroxides. The brown precipitate was
thought to be an impure hydrotalcite formed from the in situ reaction of the
magnesium salt and aluminate anion. Misra et al., [235, 236], reported that impure
hydrotalcite formed from combining magnesia [Mg(OH)2] with Bayer liquor,
while high purity hydrotalcite could be formed from calcined (500-900 ºC)
magnesia. The reduction of humate concentration in the Bayer liquor was due to
surface adsorption rather than anion intercalation. This assumption was based on
the large size of humate molecules which would not physically fit between the
hydrotalcite layers. The Queensland Alumina Ltd. (QAL) refinery has also
investigated humate removal using hydrotalcite, and found the quantity of humate
material in the liquor decreased. Again it was suggested that the positive charge of
43
the external surface of hydrotalcites are responsible for the reduction in humate
concentrations. Nigro and O’Niel, [237], investigated the use of hydrotalcite in the
removal of coloured impurities, such as ferrate, using different calcined
hydrotalcite samples between 450-650 ºC with their re-hydration in Bayer liquor.
The calcination of hydrotalcite between 450-500 ºC gave the greatest surface area
and pore volume and the most effective hydrotalcite for removal of coloured
impurities, indicating that adsorption was the predominant mechanism for ferrate
removal.
Carbonate concentrations need to be minimised in Bayer liquors for the effective
removal of impurity anions when using hydrotalcite or hydrocalumite. The high
affinity of carbonate for the interlayer region prevents efficient intercalation of
other anions. Grubbs and Valente, [240], found that hydrotalcite could be formed
without carbonate by reacting activated (calcined) magnesia with a sodium
aluminate solution containing the anion in excess. Implementation of this process
is limited as most impurities are not in excess. Studies by Perotta and Williams,
[241], found that the formation of hydrocalumite at temperatures up to 60 ºC
reduced the amount of oxalate in spent liquor. However, at higher temperatures
tri-calcium aluminate (TCA - 3CaO·Al2O3
) was the major product, with no
improvement in oxalate removal. Rosenberg et al., [242], discovered additives
that helped stabilise the hydrocalumite structure, allowing a larger range of
conditions that could be used in its formation without the undesirable formation of
TCA occurring. Large-scale impurity removal is currently not feasible, due to the
cost of recycling and recovering alumina from the LDH compounds.
44
4. Chapter summary
Seawater neutralisation of bauxite refinery residues has been employed in recent
years to reduce the pH and dissolved metal concentrations of waste water, through
the precipitation of hydrotalcite-like compounds and other Mg, Ca, and Al
hydroxides and carbonate minerals. These hydrotalcite-like compounds are able to
remove oxy-anions of transition metals through a combination of intercalation and
adsorption on the particle surfaces. Seawater neutralisation of bauxite refinery
residues has beneficial consequences for red mud management, such as greatly
reduced storage volumes and a much lower risk of potential environmental
impacts.
Used in an appropriate way, layered double hydroxides offer a potential for new
and efficient options for impurity removal in aqueous solution, including alumina
refinery Bayer liquor. The lamellar structure of LDHs can be used for the
controlled removal of a variety of species. This is achieved through their ability to
adjust the separation of the hydroxide layers, and the reactivity of the interlayer
region. Hydrotalcite has a high selectivity for carbonate anions, making it
ineffective as an anion-exchange material unless further treatment is made.
Heating to 300 ºC causes decarbonation as the carbonate anion decomposes,
resulting in an amorphous material that will absorb anions and return to its
original hydrotalcite structure.
This investigation will look at characterising the hydrotalcite structure that forms
during the seawater neutralisation process of Bauxite refinery residues using
synthetic hydrotalcite and hydrotalcite formed from Bayer liquor. This
information will then be used to investigate hydrotalcite as a potential trigger for
reversion. Other components of red mud will then be investigated to establish
which compounds in red mud contribute to reversion. Once the cause for
reversion has been established methods for minimising reversion will be explored.
The final objective of this research will be to determine the potential of synthetic
hydrotalcite, Bayer hydrotalcite, and seawater neutralised red mud for the removal
of arsenate, vanadate, and molybdate from aqueous solutions.
45
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59
CHAPTER 2
Experimental methods and analysis
techniques
60
1. Introduction
The seawater neutralisation of bauxite refinery residue results in the precipitation
of hydrotalcite-like compounds (Mg, Al, and Ca). This process is very similar to
the co-precipitation method used for the synthesis of synthetic hydrotalcite
materials. Therefore, all synthetic hydrotalcite used in these investigations will be
synthesised by the co-precipitation method.
An experiment for the identification of compounds in bauxite refinery residues
that cause reversion (increase in pH and aluminium concentration after
neutralisation) has been developed and is described in this chapter. The synthesis
and characterisation of whewellite and hydrocalumite (candidates for causing
reversion) has also been discussed.
A number of analytical techniques have been utilised in this investigation,
including X-ray diffraction, vibrational spectroscopy, thermal analysis, and
elemental analysis. Hydrotalcites, red mud, and seawater neutralised red mud have
been analysed using these techniques to enable a complete characterisation of
these materials.
61
2. Experimental methods
2.1. Synthesis of hydrotalcite with different oxy-anions
The hydrotalcites were synthesised by the co-precipitation method, which utilises
the slow addition of a caustic solution containing the oxy-anion (1) to a mixed
metal solution (2). The concentrations of anions used are given in Table 2.1,
whilst the concentrations of cations used in this investigation are given in
Table 2.2. Solution 1 contains 2M NaOH and a combination of either: 1) Na2CO3
and Na2MoO4, 2) Na2CO3 and NaVO3, or 3) Na2CO3 and Na2HAsO4·7H2O for
a combined concentration of 0.2M, respectively. The mixed metal solutions for
each Mg:Al ratio were achieved by preparing the following solutions: 2:1
hydrotalcite required 0.66M Mg2+ (MgCl2·6H2O) and 0.33M Al3+ (AlCl3·6H2O),
the 3:1 hydrotalcite required 0.75M Mg2+ (MgCl2·6H2O) and 0.25M Al3+
(AlCl3·6H2O), whilst the 4:1 hydrotalcite required 0.80M Mg2+ (MgCl2·6H2O)
and 0.20M Al3+ (AlCl3·6H2
O).
The caustic solution (2M) was added drop wise to the mixed metal solution and
was stirred at 400 rpm to ensure dissolution. After caustic addition was complete
(final pH between 8.5 and 9.5), the mixture was stirred for two hours before the
solid product was isolated via vacuum filtration with a Whatman 542 filter paper.
The precipitate was washed twice with deionised water (250 mL washing) before
being vacuum dried and placed in an oven (85 °C) overnight to dry.
2.2. Synthesis of Bayer precipitate
The Bayer precipitate, containing hydrotalcite, calcite and halite, was prepared by
the addition of seawater (Inskip Point – QLD, Australia, 2008) to Bayer liquor
(Gove refinery Australia, 2008) at a volumetric ratio of 4.5:1. Bayer liquor refers
to the combination of 1 part red mud liquor (RML) and 0.9 parts supernatant
liquor (SNL). The compositions of the two Bayer liquors are provided in
Table 2.3. The solution was stirred thoroughly for 2 hours before being vacuum
filtered and dried overnight in an oven (85 °C). An average final pH between 8.5
and 9.5 was achieved after neutralisation. Three Bayer precipitates were prepared
62
Table 2.1: Concentrations of Na2CO3, Na2HAsO4·7H2O, NaVO3, and
Na2MoO4
Synthetic hydrotalcite
used to synthesise hydrotalcites with different oxy-
anions.
Na2CO3Oxy-anion
(concentration) (concentration)
HT(CO32- 0.20 M )
Na2CO Na3 2HAsO4·7H2
HT(CO
O
32-, AsO4
3- 0.10 M ) 0.10 M
HT(AsO43- - ) 0.20 M
Na2CO NaVO3
HT(CO3
32-, VO4
3- 0.10 M ) 0.10 M
HT(VO43- - ) 0.20 M
Na2CO Na3 2MoO
HT(CO4
32-, MoO4
2- 0.10 M ) 0.10 M
HT(MoO42- - ) 0.20 M
Table 2.2: Concentration and masses used to synthesis 2:1, 3:1, and 4:1
synthetic hydrotalcites.
Desired Ratio (Mg:Al) MgCl2·6H2 AlClO 3·6H2
Conc. (M)
O
Mass (g) Conc. (M) Mass (g)
2:1 0.67 136.211 0.33 79.672
3:1 0.75 152.475 0.25 60.358
4:1 0.80 162.640 0.20 48.286
Table 2.3: Composition of Bayer liquors, determined by Potentiometric
titration.
Alumina (g/L Al2O3 Caustic (g/L Na) 2 Carbonate (g/L NaO) 2
RML
O)
5.4 14.6 n/a
SNL 2.8 3.0 9.9
63
using the same experimental procedure, and the results presented in this report are
an average of the three.
2.3. Synthesis of synthetic Bayer precipitate
The same experimental procedure used for the synthesis of Bayer precipitate is
used, however, synthetic seawater and synthetic Bayer liquor were used. Synthetic
seawater is used for two reasons: 1) so the exact composition is known, and 2) to
eliminate contaminates such as organics and solid materials found in actual
seawater.
2.3.1. Synthetic seawater (SW)
Synthetic seawater was prepared from the following AR grade salts in Table 2.4,
which represents the average composition of seawater by 99.9%. [1] The
concentrations of the individual elements are also presented in Table 2.4.
Table 2.4: Salts used to prepare synthetic seawater and relative concentrations.
Salt g/L of salt Element ppm element
NaCl 23.926 Cl 19500
MgCl2·6H2 10.830 O Na 10770
Na2SO 4.008 4 Mg 1290
CaCl2·2H2 1.519 O S 905
KCl 0.667 Ca 412
NaHCO 0.196 3 K 380
KBr 0.098 Br 67
H3BO 0.026 3 Sr 8
SrCl2·6H2 0.024 O B 4.4
NaF 0.003 F 1.3
2.3.2. Synthetic supernatant liquor (SNL)
Synthetic SNL was prepared by first making synthetic strong evaporation liquor
(SEL) and then diluting SEL by a factor of 16. Supernatant liquor (SNL) is the
64
dilute liquor which remains after the waste slurry in the disposal dams have settled
out and compacted. It is essentially diluted SEL, which is the reason behind this
synthetic procedure. Table 2.5 outlines the characteristics of synthetic SNL.
Synthetic SEL was prepared by the addition of Na2CO3
(35 g/L) to a concentrated
NaOH (4M) solution. The solution was stirred for 10 minutes before 1-2 grams of
aluminium metal was added to the caustic solution, and allowed to dissolve
completely before the addition of another 1-2 grams of aluminium metal. This
reaction is exothermic, so the liquor can boil rapidly if the aluminium is added too
quickly. This process continued until all aluminium metal had been dissolved
(31 g/L), which took around 6 hours to complete.
Table 2.5: Alumina, caustic and carbonate concentration of synthetic SNL and
real SNL, determined by Potentiometric titration.
Al
(g/L Al2O3
Caustic
) (g/L Na2
Carbonate
O) (g/L Na2
SNL 2008
O)
(QRDC sample) 2.8 3.0 9.9
Synthetic SNL 3.2 6.5 6.3
2.4. Seawater neutralisation of red mud
All Bayer liquors, slurries and solids used in these investigations were provided
by QRDC from the Gove refinery in Australia, 2008. Seawater was collected from
Inskip Point, QLD, Australia, 2008.
Red mud slurry (RMS) was prepared by mixing 300.00 g of vacuum dried red
mud with 288.93 g of red mud liquor and agitating vigorously for 1 hour. RMS
(150 mL) was then added to SNL (135 mL) and allowed to stir at 400 rpm for 10
minutes. Seawater (1.282 L) was added slowly to the stirring slurry and the pH
change of the system was monitored using a general laboratory pH probe. After
seawater addition was complete, the mixture was allowed to stir for 2 hours before
the solid component was removed via vacuum filtration with a Whatman 542
filter paper. The solid was vacuum dried and then placed in an oven (85 °C)
overnight to dry.
65
2.5. Trigger experiments
The term trigger is given to a component of red mud proposed to cause reversion.
The experiment involved the addition of different concentrations of each trigger
(Table 2.6) to 60 mL of synthetic SNL, whilst being stirred for 5 minutes. After 5
minutes, synthetic seawater (270 mL) was added to SNL at around 120 mL a
minute, and was left to stir for a further 2 hours. The pH was monitored at 15
second intervals over the full 2 hour period. After this time the solution was
vacuum filtered and dried overnight in an oven (85 °C).
Samples (20 mL) were taken every 30 minutes to monitor the ions in solution, and
the concentration of phases in the precipitate. Each sample was vacuum filtered
through a Whatman 542 filter paper and a nylon syringe filtered (0.45 μm) for ICP
analysis, while the precipitate was placed in the oven to dry.
2.5.1. Trigger materials
Triggers that were AR grade materials include calcium hydroxide (Ca(OH)2) and
sodium carbonate (Na2CO3
). The following materials were provided by QRDC
and are samples from an alumina refinery: tricalcium aluminate (TCA), sodalite,
and gibbsite.
2.5.1.1. Synthesis of hydrocalumite – Ca2Al(OH)6Cl·2H2
O
The co-precipitation method, commonly used to prepare LDHs, was used for the
preparation of hydrocalumite. The co-precipitation method involved the addition
of two solutions, where solution 1 contained 2M NaOH and a combination of
Na2CO3 to give a concentration of 0.2M, while solution 2 contained 0.66M Ca2+
(CaCl2·2H2O) and 0.33M Al3+ (AlCl3·6H2O). Solution 2 was added drop wise to
solution 1, under vigorous stirring. The precipitated compound was then
thoroughly washed to remove any residual salts and dried overnight in an oven
(85 °C).
66
Table 2.6: Concentration and mass of each trigger in 60 mL of synthetic SNL.
Trigger
Concentration
TCA BHT Ca(OH) Sodalite 2 Whewellite Na2CO Hydrocalumite 3 Gibbsite
Masses required for desired concentration
0.005 0.080g 0.146g 0.291g - 0.032g - -
0.01M
-
0.160g 0.292g 0.582g - 0.088g 0.064g 0.168g 0.047g
0.02M 0.320g - - - - - 0.337g -
0.04M 0.640g - - - - - 0.674g -
0.05M 0.799g 1.462g 0.222g 2.908g 0.438g 0.318g 0.842g 0.234g
0.075M 1.198g - - - - - - -
0.10M 1.598g 2.924g 0.445g 5.815g 0.877g 0.636g 1.684g 0.468g
0.20M - - 0.889g - - - 3.368g -
0.30M - - 1.334g - - - 5.052g -
0.40M - - 1.778g - - - 6.737g -
0.50M - - 2.223g 4.383g - - 8.421g 2.340g
1.00M - - 4.447g - - - - -
67
The X-ray diffraction pattern confirmed the formation of hydrocalumite, along
with a small quantity of calcite (Fig. 2.1).
2.5.1.2. Synthesis of whewellite – CaC2O4·H2
O
Synthetic whewellite was prepared by adding calcium chloride (1M) drop wise to
a solution of oxalic acid (1M). The mixture was stirred continuously for 2 hours
before being vacuum filtered and washed with de-ionised water. The precipitate
was placed in an oven (85 °C) overnight to dry. The X-ray diffraction pattern of
synthetic whewellite showed no impurities (Fig. 2.2).
2.6. Thermal activation and treatment of aqueous solutions
Hydrotalcites prepared for the thermally activated study were crushed to a fine
powder before being placed in a furnace and heated at 20 °C per minute to
340 °C. The samples were held at 340 °C for 1 hour before being removed and
placed in a desiccator until the sample had cooled to ambient temperature. Once
cooled the thermally activated samples were re-weighed to calculate the
percentage mass-loss. All thermally activated materials (hydrotalcite and red mud)
had a similar mass loss between 7-9 %. The thermally activated samples were
stored in a vacuum desiccator, until required for re-hydration.
Four 100 ppm solutions containing arsenate, vanadate, molybdate (individual
solutions), and a solution containing arsenate, vanadate, and molybdate (combined
solution) were prepared using AR grade sodium salts of the desired anion. Each
100 ppm aqueous solution was diluted by means of serial dilution to yield 75, 50,
25 and 5 ppm anion solutions, using ultra pure water. The pH of all solutions was
maintained at around 8.5 before treatment with thermally activated hydrotalcite.
Thermally activated hydrotalcite (0.5 g) was added to each anionic solution
(10 mL) in a small beaker (50 mL), and was stirred at 800 rpm for 30 minutes.
The mixture was vacuum filtered. The solid residue was left in the oven at 85 °C
overnight to dry, while the aqueous component was re-filtered using a syringe
filter (0.45 μm) for ICP-OES analysis (requires a very low solids concentration).
68
Figure 2.1: XRD pattern of synthesised hydrocalumite
and the corresponding reference patterns.
Figure 2.2: XRD pattern of synthesised whewellite
and the corresponding reference pattern.
69
3. Characterisation techniques
Numerous instrumental techniques have been employed to characterise the
materials in this investigation, since multiple techniques provide a more complete
analysis of these materials. The following techniques have been used throughout
this investigation:
3.1. Inductively coupled plasma optical emission spectrometry
Syringe filtered (0.45 μm) solutions were analysed neat (samples were not
diluted), due to the low concentration of anionic species in solution. Four
standards (blank-ultra pure water, 25, 50, 100 ppm) containing aluminium,
calcium, arsenate, vanadate, and molybdate were prepared in a chloride matrix.
These standards were used to prepare a calibration curve. The concentration of
each element in solution was obtained using an integration time of 3 seconds with
3 replications (1 overall replication). To ensure quality control, a sample with
known concentration of all elements (prepared using AR grade materials) was
analysed. The solution was analysed three times, with concentrations of samples
reported throughout this investigation being the average of three overall
replications. The relative amounts of each element was recorded on a Varian
Liberty 2000 ICP–OES at wavelengths of 394.400, 279.553, 393.366, 311.837,
202.032 and 188.980 nm for aluminium, magnesium, calcium, vanadium,
molybdenum, and arsenic respectively. This is an elemental technique, therefore,
in this discussion it is assumed that all arsenic, vanadium, molybdenum, and
sulfur detected by ICP is due to the concentration of the corresponding anionic
compounds arsenate (AsO43-), vanadate (VO4
3-), molybdate (MoO42-), and sulfate
(SO42-
).
70
3.2. X-ray diffraction
X-ray diffraction patterns were collected using a Philips X'pert wide angle X-ray
diffractometer, operating in step scan mode, with Cu Kα
radiation (1.54052 Å),
and parallel beam. Patterns were collected in the range 3 to 90° 2θ with a step size
of 0.02° and a rate of 30s per step. Samples were prepared as a finely pressed
powder into aluminium sample holders. Thin films, using Vaseline, were used for
experiments with minimal yields, such as the trigger experiments.
3.3. Spectroscopy
3.3.1. Fourier-transform infrared spectroscopy
Infrared spectra were obtained using a Nicolet Nexus 870 FTIR spectrometer
with a smart endurance single bounce diamond ATR cell. Spectra over the
4000-525 cm-1 range were obtained by the co-addition of 128 scans with a
resolution of 4 cm-1
and a mirror velocity of 0.6329 m/s.
3.3.2. Fourier Transform Raman spectroscopy
The Fourier Transform Raman spectroscopy (FT-Raman) analyses were
performed on powder samples pressed in a sample holder suitable for the Perkin
Elmer System 2000 Fourier transform spectrometer. The spectrometer was
equipped with a Raman accessory comprising of a Spectron Laser Systems SL301
Nd:YAG laser operating at a wavelength of 1064 nm. Spectra were taken in the
wavenumber range between 525 and 3800 cm-1
.
3.3.3. Raman microspectroscopy
The crystals of hydrotalcite were placed on the stage of an Olympus BHSM
microscope, equipped with 10x and 50x objectives and are part of a Renishaw
1000 Raman microscope system, which also includes monochromators, a filter
system and a Charge Coupled Device (CCD). Raman spectra were excited by a
71
HeNe laser (633 nm) at a nominal resolution of 2 cm-1 in the range between 100
and 4000 cm-1. Repeated acquisition using the highest magnification was
accumulated to improve the signal to noise ratio. Spectra were calibrated using the
520.5 cm-1
line of a silicon wafer.
3.3.4. Band component analysis
Spectral manipulation such as baseline correction, smoothing and normalisation
was performed using the GRAMS® software package (Galactic Industries
Corporation, Salem, NH, USA).
Band component analysis was undertaken using the Jandel ‘Peakfit’ software
package, which enabled the type of fitting function to be selected and allows
specific parameters to be fixed or varied accordingly. Band fitting was undertaken
using a Lorentz- Gauss cross-product function with a minimum number of
component bands used for the fitting process. The Lorentz- Gauss ratio was
maintained at values greater than 0.7 and fitting was undertaken until reproducible
results were obtained with squared correlations of r2
3.4. Thermal analysis
greater than 0.995.
3.4.1. Thermogravimetric analysis
Thermal decomposition of the hydrotalcites were carried out in a TA® Instrument
incorporated with a high-resolution thermogravimetric analyser (series Q500) in a
flowing nitrogen atmosphere (80 cm3
/min). Approximately 50 mg of sample was
heated in an open platinum crucible at a rate of 2.0 °C/min up to 1000 °C. The
TGA instrument was coupled to a Balzers (Pfeiffer) mass spectrometer for gas
analysis. Only selected gases such as water and carbon dioxide were analysed.
The synthesised hydrotalcites were kept in an oven for 24 hrs before TG analysis.
Thus the mass losses were calculated as a percentage on a dry basis.
72
3.4.2. Dynamic experiment
Thermal decomposition of the hydrotalcites were carried out in a Derivatograph
PC type thermoanalytical instrument (Hungarian Optical Works, Budapest,
Hungary), capable of recording the thermogravimetric (TG), derivative
thermogravimetric (DTG) and differential thermal analysis (DTA) curves
simultaneously. The sample was heated in a ceramic crucible in static air
atmosphere at a rate of 5 °C /min.
3.4.3. Controlled rate thermal analysis experiment
Thermal decomposition of HT(CO32-) was carried out in a Derivatograph PC-type
thermoanalytical instrument (Hungarian Optical Works, Budapest, Hungary)
under static air at a pre-set, constant decomposition rate of 0.10 mg/min. (Below
this threshold value the samples were heated under dynamic conditions at a
uniform rate of 1.0 °C/min). The samples were heated in an open ceramic crucible
at a rate of 1.0 °C/min-1
up to 900 °C. With the quasi-isothermal, quasi-isobaric
heating program of the instrument the furnace temperature was regulated precisely
to provide a uniform rate of decomposition in the main decomposition stage.
3.5. Electron dispersive X-ray spectroscopy
Electron dispersive X-ray microanalysis (EDX) of samples involved coating the
samples with a thin layer of evaporated carbon for conduction and examined in a
JEOL 840A analytical SEM (JEOL Ltd, Tokyo, Japan) at 25 kV accelerating
voltage. The instrument had been standardised with a set of standards before the
analysis of the hydrotalcite and red mud samples. Microanalysis of the clusters of
fine crystals was carried out using a full standards quantitative procedure on the
JEOL 840 SEM using a Moran Scientific microanalysis system (Tokyo, Japan).
Oxygen was not measured directly but calculated using assumed stoichiometries
to the other elements analysed.
73
3.6. Potentiometric titration
Bayer liquors were analysed using a 815 Metrohm Fully-automated
Potentiometric Titrator for the determination of carbonate, caustic and alumina
content. Samples (40mL) were not diluted due to the low carbonate content in the
liquors.
4. References
[1] R.B. Heslop, P.L. Robinson, Inorganic Chemistry, Elsevier, London, UK, 1961.
74
CHAPTER 3
Synthesis and characterisation of synthetic
hydrotalcites
75
1. Introduction
Hydrotalcite-like clays (more commonly known as layered double hydroxides –
LDHs) are based on the brucite structure, Mg(OH)2, where each Mg2+ ion is
octahedrally surrounded by six OH- ions. These structures crystallise in a layer
type lattice due to the relatively small twofold positively charged cations in close
proximity to the non-spherosymmetrical and highly polarisable OH- ions. [1] The
hydrotalcite structure is obtained when the Mg2+ ions are replaced by trivalent
cations of similar radius, such as Al3+
. [2] The higher charge of the aluminium
cations causes an overall positive charge on the hydroxyl layer. These layers are
maintained electrically neutral by charge compensating interlayer anions and
water molecules. The number, length, orientation and strength of the bonds
between the anions and the cationic surface all influence the thickness of the
interlayer. [3] The hydroxyl layers are stacked upon one another and are held
together by weak interactions through the hydrogen atoms. [4]
The general formula given to LDHs is [M2+1-x M3+
x(OH)2]x+Am-x/m·nH2O, where
M2+ is a divalent cation, M3+
is trivalent cation and 'A' is an interlamellar anion
with charge m-. Pure LDH phases exist for 0.2 ≤ x ≤ 0.33. Values outside the
specified x range will form:
i) boehmite (α-AlOOH) for x > 0.337,
ii) hydromagnesite (4MgCO3·Mg(OH)2·4H2
iii) a mixture of hydromagnesite and Mg(OH)
O) for 0.105 < x < 0.201, and
2
for x < 0.105. [5]
Hydrotalcite is produced when M2+ = Mg2+ and M3+ = Al3+, giving the general
formula Mg6Al2(OH)16CO3·4H2O. There are numerous ways to synthesis LDHs
including co-precipitation (most common), hydrothermal synthesis,
electrochemical methods, hydrolysis methods and urea reduction. [3, 4, 6-8] Co-
precipitation can be carried out at high or low supersaturation, the difference
being that the nucleation rate is much higher than crystal growth at high
supersaturation. However, low supersaturation results in more crystalline
structures, [2, 8] while high supersaturation produce higher yields with lower
76
crystallinity. The hydrotalcite materials synthesised in this investigation utilise
high supersaturation.
This chapter characterises hydrotalcite using a range of analytical techniques
(vibrational spectroscopy, X-ray diffraction, and thermal analysis) to obtain a
better understanding of the materials. It also explores the intercalation
characteristics of arsenate, vanadate, and molybdate into hydrotalcite of variable
cationic ratio. The decomposition temperature has been used to investigate the
thermal stability of the hydrotalcite materials. Using these techniques in
combination has enabled the mechanism for inclusion (intercalation/adsorption) to
be determined. The synthetic hydrotalcites were prepared using the same seawater
neutralisation process as that used to treat bauxite refinery residues. This ensured
uniformity in synthetic procedures for both synthetic and ‘Bayer’ hydrotalcite is
maintained.
77
2. Infrared and Raman spectroscopy
Insight into the unique structure of hydrotalcites has been obtained using a
combination of infrared and Raman spectroscopy, through the identification of
unique band positions of the hydroxyl-stretching units of Mg-OH and Al-OH,
water vibrational modes, and carbonate vibrational modes.
Water plays a unique role in the stabilisation of the hydrotalcite structure.
Hydrotalcites are layered anionic clays, where the positive layer charge is
balanced by the incorporation of anions. The neutralisation of the layer charge by
the incorporation of anionic species, along with a complex network of hydrogen
bonding involving water, the cationic surface and anions renders the hydrotalcite
structure stable. The position and intensity of the vibrational spectroscopic bands
in the hydroxyl-stretching region indicates that water is highly structured. The
position of the bands in the hydroxyl deformation region of the infrared spectrum
supports the concept of structured water between the hydrotalcite layers.
2.1. Hydroxyl stretching region
The Raman and infrared spectra of the synthesised hydrotalcites show a broad,
intense band centred at approximately 3550 and 3400 cm-1 due to the stretching
modes of hydroxyl groups in the LDH layers and water molecules (Fig. 3.1). Band
assignments have been based on the work completed by Rives [9] and Farmer
[10]. Band component analysis was used to help identify the different hydroxyl
species, however the assignment of the bands is difficult because of the complex
band profile and numerous overlapping bands. The bands at lower wavenumbers
(3200 -2800 cm-1
) are attributed to strongly hydrogen bonded water molecules to
interlayer anions. Carbonate is generally the interlayer anion, which is strongly
hydrogen bonded to interlayer water. This is generally only observed in the
infrared spectrum as water is a weak scatterer in Raman spectroscopy.
Bands situated around 3200 to 3450 cm-1 (infrared and Raman) are attributed to
the hydroxyl stretching mode of water co-ordinated to the cationic hydroxyl
surface of the hydrotalcite, while bands at 3400 to 3500 cm-1 (infrared and
78
Figure 3.1: Raman and infrared spectra of carbonate hydrotalcite in the hydroxyl stretching vibrational region.
79
Raman) are assigned to water hydrogen bonded to other water molecules in the
hydrotalcite interlayer space. Bands between 3500 and 3600 cm-1 (infrared and
Raman) are attributed to Al-OH stretching vibrations, whilst bands above
3600 cm-1 (infrared and Raman) are attributed to Mg-OH. Multiple bands in this
higher wavenumber region are attributed to water hydrogen bonded to M3OH
units (where M might be Mg or Al and any combinational permutation of these
metals). This is particularly noticeable for the M3OH units, where Mg-OH
vibrations are observed at 3698 and 3602 cm-1
, Raman and infrared respectively.
2.2. Carbonate vibrations
The structure of hydrotalcite depends on the balance of the positive surface
charges on the cationic hydroxyl surface by the negative charges of intercalated
anions. This investigation looks at the following intercalated anions: 1) carbonate,
2) carbonate/vanadate, 3) carbonate/arsenate, and 4) carbonate/molybdate
mixtures. The unperturbed carbonate ion is trigonal planar with point symmetry
D3h. Group theoretical analysis of the carbonate ion predicts four normal modes;
the ν1 symmetric stretch of A1 symmetry normally observed at 1063 cm-1, the
antisymmetric stretch of E’ symmetry observed at 1415 cm-1, the ν2 out of plane
bend at 879 cm-1, and the in-plane bend at 680 cm-1. [10] All modes are both
Raman and infrared active except for the ν2
mode, which is IR active only. This
information is summarised in Table 3.1.
Table 3.1: CO32-
Wavenumber (cm
bands. [10]
-1 Vibrational mode )
1063 Symmetric stretch ν1 CO3
1415
2-
antisymmetric stretch ν3 CO3
879
2-
out-of-plane bending ν2 CO3
680
2-
in-plane bending ν4 CO3
2-
80
Figure 3.2: Infrared spectra of the synthesised hydrotalcites,
containing arsenate, in the carbonate vibrational region.
Figure 3.3: Infrared spectra of the synthesised hydrotalcites,
containing vanadate, in the carbonate vibrational region.
HT(CO32-)
HT(CO32-,VO4
3-)
HT(VO43-)
81
The infrared spectra of the CO32- antisymmetric stretching region shows three
bands observed at around, 1360, 1390, and 1485 cm-1, for all the hydrotalcites
synthesised (Fig. 3.2-3.4). One possible interpretation is that these bands are
attributed to the carbonate anion in three different environments: a) free carbonate
anions, b) water hydrogen bonded to the carbonate, and c) carbonate bonded to
the cationic hydroxyl surface. Bands at 1485 cm-1 are attributed to the free
carbonate anion or carbonate adsorbed on the external surfaces of the layers.
Bands at around 1390 cm-1 are assigned to carbonate bonded to water in the
hydrotalcite interlayer, while the bands at around 1360 cm-1
are assigned to
carbonate bonded to the hydroxyl surface of the hydrotalcite.
The Raman spectra show the typical sharp band of the ν1 symmetric stretching
mode of the carbonate anion at 1060 cm-1 (Fig. 3.5-3.7). The spectra give a clear
representation of the amount of carbonate that enters the system through external
sources during a 2 hour synthesis period, indicated by a band at 1060 cm-1 for
samples with no carbonate in the synthesis procedure. These samples include
HT(AsO43-), HT(VO4
3-) and HT(MoO42-
).
A low intensity band is observed at 695 cm-1 and is attributed to the ν4 bending
mode of the carbonate anion. The infrared spectrum of HT(CO32-) at lower
wavenumbers displays the ν2 out of plane bending of carbonate at around
860 cm-1 and the in-plane bend at around 630 cm-1
. The presence of these bands
suggests that the carbonate anions are distorted.
2.3. Water OH deformation vibrations
Minerals and synthetic hydrotalcites containing physically adsorbed water give a
strong water deformation mode at around 1640 cm-1 (Fig. 3.2-3.4). The position of
the band is influenced by the amount of adsorbed water, the mineral type and the
exchangeable anion to which the water is bonded. This study was concerned with
liquid water, and therefore bands at 3455 and 1645 cm-1 are expected. These
bands are observed for all hydrotalcites synthesised. As the carbonate content in
the synthesised hydrotalcites decreased, the band position moved to lower
wavenumbers. The bands that occur at wavenumbers around 1645 cm-1 are
82
Figure 3.4: Infrared spectra of synthesised hydrotalcites,
containing molybdate, in the carbonate vibrational region.
Figure 3.5: Raman spectra of the synthesised hydrotalcites,
with arsenate, in the carbonate vibrational region.
HT(CO32-,MoO4
2-)
HT(CO32-)
HT(MoO4
2-)
83
indicative of water which is strongly hydrogen bonded. Such water molecules
may be hydrogen bonded to the cationic hydroxyl surface or to adjacent water
molecules. The bands that occurred at lower wavenumbers (around
1635 cm-1) for HT(VO43-) and HT(MoO4
2-
) suggest that the water molecules are
not as tightly bound as a result of a lower carbonate concentration.
2.4. Vibrations associated with arsenate
The infrared and Raman spectra of selected minerals containing arsenate have
been published by Farmer. [10] The observed vibrational bands for arsenate and
vanadate minerals are given in Table 3.2. There are four vibrations for vanadate,
namely the A1 symmetric stretching mode observed between 810 and 840 cm -1,
the E’ bending mode in the region at around 345 cm-1, the F2 antisymmetric
stretching mode between 810 and 878 cm-1, and the F2 bending mode between
398 and 463 cm-1. The F2 modes are both Raman and infrared active, whereas the
A1
and E’ modes are Raman active only.
The As-O bands associated with arsenate in the hydrotalcite structures are
observed between 900 and 700 cm-1 (Fig. 3.5). The peak maximum is situated at
around 820 cm-1 for both HT(CO32-, AsO4
3-) and HT(AsO43-). The spectrum of
the arsenate-only hydrotalcite appears to be composed of three bands: 829, 816,
and 780 cm-1. It is proposed that two of the bands are absent from HT(CO32-,
AsO43-
) due to the lower concentration of arsenate in the initial solution.
84
Figure 3.6: Raman spectra of the synthesised hydrotalcites,
with vanadate, in the carbonate vibrational region.
Figure 3.7: Raman spectra of synthesised hydrotalcites,
containing molybdate, in the hydroxyl stretching region.
HT(VO43-)
HT(CO32-,VO4
3-)
HT(CO32-)
HT(MoO42-)
HT(CO32-, MoO4
2-)
HT(CO32-)
85
Table 3.2: VO43- and AsO4
3-
bands from different sources. [10]
ν1 (A1
symmetric
stretch
)
(cm-1
(Raman active)
)
ν2
bend
(E’)
(cm-1
(Raman
active)
)
ν3 (F2
antisymmetric
stretch
)
(cm-1
(IR and Raman
active)
)
ν4 (F2
bend
)
(cm-1
(IR and
Raman
active)
)
VO4
824
3-
340 790 340
827 340 780 340
870 328 825 480
874 345 855 345
824 305 790 340
AsO4810 3-
342 810 398
837 349 878 463
2.5. Vibrations associated with vanadate
The infrared spectra of selected minerals containing pentavalent vanadium have
been published by Farmer [10] (Table 3.2). There are four vibrations for vanadate,
namely the A1 symmetric stretching mode observed between 824 and 874 cm -1,
the E’ bending mode in the region between 305 and 345 cm-1, the F2
antisymmetric stretching mode between 780 and 855 cm-1, and the F2 bending
mode between 340 and 345 cm-1. The F2 modes are both Raman and infrared
active, whereas the A1
and E’ modes are Raman active only.
The Raman spectra of the V-O bond region, 800 to 1000 cm-1, are shown in
Fig. 3.6. The absence of a peak in HT(CO32-) confirms that the band between 800
and 1000 cm-1 is indeed due to the intercalated vanadate anions. Bands at around
820 cm-1 are attributed to the A1 stretching modes of vanadate. The E’ bending
mode of vanadate in the Raman spectrum can only be seen in HT(VO43-) at
317 cm-1 (Fig. 3.9). The bands at around 340 cm-1 are attributed to the F2
bending
modes, and are observed for both vanadate hydrotalcites.
86
Figure 3.8: Raman spectra of the cation deformation
modes of arsenate containing hydrotalcites.
Figure 3.9: Raman spectra of the cation deformation
modes of vanadate containing hydrotalcites.
87
The F2 antisymmetric stretching modes for both the vanadate containing
hydrotalcites are observed at 885 cm-1 (Fig. 3.6), which is higher than bands
reported by Farmer. [10] A broad band at around 220 cm-1 appears only in the
vanadate containing hydrotalcites and therefore is possibly due to the ν4
bending
mode of the vanadate anion.
The hydrotalcites were synthesised at high pH, 8.5 to 10.5, therefore vanadate
may be present as a number of species H2VO4-, HVO4
2-, and VO43-. It is also
possible pyrovanadate anionic species are present in solution like VO3(OH)2-,
HV2O73-, and V2O7
4-
. [9] However, the sizes of these pyrovanadates would
hinder the intercalation of these anions into the hydrotalcite structure.
2.6. Vibrations associated with molybdate
The Raman spectra of the synthetic hydrotalcites with carbonate and molybdate in
the 750 to 1150 cm-1 region are shown in Fig. 3.7. The Raman spectra clearly
show the CO32- ν1 symmetric stretching modes centred upon 1060 cm-1 and the
MoO42- ν1 symmetric stretching modes centred upon 908 cm-1. [9, 10] The two
molybdate bands at 894 and 907 cm-1 are assigned to molybdate hydrogen bonded
to the interlayer water molecules and the molybdate anions chemically bonded to
the hydrotalcite hydroxyl surface, respectively. [10] These bands are observed in
both the mixed anion hydrotalcite and in the hydrotalcite synthesised with
molybdate only. The broad band at 854 cm-1 may be assigned to the molybdate ν3
antisymmetric stretching mode.
2.7. Cation deformation vibrations
Raman spectra of the hydrotalcites in the 600-100 cm-1 region (Fig. 3.8-3.10)
show two bands around 470 and 549 cm-1, which are attributed to the Mg-O and
Al-O symmetric stretching vibrations. [9] This assignment was based on the
intensities of the bands being the same, representing the same Mg:Al ratio of the
three hydrotalcites, as well as the same bands appearing in all synthesised
88
Figure 3.10: Raman spectra of the cation deformation
modes of molybdate containing hydrotalcites.
89
hydrotalcites. Raman spectra in the lower wavenumber region, below 250 cm-1,
are complex and consist of overlapping bands. These bands are not discussed in
detail, however, they are attributed to metal-oxygen bonds, lattice vibrations and
hydrogen bonds. A broad antisymmetric band is observed at around 153 cm-1
which may be resolved into component bands. This band is common for all three
hydrotalcites and is probably a hydrogen bond stretching vibration involving the
hydrotalcite OH units and water in the interlayer. [11]
3. Effect of pH and synthesis time on the intercalation/adsorption of
arsenate and vanadate from aqueous solutions
Due to the low intercalation/adsorption percentages and concentration of
molybdate in Bayer liquor, the removal of molybdate will not be included in this
investigation. This investigation will focus on the removal of arsenate and
vanadate in different synthesis conditions. The following results will look at the
formation of hydrotalcite from five solutions:
1) carbonate (0.1M) and vanadate (0.1M),
2) carbonate (0.1M) and arsenate (0.1M),
3) vanadate (0.2M),
4) arsenate (0.2M), and
5) carbonate (0.67M), vanadate (0.67M), and arsenate (0.67M).
These hydrotalcites will be referred to as:
1) HT(CO32-,VO4
3-),
2) HT(CO32-,AsO4
3-
3) HT(VO
),
43-
4) HT(AsO
),
43-), and
5) HT(CO32-,AsO4
3-,VO43-
).
3.1. Effect of synthesis pH
In order to study the influence of the synthesis pH on the ability of hydrotalcite to
remove arsenate and vanadate from solution, hydrotalcites of formula
Mg6Al2(OH)16·(An-)·xH2O were made at pH 8, 10 and 13 and were allowed to stir
90
Figure 3.11: Percentage of anions removed from solution during the synthesis of
hydrotalcites at pH 8 (green), pH 10 (blue), and pH 13 (black).
VO4
3- HVO42- H2VO4
pH: 14-13 pH: 12-10 pH:9 - 4
-
Figure 3.12: Molecular shape of the vanadate anion in the pH range 7-14.
AsO4
3- HAsO42-
pH: 14-11 pH: 10-7
Figure 3.13: Molecular shape of the arsenate anion in the pH range 7-14.
91
for 2 hours (Fig. 3.11). Increasing the synthesis pH to 13 caused a significant
reduction in the removal of arsenate and vanadate from solution. At pH 8,
essentially both anions are completely removed from solution, whilst at pH 13
only 65-85 % are removed. A slight decrease in the amount of arsenate and
vanadate removed from solution is observed at pH 10. The removal percentage of
arsenate and vanadate remained greater than 85 %, except for
HT(CO32-,AsO4
3-,VO43-), which only showed a vanadate removal percentage of
80 %. Vanadate has the lowest affinity (compared to carbonate and arsenate), and
therefore is the most vulnerable to exchange reactions involving increased OH-
concentration. The decrease in removal percentages at pH 13 is due to excess OH-
anions in solution, which compete strongly with arsenate and vanadate anions for
the hydrotalcite interlayer.
Hydrotalcites with arsenate intercalated into the structure observed a much larger
decline in percentage removal (~10 %) when synthesised at pH 10, compared to
the vanadate hydrotalcites (~3 %). At pH 8 and 10 arsenate is present in solution
as the HAsO42- anion, while the vanadate anion exists as H2VO4
- at pH 8 and as
HVO42-
at pH 10 (Fig. 3.12 and 3.13). The increased charge and decrease in size,
increases the vanadate anions affinity for the interlayer, allowing it to compete
more strongly for the interlayer with the increased hydroxide concentration. The
arsenate anion on the other hand still has a negative 2 charge at pH 10. Therefore,
its affinity is unchanged, which makes it more vulnerable to the increased
hydroxide concentration.
3.2. Chemical stability of hydrotalcites synthesised over a 2, 24, and 48 hour
period
Fifteen hydrotalcites were prepared using a solution with pH 8, using five
different anion mixtures, and allowing the hydrotalcites to age for 2, 24, and 48
hours. The hydrotalcites are summarised in Table 3.3.
The hydrotalcites synthesised with carbonate, arsenate, and vanadate are exposed
to solutions at pH 10 and 14 to determine the chemical stability of the
intercalated/adsorbed anions. ICP analysis showed the percentage of anions that
92
were re-dissolved back into solution through anion exchange reactions
(Table 3.4). After the hydrotalcite has been exposed to two different alkaline
solutions, the percentage of anions in solution indicates the stability of the anions
in the interlayer region. The lower the dissolution percentage, the more stable the
hydrotalcite structure.
Table 3.3: Fifteen 3:1 hydrotalcites prepared at pH 8 and aged for 2, 24, and
48 hours.
2 hours 24 hours 48 hours HT(CO3
2-,VO43- HT(CO)-2h 3
2-,VO43- HT(CO)-24h 3
2-,VO43-
HT(VO
)-48h
43- HT(VO)-2h 4
3- HT(VO)-24h 43-
HT(CO
)-48h
32-,AsO4
3- HT(CO)-2h 32-,AsO4
3- HT(CO)-24h 32-,AsO4
3-
HT(AsO
)-48h
43- HT(AsO)-2h 4
3- HT(AsO)-24h 43-
HT(CO
)-48h
32-,AsO4
3-,VO43- HT(CO)-2h 3
2-,AsO43-,VO4
3- HT(CO)-24h 32-,AsO4
3-,VO43-
)-48h
The results indicate that the chemical stability of the hydrotalcites is relatively
high, with the majority of anions remaining in the hydrotalcite interlayer at pH 10.
Hydrotalcites prepared with an aging time of 2 hours had the lowest anion
stability in alkaline solutions.
Table 3.4: Percentage dissolution of hydrotalcites formed over varying
synthesis periods in NaOH at pH 10 and pH 14. Note the results for
the mixed anion hydrotalcite are for the anion in bold.
% dissolution % dissolution
NaOH pH 10 2h 24h 48h NaOH pH 14 2h 24h 48h
HT(CO32-,VO4
3- 0.1 ) 0.3 0.5 HT(CO32-,VO4
3- 46.8 ) 23.2 18.9
HT(VO43- 0.1 ) 0.1 0.4 HT(VO4
3- 41.9 ) 24.7 6.0
HT(CO32-,AsO4
3- 0.6 ) 0.1 0.0 HT(CO32-,AsO4
3- 44.0 ) 4.9 3.4
HT(AsO43- 0.3 ) 0.1 0.0 HT(AsO4
3- 41.2 ) 4.2 3.8
HT(CO32-,AsO4
3-,VO43- 5.3 ) 0.9 0.1 HT(CO3
2-,AsO43-,VO4
3- 53.8 ) 16.2 7.3
HT(CO32-,AsO4
3-,VO43- 3.7 ) 1.0 0.5 HT(CO3
2-,AsO43-,VO4
3- 46.2 ) 19.6 14.3
93
3.2.1. pH 10
The following hydrotalcites were resilient to pH 10: HT(CO32-,VO4
3-),
HT(CO32-,AsO4
3-), HT(VO43-), and HT(AsO4
3-). The percentage of anions
leached back into solution is insignificant. Therefore, arsenate and vanadate in the
hydrotalcite interlayer do not undergo exchange reactions with OH- ions at pH 10.
However, the mixed anion hydrotalcite aged for 2 hours HT(CO32-,AsO4
3-,VO43-
)
did exhibit a 5.3 and 3.7 % release of arsenate and vanadate, respectively, back
into solution. A limit to the number of intercalation sites in the interlayer forces
some of the anions to be adsorbed on the external surface. The adsorbed arsenate
and vanadate anions are more susceptible to exchange reactions involving
hydroxide ions. Minimal loses are observed for the mixed hydrotalcite aged for 24
and 48 hours.
It is proposed that a higher percentage of arsenate and vanadate anions are
adsorbed on the external surface of hydrotalcites aged for 2 hours than after 24
and 48 hours. Due to competition between these anions in solution, during
synthesis, only the anions with the highest affinity would be intercalated initially.
Therefore, the lower affinity anions (arsenate and vanadate) get adsorbed on the
external surface of the structure initially. Over longer periods of time, these low
affinity anions can migrate into the hydrotalcite interlayer as the hydroxyl metal
layers rearrange to form more aligned and ordered structures. The alignment of
the layers is believed to increase the interlayer distance and the number of
intercalation sites, thus removing more anions from solution. As the intercalation
of anions in the hydrotalcite interlayer is more stable, less anionic species will be
released back into solution for hydrotalcites synthesised with greater aging times.
94
Figure 3.14: Raman spectrum in the anionic stretching region,
1200-600 cm-1
, for hydrotalcites prepared for 2 hours at pH 8.
95
3.2.2. pH 14
Exposing the fifteen hydrotalcites to a pH 14 solution significantly increased the
percentage of arsenate and vanadate anions released back into solution. The
hydrotalcites aged for 2 hours, showed the lowest interlayer stability, with the
release of over 40 % of arsenate and vanadate. The large influx of hydroxide ions
competes very strongly with arsenate and vanadate in the hydrotalcite interlayer,
forcing arsenate and vanadate out of the interlayer via exchange reactions. Again
the disordered nature of the hydrotalcites after 2 hours contributes to the high
removal of arsenate and vanadate. Hydrotalcites containing arsenate showed
lower dissolution percentages when synthesised over longer periods of time. This
is attributed to the slightly higher affinity of the arsenate anion compared with
vanadate. The vanadate hydrotalcites synthesised over a 24 hour period also
showed a large reduction in percentage dissolution. As the aging time increased to
48 hours, less vanadate anions were released back into solution. This increase in
stability is due to: 1) the re-arrangement of anions to form a more ordered
structure, and 2) a network of hydrogen bonding involving the three anions and
interlayer water.
3.3. Raman spectra of hydrotalcites synthesised
3.3.1. Carbonate vibrational region (1200-600 cm-1
)
Band positions observed for carbonate, arsenate, and vanadate anions are given in
Tables 3.1 and 3.2. Four Raman bands, attributed to the carbonate anion, are
observed in HT(CO32-)-2h-pH8 between 1200 and 600 cm-1 (Fig. 3.14). These
bands are observed at 1084, 1061, 1058, and 1030 cm-1, and are assigned to the
symmetric stretching modes of carbonate. The presence of four bands in this
region suggests that the carbonate anion is in different environments (slightly
different bonding of the CO32- anion). It is proposed that carbonate is bonded to
H2O in the interlayer or other anions in the hydrotalcite interlayer (1061 and
1058 cm-1), and also bonded to the external surface of the hydrotalcite structure
(1030 cm-1). [12]
96
Figure 3.15: Raman spectrum in the anionic stretching region,
1200-600 cm-1
, for hydrotalcites prepared for 48 hours at pH 8.
97
The Raman spectra of the HT(CO32-)-48h-pH8 shows three bands at 1062, 1052
and 971 cm-1 (Fig. 3.15) compared to the HT(CO32-)-2h-pH8, which showed four
bands. The shift of the Raman band at 1030 to 971 cm-1
suggests that the
carbonate is non hydrogen bonded or only weakly hydrogen bonded, and is
possibly acting as a space-filler in the hydrotalcite interlayer.
The Raman spectrum for HT(VO43-)-2h-pH8 hydrotalcite (Fig. 3.14) exhibited
four broad bands at 939, 907, 879 and 815 cm-1, due to the A1 stretching modes of
V-O. [13] The multiple V-O vibrational modes are believed to be due to different
bonding strengths of the V-O bond (i.e. it’s in different environments). It is
proposed that the vibrational band at 939 cm-1 is due to the V-O symmetric
stretching mode of the tetrahedral vanadate anion. At pH 8, vanadate is most
likely present as H2VO4-. The symmetry of the original tetrahedral vanadate
structure is slightly obscured by the hydrogen atoms bonded to O-, and thus will
show a shift to lower wavenumbers. It is proposed that the lower wavenumber
band at around 815 cm-1 is due to V-OH bonds, which has a weaker bond strength
compared to the other V-O bonds in the structure. Within the hydrotalcite
structure there may be numerous bands (slightly different) due to this vibration,
and this is shown by the broadness of the V-OH band at 815 cm-1. The bands at
907 and 879 cm-1 are believed to be due to the V-O- stretching modes, which are
bonded to interlayer water, other anionic species, or to the hydroxyl surface of the
hydrotalcite. Synthesis of the same hydrotalcites but at pH 13 (Fig. 3.16) showed
a significant increase in the intensity and broadness of the band at around
870 cm-1. This indicates an increase in the number of V-O- bonds, which
correspond with the change in vanadate speciation as the pH increases to 13 (pH
8: H2VO4- and pH 13: VO4
3-). The presence of a broad band at 835 cm-1, (V-OH
stretching mode), suggests that there is still HVO42-
anions present at pH 13.
The Raman spectrum for HT(AsO43-)-2h-pH8 (Fig. 3.14) exhibited four bands at
911, 876, 841 and 808 cm-1, due to A1 stretching modes of As-O. As mentioned
previously, for the corresponding vanadate hydrotalcite, these different vibrational
modes are attributed to different bonding strengths of the As-O bond. It is
proposed that the strongest bond of the tetrahedral arsenate anion is As=O, and is
assigned to the 876 cm-1 band. This band is assigned to the As=O symmetric
98
Figure 3.16: Raman spectrum in the anionic stretching region,
1200-600 cm-1, for hydrotalcites prepared for 2 hours at pH 13.
99
stretching mode. At pH 8, arsenate is predominantly present as HAsO42-. It is
proposed that the very broad band at 841 cm-1 is attributed to the symmetric
stretch of As-O-, while at lower wavenumbers, 808 cm-1, is due to As-OH
symmetric stretching vibrational modes. The broadness of the band at 841 cm-1 is
proposed to be due to the overlapping of multiple As-O- bands, which may be
bonded in slightly different ways to other species in the interlayer (water or
cationic surface), which will result in a shift in band position, and thus making the
band appear broad. The relative areas under the bands at 876 and 808 cm-1
indicates that there are approximately the same number of As-OH bonds present
in the hydrotalcite structure as there are As=O bonds, which suggests that the
anion is in the HAsO42- form (1:1 ratio of As-OH : As=O). Increasing the pH of
solution to 13 results in AsO43- being the primary arsenate species. This is clearly
visible in Fig. 3.16, as an extremely broad band at 830 cm-1, assigned to the AsO-
vibration. The intensity of this band indicates a large number of multiple bands
that are very similar, such as the three As-O- bonds in the AsO43-
The Raman spectra for the mixed hydrotalcite, HT(CO
anion. The
broadness suggests that the arsenate anions are in slightly different environments,
thus slightly shifting the individual bands to make one very broad band.
32-,AsO4
3-,VO43-)-2h-pH8
(Fig. 3.14), clearly shows the intercalation of all three anions into the hydrotalcite
structure. Bands due to carbonate are observed at 1061 and 1055 cm-1. Bands
associated with vanadate are observed at 942 and 912 cm-1, while bands observed
at 876 and 811 cm-1 are assigned to arsenate vibrational modes. The broad band at
859 cm-1 is attributed to a combination of arsenate and vanadate vibrational
modes. The assignment of the bands is clearly shown in the stacked Raman
spectra of the individual hydrotalcites and the mixed hydrotalcite. At pH 13 two
very broad bands are observed for the mixed hydrotalcite synthesised at pH 13
(Fig. 3.16) due to the large number of As-O- and V-O-
bonds.
The presence of a symmetric stretching CO32- vibrational mode centred at
1060 cm-1 observed for hydrotalcite structures synthesised using decarbonised
water, is almost certainly due to the dissolution of CO2 from the atmosphere.
100
Figure 3.17: XRD patterns of the synthesised hydrotalcites with variable cationic ratios.
101
It is proposed carbonate enters the system via the following reactions:
1: CO2(g) + OH-
(aq) → HCO3
-(aq) + H2O
2: HCO(l)
3-(aq) + OH-
(aq) → CO32-
(aq) + H2O
(l)
At higher pH levels, the concentration of OH- anions is greater, thus CO2
dissolution is more rapid, making CO32-
contamination greater.
There does not appear to be any effect on the Raman spectrum in the carbonate
vibrational region over longer synthesis times. Comparison of hydrotalcites
synthesised for 2 and 48 hours at pH 8, showed no change in the overall band
positions of carbonate, vanadate, or arsenate (Fig. 3.14 and 3.15). The only
noticeable change was in mixed hydrotalcite, HT(CO32-,AsO4
3-,VO43-)-48h-pH8,
and this is due to a decrease in intensity of the As-O- vibrational mode for the
arsenate anion. This decrease in intensity is believed to be due to a reduction in
the number of As-O- symmetric vibrations, possibly due to the arsenate anion
being primarily in the HAsO42- or even H2AsO4
-
form.
4. X-ray diffraction - XRD
The X-ray diffraction patterns of the synthesised hydrotalcites with variable
cationic ratios and the standard reference patterns are shown in Fig. 3.17. Since
the broad peaks of the synthetic samples correspond to those of the reference
hydrotalcite pattern, it can be concluded that the synthesis of the hydrotalcite
structures was successfully achieved. The hydrotalcites synthesised with
carbonate, arsenate, vanadate, or molybdate showed a single poorly crystalline
phase (Fig. 3.17). The d(003) spacing for all the synthesised hydrotalcites ranged
between 7.6 and 8.0 Å, which is commonly observed for hydrotalcite structures.
[13] Changes in the 003 reflection indicate a change in the interlayer distance of
the hydrotalcite layers, where an increase in interlayer space results in a larger
d(003) spacing. For the purpose of this investigation the arsenate, vanadate, and
molybdate hydrotalcites will be compared to the carbonate hydrotalcite with the
same Mg:Al ratio to determine changes in interlayer spacings. An increase in
basal spacing is due to larger anionic species (compared to carbonate) forcing the
102
layers of the hydrotalcite apart. An increase in interlayer distance confirms the
intercalation of anions other than carbonate.
The basal spacing for all 2:1 hydrotalcites remained relatively unchanged, with an
increase of only 0.04 Å observed for the vanadate and molybdate hydrotalcite
structures. This basal spacing increase is minimal, therefore, it is not believed
intercalation of arsenate, vanadate, or molybdate was successful. The sharp
intense peaks in the 2:1 patterns are due to contamination by NaCl. These
hydrotalcite products were washed after synthesis but it appears more washing
was required.
The 3:1 hydrotalcite series showed an increase in the 003 reflection for all anionic
species. This increase in basal spacing indicates that each anion is intercalated
into the interlamellar domain of the corresponding hydrotalcite. Arsenate,
vanadate, and molybdate anions are all sterically larger than carbonate, and
therefore the intercalation of these anionic species forces the hydroxyl layers
apart, (0.14, 0.13, and 0.26 Å, respectively). These results confirm the
intercalation of these anionic species for this particular cationic ratio. The
intercalation of molybdate showed the largest separation of the layers, and this is
due to the molybdate anion having the largest anionic radius of the three anions
investigated.
The 4:1 hydrotalcite structures appear to be more crystalline, with more well
defined and intense peaks. The d(003) value for the 4:1 HT(CO32-) is significantly
larger than the d(003) values obtained for the 2:1 and 3:1 HT(CO32-), which
suggests a greater amount of carbonate is intercalated into the 4:1 hydrotalcite. It
is proposed carbonate preferentially bonds with the Mg-OH lattice, thus a larger
quantity of carbonate is intercalated into the 4:1 hydrotalcite. The d(003) spacings
for the 4:1 hydrotalcites, containing anions other than carbonate, resemble those
values obtained for the 3:1 series. This suggests that the intercalation of arsenate,
vanadate, and molybdate may have occurred, even though there is a reduction in
the interlayer distance compared to the 4:1 carbonate hydrotalcite.
103
5. Controlled rate thermal analysis of carbonate hydrotalcite
The dynamic thermal analysis of the 3:1 HT(CO32-
) is shown in Fig. 3.18.
Table 3.5 summarises the mass loss in mg and the % mass loss over a specific
temperature range.
Table 3.5: Thermal decomposition of carbonate intercalated hydrotalcite
under dynamic conditions.
Temperature range (°C) Mass loss
(%)
28-120 9.30
120-250 10.3
250-330 3.00
330-600 21.5
600-1000 4.60
In the temperature range from ambient to about 300 °C, three overlapping stages
can be observed in the DTG curve (Fig. 3.18). It can be supposed that in this
temperature range the evolution of differently bound water occurs. In the 300 to
400 °C temperature range a sharp decomposition process can be observed due to
dehydroxylation and decarbonation of the mineral. Between 800 and 1000 °C a
slow mass-loss step is observed which is due to the degradation and melting of
residual salt.
In order to better resolve the decomposition processes, controlled rate thermal
analysis (CRTA) experiments were carried out. This analysis technique uses a
preset, constant, slow rate to provide enough time for the slow heat and mass
transfer processes to occur. This ensures that each sample is heated under
identical conditions. With the slow and constant decomposition rate of 0.10
mg/min, the decomposition is carried out under quasi-isothermal and quasi-
equilibrium conditions.
104
Figure 3.18: The dynamic thermogravimetric and differential
thermogravimetric analysis of carbonate intercalated Mg-Al hydrotalcite.
Figure 3.19: The controlled rate thermal analysis
of carbonate intercalated Mg-Al hydrotalcite.
105
The CRTA curves of 155.03 mg of sample are shown in Fig. 3.19. In the ambient
to 236 °C range three different processes can be distinguished, similar to the result
of the dynamic experiment. In this temperature range two isothermal ranges can
be observed at 67 and 192 °C. It means that gas evolution occurred under
equilibrium (isothermal) conditions. Between these two isothermal ranges,
however, a non-isothermal stage can be seen. If a decomposition process of
constant gas evolution rate is non-isothermal, it means that hidden processes
slower than the heat transport have a role. It is believed that after the first
isothermal stage the layers are collapsing, therefore more energy (i.e. higher
temperature) is needed to maintain the preset, constant rate of decomposition.
Thus, it can be concluded that dehydration is accompanied by the partial collapse
(decrease in the 001 spacing) of layers. In the temperature range between 236 and
340 °C two isotherms can be distinguished (at 323 and 336 °C). This separation of
dehydroxylation and decarbonation cannot be observed under dynamic heating
conditions. With the CRTA method a better resolution of the closely overlapping
reactions can be made. The following chemical reaction for the thermal
decomposition is proposed:
3: Mg6Al2(OH)16CO3·xH2O(s) → Mg6Al2(OH)16CO3(s) + xH2O(g)
Table 3.6 reports the decomposition process, the temperature range of this
decomposition, and the mass loss. Dehydration occurs in three stages: a)
isothermal between 29 and 77 °C, b) non-isothermal between 77 and 170 °C, and
c) isothermal between 170 and 235 °C. The calculations for the stoichiometry of
the thermal decomposition shows that the value of x in the dehydration reaction
(Appendix 1) is 6 moles (calculated 5.75). Further, the calculations show that 1
mole of water is lost in step a), 2.6 moles in step b) and 2.3 moles in step c).
It is expected that some collapse of the hydrotalcite structure occurs with this
dehydration process. The model presented suggests: a) loosely structurally bonded
water, this type of water is lost at low temperatures (in this case between 29 and
77 °C), b) water hydrogen bonded to itself in the interlayer space, and c) water
hydrogen bonded to the hydrotalcite hydroxyl surface. Type 2 water is lost
between 77 and 170 °C and Type 3 water between 170 and 235 °C. The
106
temperature required to remove type 2 and 3 water molecules shows how strongly
the water is hydrogen bonded to the hydrotalcite hydroxyl surface.
Dehydroxylation occurs in an isothermal process over the 235 to 330 °C
temperature range. Decarbonation occurs in two steps: a) an isothermal step
between 330 and 371 °C, and b) in a non-isothermal step between 371 and
541 °C.
Table 3.6: Decomposition stages under CRTA conditions.
Decomposition process
Carbonate intercalated hydrotalcite
(sample mass: 155.03 mg)
Temp. range (°C) Mass loss
mg %
Dehydration 1 (isotherm) 29-77 4.6 3.0
Dehydration 2 (non-isotherm) 77-170 11.0 7.1
Dehydration 3 (isotherm) 170-235 9.7 6.3
Dehydroxylation (isotherm) 235-330 20.5 13.2
Decarbonation 1 (isotherm) 330-371 18.4 11.9
Decarbonation 2 (non-isotherm) 371-541 4.2 2.7
6. Thermal analysis and mass spectroscopy - TGA/DTG and MS
6.1. Effect of different oxy-anions on the thermal analysis patterns of 3:1
hydrotalcite
The thermal decomposition of carbonate hydrotalcites consist of two
decomposition steps between 300 and 400 ˚C, attributed to the simultaneous
dehydroxylation and de-carbonation of the hydrotalcite lattice. Dehydroxylation
results in the collapse of the hydrotalcite structure to that of its corresponding
metal oxides, periclase (MgO) and spinel (MgAl2O4) for carbonate hydrotalcites.
107
The intercalation of oxy-anions increases the thermal stability of the hydrotalcite
structure, shown by a delay in dehydroxylation temperatures (Table 3.7). This
increased thermal stability is attributed to a substantial number of hydroxyl groups
involved in a network of hydrogen bonding involving the solvated intercalated
anions. The strength and number of hydrogen bonds associated with the
intercalated anion contributes to the overall thermal stability of the hydrotalcite
structure. Therefore, the stability of the hydrotalcite structure is dependent on the
type of anion present in the interlayer. Carbonate containing hydrotalcites have
been found to be less stable than oxy-anion (arsenate, vanadate and molybdate)
hydrotalcites. Arsenate, vanadate and molybdate anions are more stable and less
reactive than carbonate. This lower reactivity causes a delay in dehydroxylation
temperatures, thus making the hydrotalcite more thermally stable. Therefore,
hydrotalcite thermal stability is anion dependent, and can be controlled by the
incorporation of more stable and less reactive anions.
The antisymmetric shape of the DTG curve of HT(CO32-
) (Fig. 3.20) indicates the
existence of two different environments for interlamellar water: 1) free water
molecules (lower 300 ˚C), and 2) water solvating anionic species (high 300 ˚C).
The decomposition of synthetic hydrotalcites occurs in 3 steps:
1) evaporation of adsorbed water (up to 100 ˚C),
2) elimination of the interlayer structural water (up to 200 ˚C), and
3) dehydroxylation and de-carbonation of the hydrotalcite framework
(up to 400 ˚C).
The ion current curve revealed that the final dehydroxylation of hydrotalcite and
decarbonation occurred simultaneously at around 350 ˚C.
108
Figure 3.20: The thermogravimetric and differential
thermogravimetric analysis of HT(CO32-
).
Figure 3.21: The ion current curves for selected evolved
gases in the thermal decomposition of HT(CO32-).
109
Table 3.7. Summary of the TG analysis spectrum of the synthesised
hydrotalcites.
Hydrotalcite Dehydroxylation
Temperature/s (°C) Peak Shape
HT(CO32- 323 and 347 ) Antisymmetric
HT(CO32-,AsO4
3- 342 ) Symmetric
HT(CO32-,VO4
3- 351 ) Symmetric
HT(CO32-,MoO4
2- 352 and 373 ) Antisymmetric
HT(AsO43- 333 and 366 ) Antisymmetric
HT(VO43- 382 ) Symmetric
HT(MoO42- 349 ) Symmetric
6.1.1. Carbonate hydrotalcite HT(CO32-
)
Carbonate hydrotalcite has an antisymmetric peak with maxima at 323 and
347 °C (Fig. 3.20). The ion current curves for this hydrotalcite are shown in
Fig. 3.21. The mass spectrum shows evolution of OH and H2O vapour at 315 and
316 °C, respectively, thus confirming the partial dehydroxylation of the
hydrotalcite lattice at 323 ˚C. A shoulder is also observed at 350 and 352 °C (H2O
and OH vapour, respectively) attributed to the final dehydroxylation of the
hydrotalcite lattice and loss of water interacting with the carbonate anions. The
loss of OH, H2O, and CO2
from the hydrotalcite lattice at approximately the same
temperature (~350 °C) indicates that bonding between these anions exists.
6.1.2. Carbonate and arsenate hydrotalcite HT(CO32-,AsO4
3-
)
The inclusion of arsenate in the hydrotalcite structure increased the thermal
stability of the hydrotalcite, with a single mass loss being observed at 342 °C in
the dehydroxylation region (Fig. 3.22). The DTG curve of HT(CO32-) is
antisymmetric with an initial mass loss at 323 °C, attributed to the initial
dehydroxylation of the hydrotalcite lattice. The absence of this peak for
HT(CO32-,AsO4
3-
) suggests that there is increased hydrogen bonding between
110
Figure 3.22: The thermogravimetric and differential
thermogravimetric analysis of HT(CO32-,AsO4
3-
).
Figure 3.23: The ion current curves for selected evolved gases in the
thermal decomposition of HT(CO32-,AsO4
3-)
111
arsenate, carbonate, and interlayer water, which render the hydrotalcite more
thermally stable. The extensive network of hydrogen bonding between all the
anions and interlayer water, results in a single mass loss. The ion current curves
confirm the simultaneous dehydroxylation and decarbonation of the structure
(Fig. 3.23). There appears to be a shoulder at around 320 °C for OH and H2
O (ion
current curves), suggesting that a small quantity of interlayer water is not involved
in hydrogen bonding.
6.1.3. Arsenate hydrotalcite HT(AsO43-
)
The antisymmetric nature of the DTG curve between 300 and 400 °C and the
delay in decomposition temperature indicates a considerable amount of arsenate is
intercalated into HT(AsO43-) (Fig. 3.24). The ion current curves for the arsenate
hydrotalcite are given in Fig. 3.25. Comparison of HT(AsO43-) with
HT(CO32-,AsO4
3-) clearly shows that increasing the concentration of arsenate in
the hydrotalcite interlayer causes a delay in the final dehydroxylation of the
structure (366 °C), rendering the hydrotalcite thermally more stable. Even though
the final dehydroxylation step occurs at higher temperatures, the initial
dehydroxylation and the decarbonation process occur at slightly lower
temperatures (333 °C). The ion current curve shows that the peak at 333 °C is
attributed to the evolution of CO2, OH, and H2O. Therefore, the initial
dehydroxylation (weakly bonded H2O) occurs at this lower temperature, along
with the removal of carbonate (introduced through the dissolution of CO2 during
hydrotalcite preparation). There is considerably less carbonate in the structure
compared to HT(CO32-,AsO4
3-
), shown by an increased scaling factor. Due to
dehydroxylation occurring in two steps, it is proposed that the arsenate anions are
not bonded uniformly in the hydrotalcite interlayer and exists in a number of
environments.
112
Figure 3.24: The thermogravimetric and differential
thermogravimetric analysis of HT(AsO43-
).
Figure 3.25: The ion current curves for selected evolved
gases in the thermal decomposition of HT(AsO43-).
113
6.1.4. Carbonate and vanadate hydrotalcite HT(CO3,VO43-
)
The DTG curve for HT(CO3,VO43-) is given in Fig. 3.26. The results indicate that
a small amount of vanadate anions are intercalated into the hydrotalcite interlayer.
This is shown by the absence of an increase in the decomposition temperature
above 351 ºC (indicative of a network of hydrogen bonding involving anions
other than carbonate). However, an increase in the initial dehydroxylation step
indicates intercalation of minor amounts of vanadate occurred. The ion current
curves confirm that the only mass loss associated with OH and water occurs at
352 ºC (Fig. 3.27). This hydrotalcite was synthesised at pH 9, therefore the
vanadate anions exists predominately as the H2VO4- anion. Due to the relatively
large size and smaller charge density of the vanadate anions (compared to
carbonate), carbonate is intercalated preferentially, therefore reducing the amount
of vanadate anions that are intercalated. There is enough vanadate in the structure,
however, to increase the thermal stability of HT(CO3,VO43-) compared to
HT(CO32-
).
6.1.5. Vanadate hydrotalcite HT(VO43-
)
The DTG curve of HT(VO43-) exhibits a symmetric peak in the 300-400 ºC region
(Fig. 3.28). The peak maxima is at 382 ºC, therefore the intercalation of vanadate
has significantly improved the thermal stability of the hydrotalcite. Comparison
of the CO2 scaling factors for HT(CO32-,VO4
3-) and HT(VO43-) shows that the
contamination of carbonate is relatively small (Fig. 3.29). The symmetry of the
peak also suggests that a large quantity of the OH units associated with the
cationic surface are bonded or involved in a network of hydrogen bonding with
intercalated vanadate anions. The ion current curves confirm the simultaneous
dehydroxylation and decarbonation at 380 ºC. The extensive hydrogen bonding
between OH units in the hydrotalcite lattice and solvated vanadate anions are
suspected to cause the high thermal stability.
114
Figure 3.26: The thermogravimetric and differential
thermogravimetric analysis of HT(CO3,VO43-
).
Figure 3.27: The ion current curves for selected evolved
gases in the thermal decomposition of HT(CO3,VO43-).
115
Figure 3.28: The thermogravimetric and differential
thermogravimetric analysis of HT(VO43-
).
Figure 3.29: The ion current curves for selected evolved
gases in the thermal decomposition of HT(VO43-).
116
Figure 3.30: The thermogravimetric and differential
thermogravimetric analysis of HT(CO3,MoO42-
).
Figure 3.31: The ion current curves for selected evolved
gases in the thermal decomposition of HT(CO3,MoO42-).
117
6.1.6. Carbonate and molybdate hydrotalcite HT(CO3,MoO42-
)
The thermal decomposition of HT(CO3,MoO42-) resulted in an antisymmetric
peak in the dehydroxylation/decarbonation region (Fig. 3.30). The first maximum
is at 352 ºC, assigned to the partial dehydroxylation and decarbonation of the
hydrotalcite lattice. It appears that all the OH units of the hydrotalcite lattice are
involved in bonding to carbonate and to a small extent molybdate, indicated by
the increased decomposition temperature. The release of OH, H2O, and CO2 at
corresponding temperatures in the ion current curves (Fig. 3.31) confirms the
dehydroxylation and decarbonation processes. The less intense shoulder at 373 ºC
is assigned to the dehydroxylation and decarbonation of carbonate and OH units
bonded to molybdate anions. The ion current curves confirm the continued
dehydroxylation and decarbonation of the hydrotalcite lattice at this elevated
temperature. The increase in decarbonation temperature is thought to be due to
complex bonding of molybdate with carbonate and water. The thermal analysis
and ion current curves clearly show carbonate existing in two different
environments when other anionic species are present in the structure. Comparison
of the scaling factors of the CO2 ion current curve for HT(CO32-) and
HT(CO32-,MoO4
2-
) shows considerably more carbonate in the structure when
molybdate is present. It is proposed the intercalation of the larger molybdate anion
increases the interlayer distance, which then allows for additional carbonate to
enter the structure.
6.1.7. Molybdate hydrotalcite HT(MoO42-
)
The dehydroxylation and decarbonation peak in the DTG curve is symmetric
(Fig. 3.32). The single mass loss is due primarily to the dehydroxylation of the
hydrotalcite lattice, along with carbonate anions (impurity). The ion current curve
clearly shows the contamination of carbonate (Fig. 3.33), however there is
considerably less carbonate present compared to the carbonate hydrotalcite. The
peak observed at around 370 ºC for HT(CO32-,MoO4
2-) is not seen for
HT(MoO42-) indicating a reduction in the number of bonds involving molybdate,
especially with carbonate. The ion current curve shows the initial dehydroxylation
occurring at 326 ºC, however the major dehydroxylation of the lattice occurred at
118
Figure 3.32: The thermogravimetric and differential
thermogravimetric analysis of HT(MoO42-).
Figure 3.33: The ion current curves for selected evolved
gases in the thermal decomposition of HT(MoO42-).
119
352 ºC. The increased dehydroxylation temperature (by 26 ºC) is due to hydrogen
bonds with intercalated molybdate anions, which causes a stabilising effect. It is
proposed more hydroxide ions are intercalated into the structure, compared to the
other hydrotalcites, to compensate for reduced molybdate anion intercalation, due
to the increased physical size.
7. Mechanism of anion inclusion (intercalation and/or adsorption)
The d(003) is a measure of the hydrotalcite interlayer distance, and as such, any
increase to the physical size of an anion will result in a larger d(003) value.
Fig. 3.17 shows the XRD patterns of the hydrotalcites. No increases in the d(003)
spacing are observed for any of the 2:1 hydrotalcites with oxy-anions (arsenate,
vanadate, or molybdate). Therefore, the intercalation of the oxy-anions is minimal
in 2:1 hydrotalcites, even though Raman spectroscopy showed the presence of
arsenate (850-750 cm-1), vanadate (950-800 cm-1), and molybdate (920-850 cm-1)
(Fig. 3.34). Therefore, the mechanism for oxy-anions inclusion in 2:1 hydrotalcite
appears to be primarily adsorption. The 4:1 hydrotalcites showed a decrease in the
d(003) spacing, which again suggests that the primary mechanism for the inclusion
of oxy-anions for 4:1 hydrotalcites is through adsorption. The only hydrotalcites
to show increased d(003)
values were the 3:1 hydrotalcite series. Therefore, it is
proposed that adsorption is responsible for oxy-anion removal from solution for
hydrotalcites with Mg:Al cationic ratios of 2:1, 3:1, and 4:1. However,
intercalation of arsenate, vanadate, and molybdate is only possible for 3:1
hydrotalcite structures.
7.1. Effect of cationic ratio on the thermal stability of hydrotalcites with
different interlayer anions
7.1.1. Arsenate hydrotalcites
The shape of the DTG curves obtained for the 2:1, 3:1 and 4:1 arsenate
hydrotalcite series vary considerably (Fig. 3.35). The dehydroxylation /
decarbonation band becomes considerably sharper for the 4:1 hydrotalcite,
suggesting the 4:1 hydrotalcite is more crystalline. Comparison of the 2:1 arsenate
120
Figure 3.34: Raman spectra of the synthesised hydrotalcites with variable cationic ratios
121
Figure 3.35: DTG curves of the synthesised hydrotalcites with variable cationic ratio.
122
hydrotalcite and the 2:1 carbonate hydrotalcite (Fig. 3.35) reveals that the DTG
curves are almost identical. Therefore, it appears the intercalation of arsenate into
the 2:1 hydrotalcite structure does not occur under these synthesis conditions. The
Raman spectrum of the 2:1 arsenate hydrotalcite did detect the presence of
arsenate in the structure (Fig. 3.34), but the proposed mechanism for its inclusion
is adsorption. It is also observed that a smaller quantity of arsenate and a much
larger quantity of carbonate is present in the 2:1 structure, determined by the ratio
of the intensities of the bands at approximately 820 (arsenate) and 1060 cm-1
(carbonate) with the band at 555 cm-1
(Al-O-Al linkage in the hydrotalcite
structure). The increase in carbonate concentration indicates that the intercalation
of carbonate anions is more preferable than the intercalation of arsenate anions for
the 2:1 structure.
The appearance of a shoulder in the DTG curve for the 3:1 hydrotalcite at 368 ºC
indicates that arsenate is intercalated. The increase in decomposition temperature
is due to hydrogen bonding between the intercalated arsenate anion and the
hydroxyl layer surface. Intercalation of arsenate therefore increased the thermal
stability of the structure, compared to the 3:1 carbonate hydrotalcite
(decomposition temperature of 347 ºC). Comparison of the ratios of arsenate to
carbonate, detected by Raman spectroscopy is: 0.6 for the 2:1 hydrotalcite, 2.0 for
the 3:1 hydrotalcite, and 1.0 for the 4:1 hydrotalcite. Therefore, the greatest
amount of arsenate is found in 3:1 hydrotalcites. The presence of a shoulder at
decomposition temperatures above 365 ºC is believed to be due to the
dehydroxylation of hydrotalcite layers hydrogen bonded to the arsenate anion. As
the quantity of arsenate increases (indicated by Raman spectroscopy), the
intensity of the band at approximately 370 ºC on the DTG curve increases. The
3:1 hydrotalcite showed the highest quantity of arsenate in the Raman spectrum
and is observed as a relatively intense band in the DTG curve at 368 ºC. The 4:1
hydrotalcite had the second highest concentration of arsenate, and a band at
370 ºC is visible in the DTG curve. The 2:1 hydrotalcite had the lowest
concentration of arsenate and no distinguishable DTG band is observed at this
elevated temperature. Therefore, it is suggested that arsenate anions detected by
Raman spectroscopy for the 2:1 hydrotalcite are predominantly due to adsorbed
arsenate rather than intercalated arsenate anions. However, arsenate anions are
123
both adsorbed on the external surface and intercalated into the interlayer for the
3:1 and 4:1 hydrotalcite structures.
7.1.2. Vanadate hydrotalcites
The Raman bands attributed to the vanadate vibrations are observed as a broad
band between 900 and 800 cm-1 (Fig. 3.34). Comparison of the DTG curves
(Fig. 3.35) shows that the 3:1 hydrotalcite is again the most thermally stable. This
increased thermal stability is not only due to the intercalation of vanadate anions,
but also to the stability of the hydroxyl layer structure. It has been recently
reported by Yang et al., [14] that 3:1 hydrotalcite structures are more stable due to
a decrease in hydrotalcite lamellae energy, compared to 2:1 and 4:1 structures.
The decomposition temperature of the 3:1 vanadate hydrotalcite is 380 ºC, in
comparison to 344 and 329 ºC for the 2:1 and 4:1 hydrotalcites, respectively. The
increased thermal stability of the 3:1 hydrotalcite is a result of the intercalation of
vanadate anions. An increase in the intensity of the V-O symmetric stretching
modes of the vanadate anion, seen in the Raman spectrum at approximately
900 cm-1
, corresponds with increased thermal stability of the 3:1 hydrotalcite. An
increase in the concentration of vanadate anions in the interlayer region increases
the number of hydrogen bonds associated with the intercalation of this species,
and therefore increases the hydrotalcites thermal stability.
7.1.3. Molybdate hydrotalcites
The Raman spectra of the 2:1, 3:1, and 4:1 molybdate hydrotalcite confirms the
presence of molybdate anions by the appearance of bands in the 900-800 cm-1
region (Fig. 3.34). The presence of a sharp intense band at approximately
900 cm-1 is attributed to the molybdate ν1 symmetric stretching modes of the
molybdate anion. Comparison of this band with bands at approximately 545 and
465 cm-1, attributed to the Al-O-Al and Mg-O-Mg linkage in hydrotalcites, shows
the variability of the molybdate concentrations in the hydrotalcite structure.
Comparison of these bands indicates that the concentration of molybdate anions,
either intercalated or adsorbed, varies considerably with variable divalent/trivalent
124
ratio. The ratio of the 900 cm-1 band with the 548 cm-1
band for the 2:1, 3:1, and
4:1 molybdate hydrotalcites are 2.9, 1.6, and 1.5 respectively.
The hydroxyl layer charge of these hydrotalcites is as follows:
2:1 hydrotalcite: [Mg0.66Al0.33(OH2)]
3:1 hydrotalcite: [Mg
0.33
0.75Al0.25(OH2)]
4:1 hydrotalcite: [Mg
0.25
0.80Al0.20(OH2)]
0.20
Therefore, an increase in divalent:trivalent ratio increases the positive layer
charge of the hydroxyl layers of the hydrotalcite structure. The Raman spectra of
the three hydrotalcites showed that the concentration of molybdate decreased as
the divalent/trivalent cationic ratio increased. Due to the size of the molybdate
anion, incorporation of the anion with hydrotalcites is primarily through
adsorption reactions. The DTG curves do not show a considerable increase in
thermal stability, confirming that adsorption is the predominant mechanism for
the inclusion of molybdate onto the hydrotalcite structure.
The slight increase in thermal stability seen for the 3:1 hydrotalcite is due to a
small number of molybdate anions being intercalated. It is proposed that a 2-step
mechanism is involved in the intercalation of the molybdate anions: 1) carbonate
anions are initially intercalated into the structure, which increases the interlayer
distance of the hydroxyl layers, and 2) the increase in interlayer space allows the
larger molybdate anions to partially insert between the layers at the edges of the
hydroxyl sheets. The intercalation of molybdate anions is possible once the
interlayer distance is greater than the anionic diameter of the molybdate anion and
does not involve exchange reactions with carbonate. Only a small percentage of
available molybdate is intercalated due to these size restrictions.
125
8. Chapter summary
Synthetic hydrotalcites have been synthesised and characterised by a number of
analysis techniques to enable a better understanding of the hydrotalcite structure
that forms under seawater neutralisation conditions. The full characterisation of
hydrotalcites synthesised with different cationic ratios and different oxy-anions
will assist in the characterisation of ‘Bayer’ hydrotalcite formed under seawater
neutralisation conditions and Bayer liquor. The seawater neutralisation of Bayer
liquor results in the formation of hydrotalcite-like structures over a wide pH
range. Therefore, the characterisation of hydrotalcites synthesised at different pH
(within the range observed for the neutralisation process) and different cationic
ratios (dependent on pH) will allow for a more accurate identification of the type
of hydrotalcite that forms in the alumina industry. The synthesis of hydrotalcites
under controlled conditions also gave an insight into the effect of pH and time on
the removal (intercalation and/or adsorption) of oxy-anions commonly found in
Bayer liquors.
The ability of hydrotalcites, synthesised at different pH and with different cationic
ratios, to remove arsenate and vanadate was analysed. This investigation has
shown that the synthesis of hydrotalcites in highly alkaline solution reduces the
effectiveness of the structure to remove oxy-anions from solution. It has been
proposed that the reduction is caused by an influx of OH ions competing for the
hydrotalcite interlayer. The Mg:Al ratio had a minimal effect on the overall
removal of vanadate and arsenate from solution, with removal exceeding 90 % in
each case. Hydrotalcites containing vanadate and arsenate are stable for solutions
up to pH 10, however, exposure of these hydrotalcites to highly alkaline solutions
does result in the exchange of a considerable amount of vanadate and arsenate
anions for hydroxyl anions.
X-ray diffraction, infrared and Raman spectroscopy confirmed the formation of
hydrotalcite in regards to the position of bands compared with known values. The
Raman spectra (1200 – 700 cm-1 region) were found to be useful in the
identification of oxy-anion inclusion, and in identifying the extent of carbonate
contamination in oxy-anion only solutions. Combining the results of X-ray
126
diffraction, Raman spectroscopy and thermal analysis enabled the identification of
the mechanism for inclusion of the three oxy-anions investigated. The
predominant mechanism for the removal of these anionic species from solution is
adsorption for 2:1 and 4:1 hydrotalcites. 3:1 hydrotalcites remove oxy-anions by a
combination of adsorption and intercalation processes. The small amount of
intercalated molybdate anions in the 3:1 hydrotalcites is believed to be due to a
2-step mechanism.
Thermal analysis techniques have shown that the decomposition of the
synthesised hydrotalcites occurred in 3 steps: 1) evaporation of adsorbed water
(up to 100 ˚C), 2) elimination of the interlayer structural water (up to 200 ˚C), and
3) dehydroxylation and de-carbonation of the hydrotalcite framework (up to
400 ˚C). Results have shown that hydrotalcites with divalent/trivalent cationic
ratios of 3:1 are thermally more stable than the corresponding 2:1 and 4:1
structures. The antisymmetric shape of the DTG curves in thermal analysis
experiments indicates the existence of two types of interlamellar water molecules,
those that are free and those solvating the anion species. The free water in the
interlayer is removed at considerably lower temperatures than those that are
solvated and involved in a network of hydrogen bonding. An increase in thermal
stability of hydrotalcites with an anionic species other than carbonate is due to an
increase in the number of hydrogen bonds associated with the intercalated
solvated anions and the cationic surface. The intercalation of vanadate anions into
hydrotalcite showed the greatest increase in thermal stability.
The results obtained in this chapter will be used to assist in the characterisation of
Bayer hydrotalcite, synthesised using seawater and Bayer liquors, in the following
chapter. Some of the experimental techniques used for the characterisation of
synthetic hydrotalcites cannot be used for Bayer hydrotalcite (possibility of
organics), for example TG-mass spectroscopy. Therefore, the results found for
synthetic hydrotalcite will be used for comparison.
127
9. References
[1] A. Vaccari, Preparation and catalytic properties of cationic and anionic clays, Catalysis
Today. 41 (1998) 53.
[2] W.T. Reichle, Synthesis of anionic clay minerals (mixed metal hydroxides, hydrotalcite),
Solid State Ionics. 22 (1986) 135-141.
[3] F. Cavani, F. Trifiro, A. Vaccari, Hydrotalcite-type anionic clays: preparation, properties
and applications, Catalysis Today. 11 (1991) 173-301.
[4] F. Trifiro, A. Vaccari, in: J.L. Atwood, J.E.D. Davies, D.D. MacNicol, F. Vogtle, J.M.
Lehn, G. Alberti, T. Bein (Eds), Solid-State Supramolecular Chemistry: Two- and Three-
Dimensional Inorganic Networks., Pergamon, Oxford, 1996, pp. 251-291.
[5] S. Miyata, Physicochemical properties of synthetic hydrotalcites in relation to
composition, Clays and Clay Minerals. 28 (1980) 50-56.
[6] J.M. Fernandez, M.A. Ulibarri, F.M. Labajos, V. Rives, The effect of iron on the
crystalline phases formed upon thermal decomposition of Mg-Al-Fe hydrotalcites,
Journal of Materials Chemistry. 8 (1998) 2507-2514.
[7] S. Miyata, The synthesis of hydrotalcite-type compounds and their structures and
physiochemical properties, Clays Clay Minerals. 23 (1975) 369-375.
[8] W.T. Reichle, Anionic clay materials, ChemTech. 16 (1986) 58-63.
[9] V. Rives, Layered Double Hydroxides: Present and Future, Nova Science, New York,
2001.
[10] V.C. Farmer, Editor, The Infrared Spectra of Minerals, Mineralogical Society London,
UK, 1974.
[11] R.L. Frost, M.L. Weier, J.T. Kloprogge, Raman spectroscopy of some natural
hydrotalcites with sulfate and carbonate in the interlayer, Journal of Raman Spectroscopy.
34 (2003) 760-768.
[12] S.J. Palmer, T. Nguyen, R.L. Frost, Synthesis and Raman spectroscopic characterisation
of hydrotalcite with CO32- and VO3
-
[13] G.W. Brindley, G. Brown, Editors, Mineralogical Society Monograph, No. 5: Crystal
Structures of Clay Minerals and Their X-ray Identification, 1980.
anions in the interlayer, Journal of Raman
Spectroscopy, 38 (2007) 1602-1608.
[14] Z. Yang, H. Zhou, J. Zhang, W. Cao, Relationship between Al/Mg Ratio and the Stability
of Single-layer Hydrotalcite, Acta Physico-Chimica Sinica. 23 (2007) 795-800.
128
CHAPTER 4
Synthesis and characterisation of Bayer
hydrotalcites
129
1. Introduction
The seawater neutralisation of aluminate solution studies performed by Smith et
al., [1, 2], reported that the exact composition of the precipitate was dependent on
the precipitation conditions. The composition of the Bayer hydrotalcite
(hydrotalcite formed from sodium aluminate solutions) is dependent on the pH;
hydrotalcite formed at high pH (pH > 13) has a Mg:Al ratio of 2:1 (Eq. 1), while
those precipitated at pH 8 have a Mg:Al ratio of 4:1 (Eq. 2). At high pH a more
stable microcrystalline carbonate hydrotalcite (Mg4Al2(CO3)(OH)12·xH2O)
forms, due to adsorbed carbon dioxide (CO2) from the atmosphere producing a
saturated carbonate solution. At lower pH (pH < 9.5) a less well defined crystal
structure forms. Due to the decrease of available carbonate in solution, increased
intercalation of other anions into the hydrotalcite structure
(Mg8Al2Cl(CO3)0.5(OH)20·xH2O) is possible. The decrease in available carbonate
is due to lower pH, resulting in a lower adsorption of CO2
and a decrease in
available carbonate anions for intercalation, Eq. 3.
1. 4MgCl2(aq) + 2NaAl(OH)4(aq) + NaOH(aq) + Na2CO3(aq)
→ Mg
4Al2(CO3)(OH)12·xH2O(s) + 8NaCl
(s)
2. 8MgCl2(aq) + 2NaAl(OH)4(aq) + 12NaOH(aq) + ½Na2CO3(aq)
→ Mg
8Al2Cl(CO3)0.5(OH)20·xH2O(s) + 15NaCl
(s)
3. CO2(g) + 2Na2+(aq) + 2OH-
(aq) → 2Na2+(aq) + CO3
2-(aq) + H2O
(l)
From the work by Smith et al., [1], seawater neutralised red mud would consist of
both the 2:1 and 4:1 hydrotalcite. Carbonate is the predominant anion intercalated
into the hydrotalcite interlayer, which hinders the intercalation of other anionic
species. Increase in temperatures showed a slight increase in adsorption
efficiency, [2] attributed to the decrease in carbonate through the conversion of
carbonate to CO2
at higher temperatures.
This chapter details the characterisation of precipitates formed during seawater
neutralisation of bauxite refinery residue liquors, specifically Bayer hydrotalcite.
130
Figure 4.1: Comparison of red mud and seawater neutralised red mud XRD patterns.
131
These characterisations are based on results obtained and reported for synthetic
hydrotalcite synthesised using neutralisation conditions. Bayer precipitates formed
at variable temperatures have also been characterised. The neutralisation of
bauxite refinery residues generally occurs between 50 and 60 °C, however the
formation of hydrotalcite continues as the residue cools. The formation of Bayer
hydrotalcite is one of the mechanisms for the removal of oxy-anions from bauxite
refinery liquors, and therefore has been characterised. Bayer hydrotalcite is
synthesised using seawater, and therefore a high sulfate concentration is present.
The main difference between synthetic and Bayer hydrotalcite is the presence of
intercalated sulfate anions.
Bauxite refinery residues contain a large proportion of unreactive hematite and
silica particles, so hydrotalcite that forms during neutralisation cannot be
separated and therefore can not be analysed. This chapter looks at synthesising
Bayer hydrotalcite from Bayer liquor in the absence of the solid components of
bauxite refinery residues. Bayer hydrotalcite is believed to be similar to
hydrotalcite formed in the neutralisation of bauxite refinery residues.
2. Identification of hydrotalcite formation in seawater neutralised red mud
2.1. X-ray diffraction
Bauxite refinery residue (called “red mud”) is a highly complex residue with
numerous mineralogical phases. [3] A summary of the phases present in the red
mud used in this investigation is given in Table 4.1. Comparison of the XRD
patterns of red mud (RM) and seawater neutralised red mud (SWN-RM)
confirmed the formation of hydrotalcite, shown by a broad band at approximately
12º 2θ (Fig. 4.1). This peak is the characteristic d(003) peak for hydrotalcite
(Mg6Al2(OH)16(CO3)·4H2O). [4] The weak intensity of the hydrotalcite peak is
due to the overshadowing of the sharper and more crystalline mineralogical
phases present. The broadness of the hydrotalcite peak indicates poor crystallinity.
Bayer hydrotalcite prepared in the absence of RM, exhibited the same broadness.
The neutralisation process produces hydrotalcite-like compounds through the
neutralisation of free OH- with Mg, Al, and Ca to form hydroxycarbonates.
132
Figure 4.2: Thermal analysis of an Australian red mud.
Figure 4.3: Thermal analysis of seawater neutralised red mud.
133
Table 4.1: Quantitative XRD analysis of red mud.
Red mud component Formula %
Hematite Fe2O 65.2 3
Sodalite Na8(Al6Si6O24)Cl 6.3 2
Anatase TiO 4.9 2
Boehmite AlO(OH) 3.2
Gibbsite Al(OH) 2.0 3
Calcite Ca(CO) 1.5 3
Quartz SiO 1.1 2
Calcium aluminate hydroxide Ca3Al2(OH) 0.8 12
Rutile TiO 0.8 2
Iron sulfate Fe2(SO4) 0.5 3
Amorphous content n/a 14.3
2.2. Thermal analysis
The comparison of thermal analysis patterns of RM and SWN-RM clearly shows
the formation of hydrotalcite (RM-hydrotalcite) (Fig. 4.2 and 4.3). The primary
mass loss (4.24 %) observed for RM occurred at 219 °C, and is attributed to the
loss of chemically adsorbed water to the aluminium phases found in red mud
(boehmite and gibbsite). [5, 6] Hematite, which makes up to 65 % of the total
composition of red mud, is thermally stable in the heating range used in this
investigation. Therefore, only small mass losses are observed in the DTG curves.
The mass losses in the DTG curve of red mud are attributed to the dehydration of
the aluminium phases (165 – 325 °C) and the decarbonation of calcite (400 –
534 °C). The small shoulder observed at around 280 °C is believed to be due to
the dehydroxylation of calcium aluminate hydrate.
134
Figure 4.4: XRD pattern of precipitate formed during the SWN of Bayer liquor.
135
The thermal analysis of SWN-RM showed two primary mass losses (214 and
289 °C) due to the dehydration of aluminium phases (3.43 %) and the
dehydroxylation and decarbonation of hydrotalcite (3.49 %). The increased mass
loss up to 185 °C is due to the dehydration of hydrotalcite (4):
4. Mg6Al2(OH)12(CO3)·4H2O(s) → Mg6Al2(OH)12(CO3)(s) + 4H2O(g)
An increased mass loss is also observed between 400-550 °C, due to the
additional formation of calcium carbonate species.
3. Bayer hydrotalcites formed during the seawater neutralisation of bauxite
refinery residues
3.1. X-ray diffraction
The XRD pattern of the precipitates formed by the SWN of Bayer liquor and the
corresponding reference patterns are shown in Fig. 4.4. Three mineralogical
phases are detected: 1) hydrotalcite, 2) calcite (CaCO3), and 3) aragonite
(CaCO3). The full width half maximum (FWHM) of the d(003)
hydrotalcite peak
indicates that small crystallites formed. Two phases of calcium carbonate are
formed from calcium cations in seawater and carbonate in Bayer liquor. Sodium
chloride is present due to the evaporation of the residual seawater during mud
drying.
The basal spacing for Bayer hydrotalcite is 7.76 Å, when prepared at room
temperature. Synthetic carbonate hydrotalcites, prepared using SWN conditions,
have basal spacings of around 7.66 Å. The increase in the d(003) spacing obtained
for the Bayer hydrotalcite suggests that anions other than carbonate are
intercalated into the structure. These larger anionic species probably include
sulfate (seawater) and oxy-anions of transition metals, such as arsenate and
vanadate (Bayer liquors).
136
3.2. EDX analysis
The major elements detected using EDX are magnesium, aluminium, sodium,
calcium and chlorine. Deviations in the Mg:Al ratios are expected for different
concentrations of Bayer liquor, since the concentration of aluminium in Bayer
liquors differs widely between refineries. An average Mg:Al ratio of 3.4:1 is
observed for the Bayer liquors used in this investigation (Table 4.2). It is thought
a mixture of different hydrotalcite species form during the SWN process. The
formation of hydrotalcite structures is highly pH dependent, with lower M2+:Mg3+
ratios obtained at higher pH values. The SWN process consists of a large pH
range (pH values starting at 13 and finishing around pH 8.5), which suggests that
a mixture of 3:1 and 4:1 hydrotalcite structures form. The broadness of the 003
reflection in the XRD pattern (Fig. 4.4) suggests that overlapping of similar types
of hydrotalcites is possible.
Table 4.2: EDX analysis of the molar ratio of the three Bayer precipitates.
Bayer HT synthesised at 55 °C 1 2 3 Average Ratio
Mg 10.08 2+ 11.23 11.74 11.02 3.44
Al 2.81 3+ 3.25 3.56 3.21
3.3. ICP-OES analysis
The concentrations of aluminium, arsenate, vanadate, and molybdate were
analysed before and after the SWN process to determine the percentage removal
of each ion. The initial and final concentrations, and percentage removal for each
species, are given in Table 4.3. It is proposed the removal of aluminium cations
from solution are due to the formation of the hydrotalcite hydroxyl layers. The
removal of aluminium from Bayer liquors is essential for the safe disposal and
storage of these refinery residues and it appears that the SWN process is a cheap
and effective way of removing aluminium and oxy-anions from bauxite refinery
residue liquors.
137
The SWN process removes a significant percentage of arsenate from Bayer liquor
(93.34 %). The mechanism for removal is proposed to be the intercalation and/or
the adsorption of the anions into/onto the positive hydrotalcite surface. The large
reduction in concentration is believed to be due to the low initial concentration of
arsenate in the liquor and the relatively high affinity of arsenate for the interlayer
region. The pH of solution during the neutralisation process suggests that
vanadate exists as VO43-, HVO4
2-, and H2VO4-, while arsenate could exist as
AsO43- and HAsO4
2-
. A larger concentration of vanadate species is present in the
liquor, therefore, the percentage removal is not as high as arsenate. However, a
significant amount of vanadate species is still removed (56.8 %). The percentage
removal of molybdate is insignificant, due to the low concentration of molybdate
in the Bayer liquor (less than 2 ppm) and the relatively low affinity for
hydrotalcite intercalation.
Table 4.3: Percentage removal of ions during the SWN of Bayer liquors.
Aluminium Arsenate Vanadate Molybdate
Initial
conc.
(ppm)
1490 ± 5.0 % 6.70 ± 4.8 % 32.5 ± 5.1 % 1.65 ± 4.9 %
Final
conc.
(ppm)
0.295 ± 5.0 % 0.446 ± 4.8 % 14.0 ± 5.1 % 1.16 ± 4.9 %
%
removal 99.9 ± 5.0 % 93.3 ± 4.8 % 56.8 ± 5.1 % 29.5 ± 4.9 %
Removal of these oxy-anions is essential before these refinery residues can be
safely disposed. It must be noted, that the mechanism for removal of these anionic
species may be due to a combination of intercalation and adsorption processes. It
is proposed that smaller anionic species are predominantly intercalated, while the
larger anionic species are adsorbed onto the external surface of the hydrotalcite
structure.
138
Figure 4.5: Infrared and Raman spectra of the
Bayer precipitate in the hydroxyl stretching region.
Figure 4.6: Infrared spectrum of Bayer hydrotalcite
in the carbonate vibrational region.
139
3.4. Raman and infrared spectroscopy
The infrared and Raman spectra of Bayer precipitate observed broad intense
bands centred at around 3400 cm-1
synthetic hydrotalcites. Bayer hydrotalcite did show additional infrared bands in
the lower hydroxyl stretching region at 3047, 2908, and 2752 cm
(Fig. 4.5). The broad bands are attributed to the
stretching modes of hydroxyl groups in the hydroxyl layers and water molecules
associated with the hydrotalcite structure. The positions of the bands are in a
similar region as corresponding bands in the synthetic hydrotalcites (Chapter 3).
There are slight shifts to lower wavenumbers, compared to the synthetic
hydrotalcites, indicating a weakening of the bonds in Bayer hydrotalcite. It is
proposed that the structure of Bayer hydrotalcites is slightly less stable than the
-1. These bands
are attributed to hydrogen bonding between water molecules and interlayer
carbonate and sulfate anions. The other bands are due to OH stretching vibrations
of water coordinated to: 1) other interlayer water, 2) the OH cationic surface, and
3) separate water molecules bound to M3
OH units (where M might be Mg or Al
and any combinational permutation of these metals). This is discussed further in
Chapter 3.
The infrared spectra of the CO32- antisymmetric stretching region (Fig. 4.6) shows
four bands at 1491, 1459, 1403, and 1361 cm-1. Multiple bands indicate the
carbonate is in multiple environments. Results from XRD (Fig. 4.4) revealed that
calcite and aragonite (calcium carbonates) precipitate along with Bayer
hydrotalcite. The formation of these carbonate species is due to the presence of
calcium in seawater. Therefore, the bands are assigned to carbonate in the two
forms of calcium carbonate, and carbonate in the hydrotalcite interlayer. The band
at 1491 cm-1 is attributed to the ν3 mode of aragonite, 1459 cm-1 is assigned to the
ν3 mode of calcite, 1403 cm-1 is assigned to carbonate bonded to water in the
hydrotalcite interlayer, while the band at 1361 cm-1 is assigned to carbonate
bonded to the hydroxyl surface of hydrotalcite. The ν1 mode of aragonite is
observed at 1086 cm-1, and at 1118 cm-1
for calcite.
140
Figure 4.7: Raman spectrum of Bayer precipitate in the 1150 to 950 cm-1
region.
Figure 4.8: Raman spectrum of Bayer precipitate in the 800 to 200 cm-1 region.
141
The position of the bands in the 1600 cm-1 region indicates that a number of
anions are bonded with interlayer waters. The larger bands at 1655 and 1631 cm-1
are probably due to sulfate and carbonate bridging bonds. [7] The other bands are
attributed to carbonate and sulfate bands in different environments.
The lower wavenumber region in the Raman spectrum for the Bayer precipitate is
shown in Fig. 4.7 (1150-950 cm-1) and Fig. 4.8 (800-200 cm-1). The intense peak
at 1085 cm-1 is attributed to carbonate vibrations in both phases of calcium
carbonate, aragonite and calcite. Bands at 280 and 711 cm-1 are assigned to
calcite, while the band at 703 cm-1 is due to aragonite. These bands are attributed
to the ν4 planar bending modes of carbonate. Bands at 1062 and 1075 cm-1 are
assigned to the carbonate ν1 symmetric stretching modes. These carbonate
vibrational bands are assigned to carbonate bound to the hydrotalcite hydroxyl
surface (1075 cm-1) and the band at 1062 cm-1 is assigned to carbonate bonded to
interstitial water (commonly observed in synthetic hydrotalcites). In the lower
wavenumber region, bands at 552 and 465 cm-1
are attributed to the Al-O-Al and
Mg-O-Mg linkage bonds, respectively, commonly observed in hydrotalcites.
The observation of bands at 981 and 992 cm-1 (Fig. 4.7) confirms the presence of
sulfate anions (ν1 symmetric stretch of sulfate). The sulfate bands may have
originated from sulfate anions intercalated and/or adsorbed into/onto the
hydrotalcite structure. The absence of bands in the region 900-800 cm-1
suggests
that the intercalation of arsenate, molybdate, and vanadate is limited or is
overshadowed by the much more intense carbonate and sulfate bands. ICP
analysis showed that only a small concentration of these oxy-anions are present in
Bayer liquor, presumed to be lower than the detection limit of FT-Raman
spectroscopy. The intensity of these bands and the absence of any other sulfate
species in the XRD pattern, indicates that the sulfate detected is associated with
the hydrotalcite structure. The appearance of two sulfate bands, suggests that the
sulfate anions are present in two different environments: 1) sulfate bonded to the
cationic surface of the hydroxyl layer, and 2) sulfate bonded to interlayer water.
142
Figure 4.9: DTG curves of Bayer precipitate,
hydrotalcite, calcium carbonate, and seawater.
Figure 4.10: TG/DTG curve of the Bayer precipitate.
143
3.5. Thermogravimetric Analysis
The Bayer precipitate TG/DTG curves are slightly more complex than the
synthetic hydrotalcite samples. Due to the possibility of organic compounds being
present, a mass spectrometer was not used to identify each component that was
being evolved at the corresponding decomposition temperatures. Therefore, based
on the results determined by XRD, samples of other components of the precipitate
were analysed to identify the decomposition steps (Fig. 4.9). The thermal analysis
patterns of the synthetic hydrotalcites are also used to assist in the identification of
decomposition steps. The thermal analysis of the precipitate (Fig. 4.10) showed
four main decomposition steps. The DTG curves for synthetic hydrotalcite
(Mg6Al2(OH)16CO32-·5H2
O), calcium carbonate, and seawater were band
component fitted to identify the different decomposition steps.
The decomposition steps are:
1) the loss of adsorbed water on the surface of the precipitate (up to 100 °C),
2) dehydroxylation and decarbonation of the hydrotalcite structure (between 200
and 400 °C),
3) the decomposition of calcium carbonate (500-700 °C), and
The largest mass loss, 24.73 %, is assigned to the removal of adsorbed water on
the external surface of the precipitate. The absence of a peak at slightly higher
temperatures, 100-200 °C, suggests that the hydrotalcite does not contain a large
quantity of weakly bonded interlayer water. Two mass loss steps are observed at
313 °C and 360 °C. The broadness of the peak at 313 °C is due to the removal of
interlayer water existing in different environments. Interlayer water experiences
different bonding within the interlayer, thus there are multiple water units with
slightly different bonding strengths. Water which is strongly hydrogen bonded
will be removed at higher temperatures, whilst weaker bonds will be removed at
lower temperatures. Therefore, multiple water molecules with slightly different
bond strengths will result in a broad decomposition band. The sharp intense band
(360 °C) is believed to be due to a large number of water and carbonate units in
similar environments being removed. The decomposition temperatures obtained
are in good agreement with the synthetic hydrotalcite samples (Chapter 3). The
144
presence of these decomposition bands suggests that the bands obtained are
associated with the dehydroxylation and decarbonation of the hydrotalcite
structure. The ion current curve for the synthetic hydrotalcite showed the
evolution of water vapour at 316 °C, confirming the loss of OH units at
313 °C, and the evolution of CO2
at 350 °C, confirming the loss of carbonate
anions in the interlayer. The decomposition steps of Bayer hydrotalcites are:
Mg
3:1 hydrotalcite:
6Al2(OH)16(CO32-,SO4
2-)·xH2O(s) Mg6Al2(OH)16(CO32-,SO4
2-)(s) + xH2O
Mg
(g)
6Al2(OH)16(CO32-,SO4
2-)(s) → MgAl2O4(s) + 5MgO(s) + (CO2,SO2) (g) + 8H2O(g) +
O2(g)
Mg
4:1 hydrotalcite:
8Al2(OH)18(CO32-,SO4
2-)·xH2O(s) Mg8Al2(OH)18(CO32-,SO4
2-)(s) + xH2O
Mg
(g)
8Al2(OH)18(CO32-,SO4
2-)(s) → MgAl2O4(s) + 7MgO(s) + (CO2,SO2) (g) + 9H2O(g) +
2O
2(g)
The delay in the decarbonation temperature for the Bayer precipitate, compared to
synthetic hydrotalcite HT(CO32-), suggests that the Bayer hydrotalcite that forms
is thermally more stable. It is proposed that intercalation of other anions into the
Bayer hydrotalcite, such as sulfate, arsenate and vanadate, increased the
structures’ thermal stability. This is observed for synthetic hydrotalcites
containing oxy-anions and is due to a substantial number of hydroxyl groups
involved in a network of hydrogen bonds involving the negatively charged anions.
As previously mentioned, the increase in the d(003)
spacing compared to the
carbonate hydrotalcite suggests that these larger anionic species are intercalated
into the structure. The ICP results also support this theory, shown by a reduction n
anionic concentration in solution.
The decomposition of calcite and aragonite is believed to occur at temperatures of
around 600 ºC. It is proposed that the mass loss of 3.21 % is due to the evolution
of CO2
from these calcium carbonate species. The mass loss at 517 ºC of 6.21 %
is believed to be due to the evolution of water vapour from calcium hydroxide
species that may also be present in the precipitate.
145
4. The effect of synthesis temperature on the formation of hydrotalcites in
Bayer liquor
4.1. X-Ray Diffraction
The X-ray diffraction patterns of the precipitates and the corresponding
reference patterns are given in Fig. 4.11. Multiple phases are detected and in
different proportions for the four Bayer precipitates that formed. The most
significant phase is hydrotalcite, reference pattern (01-089-0460), identifiable
as the broader peaks in the pattern. The crystallinity of these Bayer
hydrotalcites decreases with increasing temperature, clearly shown by the
broadening and overlapping of 2 peaks at 60° 2θ. The d (003)
spacing of the
synthesised Bayer hydrotalcites (BHT) from 0 °C to 75 °C are 7.71, 7.82, 7.93,
and 7.79 Å, respectively.
The elemental composition (EDX) of the Bayer precipitate suggests the SWN
process produces hydrotalcites with a Mg:Al ratio between 3 and 4, with the
average value of 3.5, independent of temperature up to 55 °C (Table 4.4). The
Mg:Al ratios obtained at 0, 25, and 55 °C are 3.4, 3.8, and 3.4 respectively.
The precipitate formed at 75 °C resulted in a Mg:Al ratio of 6.8. This large
increase in Mg:Al ratio is proposed to be due to the co-precipitation of
hydromagnesite, however, it is thought a Mg:Al ratio of around 3.5 is still
obtained at 75 °C, for the hydrotalcite that formed. This assumption is based
on only minimal changes being observed for other analysis techniques used to
characterise the precipitate. The elements detected using EDX are Mg, Al, Ca,
S, O, C, and Cl. No significant changes are observed in sulfur concentrations,
therefore the synthesis temperature does not appear to have an effect on the
uptake of sulfate anions.
The basal spacing for the Bayer hydrotalcites increased with increasing
temperature up to 55 °C, suggesting that the removal ability increased. The
predominate anions intercalated into the interlayer region are carbonate, sulfate,
and water molecules. Carbonate and sulfate both have very high affinities for the
146
Figure 4.11: XRD patterns of Bayer precipitates synthesised at
different temperatures via the SWN process.
147
Table 4.4: EDX results of the molar ratio of Bayer precipitates synthesised
at 0, 25, 55, and 75 °C.
Bayer HT synthesised at 0 °C 1 2 3 Average Ratio
Mg 10.60 2+ 10.39 11.18 10.72 3.34
Al 3.15 3+ 3.06 3.41 3.21
Bayer HT synthesised at 25 °C 1 2 3 Average Ratio
Mg 12.43 2+ 11.27 12.32 12.01 3.77
Al 3.30 3+ 2.91 3.35 3.19
Bayer HT synthesised at 55 °C 1 2 3 Average Ratio
Mg 10.08 2+ 11.23 11.74 11.02 3.44
Al 2.81 3+ 3.25 3.56 3.21
Bayer HT synthesised at 75 °C 1 2 3 Average Ratio
Mg 13.39 2+ 14.86 14.29 14.18 6.75
Al 1.93 3+ 2.35 2.02 2.10
hydrotalcite interlayer. An increase in temperature is believed to cause slightly
more disordered structures. This disorder causes the hydroxyl layers to be slightly
mis-aligned, which provides greater space between the interlayer and allows
additional anions and water molecules to be intercalated, further increasing the
basal spacing. The inclusion of larger anionic species drives the hydroxyl layers
further apart, thus causing an increase in basal spacing. Crystalline structures are
aligned with finite interlayer distances. However, Bayer hydrotalcites formed at
75 °C, showed a reduction in interlayer distance. This reduction is believed to be
due to the slight dehydration of the interlayer region at these increased
temperatures.
Another predominant phase that formed during the SWN process is aragonite
(CaCO3), which has an orthorhombic crystal system. Only a small quantity of
aragonite formed at 75 °C, however, the formation of other carbonate species
148
Figure 4.12: Raman and infrared spectra of Bayer precipitates
in the hydroxyl stretching region.
149
appear to be favoured, hydromagnesite (Mg5(CO3)4(OH)2·4H2O) and calcium
carbonate hydrate (CCH). Aragonite appears to form predominantly at 25 and
55 °C. The concentration of CCH increased with increasing synthesis
temperature, clearly shown by the sharp peak overlapping the d(006)
peak of
hydrotalcite at approximately 18 ° 2θ.
4.2. Vibrational spectroscopy
4.2.1. Hydroxyl stretching and bending vibrations
The Raman spectra and infrared spectra of the OH stretching region of the
Bayer precipitates are shown in Fig. 4.12. Both Raman and infrared band
profiles in the hydroxyl stretching region are broad, consisting of multiple
overlapping bands. The Raman bands at 3570, 3565, 3574, and 3566 cm-1
, for
the four precipitates in order of increasing temperature, are assigned to the OH
stretching vibrations of –MgOH, while bands at 3445, 3451, 3438, and
3438 cm-1
are assigned to the OH stretching vibrations of –AlOH in
hydrotalcite and to a small extent hydromagnesite for the 75 °C precipitate.
The bands in the infrared spectra are at slightly higher wavenumbers and
include a couple of additional bands. The infrared spectra of BHT @ 25°C and
BHT @ 55°C are quite similar in appearance, whereas, the spectrum for
BHT @ 0°C has one less band, while BHT @ 75°C has two additional sharp
small intensity bands. It is proposed that the absence of the additional band for
BHT @ 0°C is due to the overlapping of the broad intense peaks, while the
additional sharp peaks in BHT @ 75°C is due to two different kinds of OH
groups in hydromagnesite. The bands at 3519 and 3648 cm-1 are assigned to
the two types of OH groups, while 3519 cm-1 is attributed to OH units involved
in hydrogen bonding, while the band at 3648 cm-1 are not. [7] According to the
results obtained by XRD, hydromagnesite only forms at 75 °C. This reinforces
that these peaks are due to hydromagnesite.
150
Figure 4.13: Infrared spectra of Bayer
precipitates in the 1800 - 1200 cm-1 region.
151
The Raman and infrared bands in the region 3400 to 3200 cm-1 are assigned to
the OH stretching vibrations of water coordinated to the cations in the brucite-
like layers. BHT @ 0°C exhibited four bands in this lower wavenumber
region, 3097, 2916, 2769, and 2596 cm-1, while the other Bayer precipitates
exhibited three bands. Infrared bands at 3097 and 2916 cm-1 for BHT @ 0°C,
3034 and 2899 cm-1 for BHT @ 25°C, 3078 and 2932 cm-1 for BHT @ 55°C,
and 3172 and 2950 cm-1 for BHT @ 75°C are believed to be attributed to water
hydrogen bonded to interlayer anions. Infrared bands at around 2700 and
2500 cm-1 are assigned to the calcium carbonate species, aragonite and calcium
carbonate hydrate. The infrared bands at 2769 and 2703 cm-1, BHT @ 0°C and
BHT @ 75°C respectively, are believed to be due to adsorbed water hydrogen
bonded with carbonate associated with hydromagnesite, while bands at around
2550 cm-1
are assigned to water hydrogen bonded to carbonate, associated with
aragonite. These assumptions are base on XRD results, which showed the
presence of aragonite in all samples, while hydromagnesite is only found in
BHT @ 75°C.
The Raman spectrum of BHT @ 55°C appears to be more compact than the
other three samples. It is suggested that a smaller quantity of water is
associated with this sample, in particular interlayer water. The d003 spacing,
found by XRD, showed this Bayer hydrotalcite had the largest interlayer
distance, which would indicate a larger quantity of water and anions in the
interlayer region. However, the absence of the water band at 2950 cm-1
,
suggests that the increase in interlayer distance is an increase of intercalated
anions rather than water.
The water deformation modes are observed in the infrared spectra at around
1650 cm-1 (Fig. 4.13). The Bayer precipitates show water deformations modes
at 1649, 1649, 1651, and 1657 cm-1, with increasing temperature, attributed to
interlayer water hydrogen bonded to interlayer anions. The position of this
band shifts to slightly higher wavenumbers for the precipitates formed at 55
and 75 °C, indicating a weakening of the hydrogen bond. The position of these
bands suggests that interlayer water is hydrogen bonded to carbonate and
sulfate. [8]
152
Figure 4.14: Raman spectra of Bayer
precipitates in the 1200 - 900 cm-1
region.
153
4.2.2. Carbonate vibrational region
The Raman spectra in the 1200 to 900 cm-1 region has multiple bands at
around 1085 cm-1 attributed to the CO32- symmetric stretching vibrations
(Fig. 4.14). The Raman band profiles for the carbonate symmetric stretching
vibrations clearly show the formation of calcite and aragonite at varying
temperatures. The most simplistic profile is observed for BHT @ 25°C, with a
sharp intense band at 1085 cm-1, and a broad shoulder at around 1060 cm-1.
The sharp band at 1085 cm-1, observed in all precipitates, is assigned to the
symmetric stretching mode of carbonate in aragonite. The bands at around
1060 cm-1 are assigned to hydrotalcite and are due to the symmetric stretching
mode of carbonate anions bonded to interlayer water. Raman bands at around
1090 and 1100 cm-1
for the three other precipitates are assigned to calcite and
CCH. These values are in good agreement with literature. [7] The shift towards
higher wavenumbers indicates weaker hydrogen bonding of the carbonate ion
occurs at increased temperatures. This is in harmony with the position of the
water deformation bands.
The overall band profile in the infrared spectra for the carbonate antisymmetric
vibrational region (Fig. 4.13) consists of two or three overlapping bands.
Determination of these bands proved to be more difficult than those in the
Raman spectra, however the following assignments have been made based on
literature and XRD results found in this study. Bands at around 1400 and
1360 cm-1 are assigned to carbonate incorporated into the hydrotalcite
interlayer. BHT @ 75°C has a significantly different shape, due to the sharp
peaks at around 1480 and 1420 cm-1. These sharp bands are assigned to the
carbonate antisymmetric stretching mode of hydromagnesite. Infrared bands
around 1420 cm-1 are assigned to calcite. [9] The broader bands at 1477, 1477,
1486, and 1484 cm-1, with increasing temperature, are assigned to aragonite,
reported in literature to be situated at 1493-70 and 1450-30 cm-1
. [9] The
154
Figure 4.15: Raman spectra of Bayer
precipitates in the 900 - 200 cm-1
region.
155
shoulder at around 1520 cm-1 is also assigned to aragonite and possibly a
minor contribution from carbonate in hydrotalcite. XRD results showed only a
minor quantity of aragonite present in BHT @ 75°C, and the absence of the
peak at 1520 cm-1 for this precipitate, indicates that the peak at 1520 cm-1
is
predominantly due to the antisymmetric stretch of carbonate in aragonite.
There appears to be two peaks in the 1000 to 900 cm-1 region for all four
precipitates, proposed to be intercalated sulfate. A tetrahedral ion, such as
sulfate, has four modes of vibration when it retains its full (Td) symmetry;
these are the symmetric stretching (ν1) modes observed at 983 cm-1, the ν2
bending mode observed at 450 cm-1, the ν3 mode at 1105 cm-1, and the ν4
mode at 611 cm-1. The ν1 and ν2 modes are Raman active only, whereas the ν3
and ν4 modes are both infrared and Raman active. [7] The Raman bands at 990
and 980 cm-1 are assigned to the ν1 S-OH stretch of the sulfate anions in the
hydrotalcite interlayer. The two different peaks suggest that sulfate anions
exist in two different environments within the hydrotalcite interlayer. The band
at 990 cm-1 is proposed to be sulfate anions bonding with the cationic surface
of the brucite-like sheets, while the band at 980 -1
is assigned to sulfate anions
hydrogen bonded to interlayer water. The intercalation of other anionic species
has not been identified using these spectroscopic characterisation techniques.
4.2.3. Cation OH deformation modes
An intense Raman band at 550 cm-1 with a shoulder at 565 cm-1 (Fig. 4.15) is
assigned to the Al(OH)6 unit in hydrotalcite due to the vibration of aluminium-
oxygen bonds. No apparent changes are detectable in these bands, and
therefore, it is proposed that the synthesis temperature has a minimal effect on
the brucite-like sheets of the hydrotalcite structure. Raman bands at around
470 cm-1
are assigned to the Mg-O-Mg linkage bonds in hydrotalcite, and
again no changes in spectra have been observed as a result of change in
temperature.
156
a: Bayer precipitate formed at 0 °C
b: Bayer precipitate formed at 25 °C
c: Bayer precipitate formed at 55 °C
d: Bayer precipitate formed at 75 °C
Figure 4.16: Thermal analysis of Bayer precipitates formed at 0, 25, 55, and 75 °C.
157
4.3. Thermal analysis –TG and DTG
4.3.1. Decomposition between 30 – 230 °C
The four precipitates, in increasing order of synthesis temperature, show a
common broad band stretching from 30 to 175 °C (Fig. 4.16). The broad band
appears to be due to a number of overlapping bands, primarily situated at around
50 and 140 °C. The first band is assigned to the removal of adsorbed water from
the external surfaces of the different precipitates that formed. The second band is
believed to be due to the evolution of water originating from free interlayer water
in Bayer hydrotalcite. The separation of the bands is most clearly seen in
Fig. 4.16a, where a band assigned to adsorbed water is observed at 50 °C with a
mass loss of 3.17 %, while the second band observed at 156 °C, assigned to the
removal of interlayer water, has a mass loss of 9.85 %. The combined mass loss
between 0 and 200 °C decreased with increasing synthesis temperature; 13.02,
11.94, 11.17, and 9.09 %, respectively. It is believed that at increasing
temperature, less water is associated with these structures due to the slight
evaporation / dehydration of the interlayer region. It is observed that the adsorbed
water evolution temperature decreases with increasing synthesis temperatures as
well. This suggests there is less hydrogen bonding (less interlayer water),
involved in Bayer hydrotalcites formed at elevated temperatures, and therefore,
rendering them slightly less stable.
The average molecular formula of Bayer hydrotalcite is
Mg7Al2(OH)18(CO32-,SO4
2-)·xH2
O. The amount of interlayer water associated
with these hydrotalcites is suggested to be between 4 and 6 moles of water
(Chapter 3). The decomposition step for the removal of interlayer water is
proposed to be as follows:
5. Mg7Al2(OH)18(CO32-,SO4
2-)·xH2O(s)
→ Mg
7Al2(OH)18(CO32-,SO4
2-)(s) +
xH2O
(g)
158
Figure 4.17: Stacked DTG curves of the Bayer precipitates in the
dehydroxylation/decarbonation region.
159
The peak at around 204 °C, for the 75 °C precipitate, is believed to be due to the
evolution of water vapour associated with the dehydration of hydromagnesite. The
formula used to represent hydromagnesite is based on the reference formula
identified by XRD. The dehydration of hydromagnesite is as follows:
6: Mg5(CO3)4(OH)2·4H2O(s) → Mg5(CO3)4(OH)2(s) + 4H2O(g)
4.3.2. Decomposition between 250 – 400 °C
Numerous studies on the decomposition of hydrotalcites report the
dehydroxylation of the brucite-like layers and the decarbonation of the interlayer
region occurring at temperatures generally between 300 and 400 °C. [10-14] The
possibility of organics, present in Bayer liquor, in the samples prevented mass
spectroscopy data from being obtained on the evolved gases, therefore,
assignments of the peaks in this chapter will be determined from synthetic
hydrotalcites prepared using SWN conditions (Chapter 3). The DTG curves of the
Bayer precipitates are compared with pure and synthetic compounds, determined
to be present in the precipitate by XRD (Fig. 4.9). This figure will be used in the
analysis of the mass losses observed for Bayer hydrotalcites.
The decomposition temperature, between 250-400 °C (Fig. 4.17), decreases with
increased synthesis temperatures: 380, 381, 376, and 369 °C. This decrease in
decomposition temperature indicates that Bayer hydrotalcites formed at 0 and
25 °C are more stable than those formed at 55 and 75 °C. Synthetic hydrotalcite,
with only carbonate intercalated into the structure has a decomposition
temperature of 350 °C. [15] The Bayer hydrotalcites synthesised in this
investigation all obtained much higher decomposition temperatures. This increase
in stability is believed to be due to the intercalation of sulfate. The intercalation of
sulfate increases the stability of the structure due to an increase in the number of
hydroxyl groups involved in hydrogen bonding between the sulfate anions and the
cations in the brucite-like layers.
160
The decomposition step at around 380 °C for Bayer hydrotalcites synthesised at 0
and 25 °C, and at 369 °C for the Bayer hydrotalcite synthesised at 75 °C, are
assigned to the simultaneous dehydroxylation and decarbonation of the
hydrotalcite structures. The Bayer hydrotalcite synthesised at 55 °C, however, is
antisymmetric in shape. It is believed two types of interlamellar water molecules
are present in this structure: 1) water molecules bonded to the cationic brucite-like
surface (low 300 °C), and 2) water molecules solvated between intercalated
anionic species (high 300 °C). The lower thermal stability of BHT @ 55°C is due
to lower water content in the hydrotalcite interlayer, thus reducing the number of
hydrogen bonds in the structure. The DTG curves of the four hydrotalcites have
been peak fitted and stacked (Fig. 4.17).
Four bands are present for three of the precipitates; BHT @ 0°C exhibited bands
at 311, 339, 366, and 380 °C, BHT @ 25°C exhibited bands at 302, 342, 369, and
382 °C, and BHT @ 55°C exhibited bands at 300, 343, 359, and 377 °C.
BHT @ 75°C exhibited three bands at 314, 353, and 367 °C. The bands at lower
decomposition temperatures, at around 300 °C, are assigned to the removal of
weakly bonded interlayer water. The bands at around 340 °C are assigned to the
initial dehydroxylation of the brucite-like layers of the Bayer hydrotalcite
structures. There appears to be a slight delay in decomposition temperature as the
synthesis temperature increased. This indicates that the hydroxyl layers of Bayer
hydrotalcite become slightly more stable with increased synthesis conditions. The
Bayer precipitate formed at 55 °C showed a very large broad band at 343 °C,
compared to the other precipitates. Raman spectroscopy results indicate a smaller
quantity of interlayer water is present in precipitates formed at 55 °C, shown by
the absence of a band at 3000 cm-1
(Fig. 4.12). A reduction in the amount of
interlayer water is suggested to make the dehydroxylation process easier, due to a
reduction in the number of hydrogen bonds.
The bands situated at around 360 °C are believed to be due to the slight
decarbonation of aragonite. It is thought that a small phase transition occurs,
which results in a small mass loss. Aragonite is assigned to these bands based on:
1) XRD showed an increase in the amount of aragonite in the sample up to 55 °C
and minimal amounts in the 75 °C precipitates, 2) an increase in intensity of the
161
DTG peak at around 360 °C as synthesis temperature increased to 55 °C, and 3)
the absence of a DTG band at 360 °C for the precipitate formed at 75 °C.
The final bands at 380, 382, 377, and 367 °C, are assigned to the simultaneous
dehydroxylation and decarbonation of the Bayer hydrotalcites. The removal of
interlayer sulfate anions also occurs during this decomposition step. The presence
of sulfate anions is believed to have increased the stability of the hydrotalcite
structures, through a highly complex network of strong hydrogen bonds between
the interlayer anions, water, and the cationic surface of the brucite-like layers. The
decrease in thermal stability of the hydrotalcite formed at 75 °C is believed to be
due to the dehydration of the interlayer region during synthesis, which is
supported by the decrease in interlayer distance of this hydrotalcite found by XRD
techniques. A reduction in the number of interlayer anions that form complex
networks of hydrogen bonding renders the structure more thermally unstable. The
full dehydroxylation and decarbonation decomposition steps of Bayer hydrotalcite
are as follows:
7. Mg7Al2(OH)18(CO32-,SO4
2-)(s)
→ MgAl
2O4(s) + 6MgO(s) + (CO2,SO2)(g) + 9H2O(g) + O
2(g)
The additional band observed at 424 °C, in the 75 °C precipitate, is due to the
dehydroxylation of hydromagnesite, shown below. The dehydroxylation of
synthetic hydromagnesite has been reported to occur between 375-450 °C. [16,
17]
8. Mg5(CO3)4(OH)2(s) → 4MgCO3(s) + MgO + H2O(g)
4.3.3. Decomposition between 400 – 650 °C
This temperature region can be separated into two sections; 1) between 400 and
550 °C, and 2) between 550 and 650 °C. The first region is assigned to the
decomposition of MgCO3 (final decomposition step of hydromagnesite):
4MgCO3(s) → 4MgO(s) + 4CO2(g). The peaks are observed at 425 and 492 °C,
162
and are only present in the 75 °C precipitate, confirming the observations obtained
by XRD.
The sharp intense peak at around 600 °C is assigned to the decarbonation of
calcium carbonate species. The mass loss step and peak maxima are as follows:
BHT @ 0°C peak maximum at 631 °C and a mass loss of 8.66 %, BHT @ 25°C
peak maximum at 632 °C and a mass loss of 11.85 %, BHT @ 55°C peak
maximum at 617 °C and a mass loss of 11.99 %, and BHT @ 75°C peak
maximum at 601 °C and a mass loss of 7.67 %. The stability of aragonite appears
to decrease at elevated temperatures. The decarbonation of aragonite is as follows:
9. Ca(CO3)(s) → CaO(s) + CO2(g)
163
5. Chapter summary
The combination of XRD and thermal analysis techniques successfully
identified the formation of hydrotalcite formed from the seawater
neutralisation of red mud. The complexity of red mud residues makes it
difficult to separate the individual components of these residues. Therefore,
hydrotalcite was synthesised and characterised in the absence of red mud.
These hydrotalcites are referred to as ‘Bayer’ hydrotalcite.
EDX found Bayer hydrotalcite prepared in this investigation have an average
Mg:Al ratio of 3.4:1. It is proposed that different hydrotalcites formed (3:1 and
4:1 structures) due to the wide range of pH values that occur during the
neutralisation process. The SWN process removes significant levels of arsenate
and vanadate from solution, determined by ICP-OES, while the removal of
molybdate is insignificant due to the initial low concentration in the liquor. The
removal of these oxy-anions from solution could not be confirmed by Raman
spectroscopy due to detection limits and overshadowing by the large carbonate
and sulfate bands.
XRD showed that the crystallinity of the Bayer hydrotalcite decreased with
increasing temperature. Elevated synthesis temperatures caused the formation of
several other phases to become more favourable, such as hydromagnesite at
75 °C. The formation of hydromagnesite removes magnesium ions from solution,
which is essential for the formation of Bayer hydrotalcite during the SWN
process. The formation of Bayer hydrotalcite is the only know mechanism for the
removal of hydroxide and aluminium ions from Bayer residue, therefore SWN
temperatures need to be kept below 75 °C, to ensure hydrotalcite formation is
maximised. The formation of aragonite is favourable at temperatures between 25
and 55 °C. The interlayer distance of Bayer hydrotalcite increased with
temperature (up to 55 °C), with a maximum d(003)
spacing of 7.93 Å. At 75 °C the
interlayer distance reduced to 7.79 Å believed to be due to the dehydration of the
structure during synthesis.
164
The presence of bands at 3000 cm-1
in the Raman spectra, indicate Bayer
hydrotalcites have large quantities of interlayer water. However, the absence of
this peak in BHT synthesised at 55 °C and a large basal spacing, suggests that
the interlayer region contains a lower percentage of interlayer water and a
higher percentage of interlayer anions. Therefore, the precipitation of Bayer
hydrotalcites at 55 °C appears to be the most effective at removing dissolved
anions from Bayer liquors.
The position of the water deformation modes indicate that interlayer water is
hydrogen bonded to carbonate and sulfate. The intercalation of sulfate anions
is confirmed by the presence of Raman bands at around 990 and 980 cm-1
. The
intercalation of other anionic species is not identified by the techniques used in
this study.
Bayer hydrotalcite showed the same decomposition steps previously observed for
synthetic carbonate hydrotalcite, however, a delay in decomposition is observed.
This increase in thermal stability is due to a substantial network of hydrogen
bonding between the cationic surface of the layers and solvated intercalated
sulfate anions.
The next chapter deals with the phenomenon ‘reversion’, defined as an
increase in pH and aluminium concentration after the neutralisation process.
The cause for reversion has been thought to involve Bayer hydrotalcite,
therefore, a full characterisation has been completed. The following chapter
identifies the primary compounds within red mud that contribute to reversion
and those that don’t cause reversion but change the final pH of the neutralised
bauxite refinery residues.
165
6. References
[1] H.D. Smith, G.M. Parkinson, Seawater Neutralisation: Factors affecting adsorption of
anionic chemical species 7th International Alumina Quality Workshop, Perth, Australia,
2005.
[2] H.D. Smith, G.M. Parkinson, R.D. Hart, In situ absorption of molybdate and vanadate
during precipitation of hydrotalcite from sodium aluminate solutions, Journal of Crystal
Growth. 275 (2005) 1665-1671.
[3] P. Castaldi, M. Silvetti, L. Santona, S. Enzo, P. Melis, XRD, FTIR, and thermal analysis
of bauxite ore-processing waste (red mud) exchanged with heavy metals, Clays and Clay
Minerals. 56 (2008) 461-469.
[4] J.T. Kloprogge, D. Wharton, L. Hickey, R.L. Frost, Infrared and Raman study of
interlayer anions CO32-, NO3
-, SO42- and ClO4
-
[5] V.M. Sglavo, R. Campostrini, S. Maurina, G. Carturan, M. Monagheddu, G. Budroni, G.
Cocco, Bauxite ‘red mud’ in the ceramic industry. Part 1: thermal behaviour, Journal of
the European Ceramic Society. 20 (2000) 235-244.
in Mg/Al-hydrotalcite, American
Mineralogist. 87 (2002) 623-629.
[6] G. Mariotto, E. Cazzanelli, G. Carturan, R. Di Maggio, P. Scardi, Raman and x-ray
diffraction study of boehmite gels and their transformation to α- or β-alumina, Journal of
Solid State Chemistry. 86 (1990) 263-274.
[7] V.C. Farmer, Editor, The Infrared Spectra of Minerals, Mineralogical Society, London,
UK, 1974.
[8] V. Rives, M. Angeles Ulibarri, Layered double hydroxides (LDH) intercalated with metal
coordination compounds and oxometalates, Coordination Chemistry Reviews. 181 (1999)
61-120.
[9] J.A. Gadsden, Infrared Spectra of Minerals and Related Inorganic Compounds,
Butterworth, Sevenoaks, England, 1975.
[10] E. Kanezaki, 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, Materials Research Bulletin. 33
(1998) 773-778.
[11] L. Pesic, S. Salipurovic, V. Markovic, D. Vucelic, W. Kagunya, W. Jones, Thermal
characteristics of a synthetic hydrotalcite-like material, Journal of Materials Chemistry. 2
(1992) 1069-1073.
[12] G.W. Brindley, S. Kikkawa, Thermal behaviour of hydrotalcite and of anion-exchanged
forms of hydrotalcite, Clays and Clay Minerals. 28 (1980) 87-91.
[13] T. Lopez, E. Ramos, P. Bosch, M. Asomoza, R. Gomez, DTA and TGA characterization
of sol-gel hydrotalcites, Materials Letters. 30 (1997) 279-282.
[14] G.W. Brindley, S. Kikkawa, Formation of mixed magnesium and aluminium hydroxides
with interlayer nitrate and carbonate ions, Thermochimica Acta. 27 (1978) 385-386.
166
[15] S.J. Palmer, A. Soisonard, R.L. Frost, Determination of the mechanism(s) for the
inclusion of arsenate, vanadate, or molybdate anions into hydrotalcites with variable
cationic ratio, Journal of Colloid and Interface Science. 329 (2009) 404-409.
[16] Y. Sawada, J. Yamaguchi, O. Sakurai, K. Uematsu, N. Mizutani, M. Kato,
Thermogravimetric study on the decomposition of hydromagnesite 4
MgCO3.Mg(OH)2.4H2
[17] V. Vagvolgyi, R.L. Frost, M. Hales, A. Locke, J. Kristof, E. Horvath, Controlled rate
thermal analysis of hydromagnesite, Journal of Thermal Analysis and Calorimetry. 92
(2008) 893-897.
O, Thermochimica Acta. 33 (1979) 127-140.
167
CHAPTER 5
REVERSION:
Identification and consequence of triggers
Minimisation of reversion
168
1. Introduction
Reversion is a term used to describe the increase in pH and dissolved metal
concentrations in solution after the seawater neutralisation (SWN) process. The
seawater neutralisation process consists of the following components: 1) seawater,
2) supernatant liquor (SNL), 3) RML (Bayer liquor), and 4) red mud slurry
(RMS). The volume of seawater used for the neutralisation of the slurry is
dependent on the concentration of aluminium in Bayer liquor. One of the
objectives of the neutralisation process is to remove aluminium from the liquor
residue, through the formation of Bayer hydrotalcite. The other aim is to add
sufficient seawater to permanently reduce the pH below 8.9. However, an increase
in pH after the neutralisation point (i.e. after all the seawater has been added) can
occur if insufficient seawater is added, and this is referred to as “reversion”. The
calculation of seawater volumes required to neutralise Bayer residue is performed
according to the dissolved levels of caustic and alumina in the liquor. Therefore,
the presence of solid-phase compounds that can increase this neutralisation
requirement will negatively impact the process.
The red mud slurry is prepared by the addition of pre-determined ratios of dry red
mud, SNL, and RML to produce slurry comparable with bauxite residues
produced in the Gove refinery. The neutralisation process involves the
combination of ambient seawater to hot red mud (75 °C), resulting in a reaction
temperature between 50 and 60 °C.
This chapter investigates the components of red mud that are most likely to cause
reversion (increase in pH and aluminium concentration after seawater
neutralisation). This was achieved by measuring the pH and elemental
concentrations of desired species over a period of time. An introduction on the
initial findings of the seawater neutralisation process is presented in the beginning
of this chapter. Synthetic liquors that resemble Bayer liquor are used in this
investigation to reduce the complexity of the red mud system. A number of
components of red mud have been identified as triggers for reversion, whilst
others have been shown to have other implications on the final composition of the
neutralised solution.
169
The minimisation of reversion is essential for the safe disposal and containment of
bauxite refinery residues. The seawater neutralisation process is used to reduce
both the alkalinity and metal concentrations of the residue, however, pH and metal
reversion may increase levels above recommended specifications. Therefore,
methods for minimising reversion have been devised to ensure the safe storage
and disposal of bauxite refinery residues.
1.1. pH reversion
The ‘neutralisation point’ is used to describe the pH value obtained after all the
seawater has been added to the slurry. Note the use of neutralisation in this
research work refers to the neutralisation point and not pH 7. The “end point” is
dependent on the initial causticity of the red mud slurry, and the quantity of
seawater used to neutralise. The initial pH of the slurry is dependent on the red
mud and Bayer liquors used to prepare the slurry. Therefore, slight differences in
the initial pH of the slurry are expected for different experiments.
The shape of the pH versus time graph shows that there are a number of different
reactions occurring during the SWN process (Fig. 5.1). The initial decrease in pH
signifies the formation of magnesium and calcium hydroxides, Ca(OH)2 and
Mg(OH)2
, equations 1 and 2 respectively. The consumption of hydroxide ions in
the formation of these hydroxides causes the pH of solution to decrease. As the
pH decreases, precipitates of hydroxycarbonates of aluminium, calcium, and
magnesium form. Amongst these hydroxycarbonates, hydrotalcite-like
compounds are favoured. [1, 2] The increase in pH after the neutralisation point is
believed to be due to the dissolution of species in red mud once the pH of the
slurry has been lowered.
1. CaCl2(aq) + 2NaOH(aq) Ca(OH)
2. MgCl2(s)
2(aq) + 2NaOH(aq) Mg(OH)
2(s)
After the neutralisation point the pH begins to rise rapidly for 5-7 minutes, before
the rate of reversions slows and plateaus (60 minutes) (Fig. 5.1). Once the pH has
reached plateau, all formation and dissolution reactions are believed to be in
170
Figure 5.1: Seawater neutralisation curve of a red mud slurry
obtained from a Gove refinery in 2008.
Figure 5.2: Effect of the volumetric seawater neutralisation ratio on pH reversion.
171
equilibrium. The average neutralisation point over 5 reactions for SWN-RMS
(Gove 2008) at 55 °C is 9.40 (SW:RMS volumetric ratio of 4.5:1). The pH after
120 minutes increased on average by 1.1 pH units, to a final pH of 10.53 ± 0.025.
This is a pH increase of 11.9 % after the neutralisation point. Examination of the
pH over one week showed no further increases outside of instrumental error
(Fig. 5.1).
1.1.1. Effect of volumetric seawater to RMS ratio
Red mud slurry neutralised with a seawater neutralisation volumetric ratio of 4.5
clearly shows pH reversion (Fig. 5.1). Increasing the volumetric seawater
neutralisation ratio not only reduced the final pH but appears to have eliminated
any signs of reversion (Fig. 5.2). The increased magnesium concentration at ratios
greater than 5 is shown to significantly reduce the extent of reversion, with
volumetric ratios greater than 8 showing no signs of reversion. The mechanism
for the removal of aluminium from solution is due to the formation of additional
Mg,Al hydrotalcites, however the sheer volume of seawater also has a dilution
effect.
1.1.2. Effect of temperature
The neutralisation process was carried out at four different temperatures: 1) 5 °C,
2) room temperature, 3) 55 ºC, and 4) 75 ºC on the same 2008 Gove slurry. Note
the final pHs of these slurries are higher than the 8.5 to 9.5 range specified. This is
because a lower volumetric ratio (seawater:slurry) is used (4.5:1) compared to that
generally used in the refinery (between 8 and 10). A plot of time versus pH
revealed that pH reversion is temperature dependent, where an increase in
temperature increases the rate of reversion (Fig. 5.3). The increased rate of
reversion is in agreement with the Arrhenius equation (Eq. 3), whereby increasing
temperature increases the rate of the reaction.
172
Figure 5.3: Effect of temperature on pH reversion.
173
Table 5.1: Summary of pH during the SWN-RMS at 5, 25, 55, and 75 °C.
5 °C 25 °C 55 °C 75 °C
Neutralisation point 10.43 9.81 9.36 8.54
Plateau reached (minutes) 2.5 mins 16 mins 2.0 mins 1.25 mins
% decrease 24.91 % 25.29 % 25.83 % 29.65 %
5 minutes 10.45 9.88 9.55 8.98
15 minutes 10.50 9.82 10.04 9.38
30 minutes 10.55 9.87 10.30 9.55
60 minutes 10.54 10.18 10.45 9.60
120 minutes 10.52 10.60 10.52 9.59
240 minutes 10.53 10.79 10.51 9.60
% increase 0.95 % 9.08 % 10.94 % 11.04 %
Table 5.2: Percentage increase of aluminium, arsenate, vanadate, and molybdate 60 minutes after neutralisation, determined by ICP-OES.
Aluminium Arsenate Vanadate Molybdate
Neutralisation point (ppm) 877.6 1.050 10.49 1.470
Final pH (ppm) 940.5 2.830 10.76 2.700
% Increase 6.690 62.90 2.395 45.56
174
3. k = Ae
where k is the rate constant, T is temperature (K), Ea is activation energy, A is the
pre-exponential factor, and R is the gas constant.
Ea/RT
A summary of the pH values observed during the four hours of the reaction for
each temperature is given in Table 5.1. Increasing the temperature of the
neutralisation process caused a greater reduction in pH, with 5 °C having a
neutralisation point of 10.43, while at 75 °C a neutralisation point of 8.54 is
obtained. Increasing the neutralisation temperature increases the rate of formation
of hydrotalcite and other compounds responsible for reducing the pH of the slurry.
At lower temperatures (5 °C) the extent of reversion is minimal, with an increase
of less than 0.10 pH units after the neutralisation point. However, the pH of
solution remains highly alkaline over the entire process (pH greater than 10.40).
Increasing the temperature to 25 °C, resulted in a much greater decline in pH
(neutralisation point equal to 9.81), before a slow increase in pH is observed 20
minutes after the neutralisation point. A final pH of 10.79 is obtained four hours
after neutralisation. A total increase of 9.08 % is observed after the neutralisation
point (9.81). Increasing the neutralisation temperature to 55 and 75 °C resulted in
a rapid increase in pH after the neutralisation point (reversion occurred in less
than 2 minutes). The increase in pH is much sharper and there is no lag time
between the neutralisation point and the beginning of reversion. Increasing the
neutralisation temperature causes all formation and dissolution reaction rates to
increase, thus causing the pH to plateau in a shorter amount of time. For both
temperatures the percentage increase is relatively similar.
It appears that a higher portion of pH reducing compounds form at 75 °C, shown
by an increase in the % decrease in pH for 75 °C compared to 55 °C. This
indicates a larger concentration of hydroxyl ions have been removed from
solution. Chapter 4 found that increasing the neutralisation temperature to 75 °C
caused the formation of hydromagnesite, which would contribute to the lowering
of the pH of the slurry. However, the formation of hydromagnesite decreases the
amount of aluminium available for hydrotalcite formation, which ultimately
reduces impurity removal. Both compounds compete for magnesium ions in
175
solution. Therefore, increased seawater volumes are required to ensure both
hydrotalcite and hydromagnesite form at higher temperatures.
1.2. Reversion of dissolved metals
The release of aluminium, arsenate, vanadate, and molybdite via reversion has
been investigated. The concentration of each species before and after the
neutralisation process was analysed using ICP-OES. Batch samples were collected
throughout the SWN process at 15 minute intervals. The percentage increase of
each species after the neutralisation point is provided in Table 5.2.
There is a significant increase in aluminium (7 %) in solution after the SWN
process, with an increase of around 65 ppm. The concentration of re-dissolved
aluminium is due to the dissolution of species found in red mud solids and
possibly the newly formed Bayer hydrotalcite-like structures. A large percentage
increase, relative to the initial concentration after neutralisation, is observed for all
the oxy-anions, with increases of 63, 46, and 24 % for arsenate, molybdate and
vanadate, respectively. However, in regards to the initial concentration of these
oxy-anions in solution, these increases are minimal. It is proposed that the
removal of these anionic species is through adsorption and intercalation reactions
involving the newly formed hydrotalcite structures. It is believed anion exchange
reactions that occur after neutralisation are responsible for the removal of oxy-
anions from the hydrotalcite surface and interlayer. It is also proposed that the
oxy-anions adsorbed on red mud particles are susceptible to exchange reactions
that release the anions back into solution. These exchange reactions are facilitated
by the exchange of the oxy-anion for a more preferable anion, such as carbonate
or sulfate.
1.3. Identification of the source of reversion
Conducting the SWN process in the presence of only one of either SNL, RML, or
RMS, can assist in the identification of the source of reversion (pH). Reversion
does occur in the presence of RMS or RML, but not in the presence of SNL.
However, neutralisation of equal volumes of RML and SNL did not show
176
Figure 5.4: pH plot for the SWN of synthetic SW and SNL.
Figure 5.5: Aluminium concentration after neutralisation,
determined by ICP-OES.
177
reversion. This suggests that a component of RML contributes to reversion, but is
dependent on the concentration of RML in solution. These observations suggest
that reversion is exclusively due to component(s) of RMS and RML, where red
mud is a highly complex mixture containing approximately 12-15 different
mineralogical phases. These observations indicate that reversion is dependent on
the concentration of the compound causing reversion.
1.4 Identification of triggers causing reversion
The term “trigger” is used to describe a compound that causes reversion. The
benefits of reversion being absent in SNL means that different triggers can be
tested to see if they cause reversion in the absence of red mud, reducing the
complexity of the system. The SWN of SNL with the addition of a trigger will
enable each species to be either confirmed or eliminated as a cause of reversion.
2. Triggers causing reversion
Reversion is defined as any increase in pH and aluminium concentration caused
by the presence of a trigger in relation to results obtained for seawater neutralised
synthetic supernatant liquor solutions. The synthetic seawater neutralisation of
SNL will be referred to as the blank sample.
2.1. Seawater neutralised SNL - blank
2.1.1. pH
It can be clearly seen that pH reversion is not present during the neutralisation of
SNL (Fig. 5.4). SNL has an initial pH of around 12, and a pH of 8.35 after
neutralisation. The dramatic decrease in caustic concentration is a result of the
formation of Bayer hydrotalcite (Eq. 4) and hydrocalumite (Ca2Al(OH)6Cl·2H2
O)
(Eq. 5), removing hydroxyl, aluminium and calcium ions from solution.
178
Figure 5.6: Magnesium concentration after neutralisation,
determined by ICP-OES.
Figure 5.7: Calcium concentration after neutralisation,
determined by ICP-OES.
179
4. 6MgCl2(aq) + 2NaAl(OH)4(aq) + 8NaOH(aq) + Na2CO3(aq)
→ Mg
6Al2(OH)16(CO32-)·xH2O(s) + 12NaCl
(s)
5. CaCl2(aq) + NaAl(OH)4(aq) + Ca(OH)2(aq) + 2H2
→ Ca
O
2Al(OH)6Cl·2H2O(s) + NaCl
(s)
2.1.2. ICP-OES
All aluminium is removed from solution during the addition of seawater
(Fig. 5.5), whilst there is still a large concentration of magnesium left in solution
(Fig. 5.6). Therefore, the limiting factor for the formation of hydrotalcite is the
concentration of dissolved aluminium in solution. These experiments were
conducted at 55 °C, and as such hydromagnesite does not form in high quantities.
The absence of any re-dissolved aluminium (after 2 hours) indicates that BHT
hydroxyl layers are stable under these conditions, and therefore does not
contribute to aluminium reversion.
It is suspected that a minor quantity of hydrocalumite forms. However, it is
proposed that hydrocalumite is unstable at low pH values and dissolves back into
solution. The aluminium ions that are released into solution react with excess
magnesium to form the more stable Mg,Al hydrotalcite. The dissolution of
hydrocalumite is proposed to cause the increase in calcium concentration in
solution after 30 minutes (Fig. 5.7). An increase in aluminium is not observed due
to the immediate formation of hydrotalcite.
The remaining magnesium in solution decreased steadily by 25 % over the 2 hour
period (Fig. 5.6). The majority of the remaining magnesium after Bayer
hydrotalcite formation is used in the formation of Mg-calcite. It is also thought a
small amount of hydromagnesite (Mg5(CO3)4(OH)2·4H2
O) forms. The thermal
analysis of BHT (Fig. 5.8) showed a mass loss at 458 (Eq. 6) and 511 °C (Eq. 7),
typical of the thermal decomposition of hydromagnesite. [3, 4]
6. 458 °C: Mg5(CO3)4(OH)2(s) → 4MgCO3(s) + MgO(s) + H2O(g)
7. 511 °C: MgCO
3(s) → MgO(s) + CO2(g)
180
Figure 5.8: TG analysis of Bayer precipitate.
Figure 5.9: Sulfate concentration after neutralisation, determined by ICP-OES.
181
The loss of mass at 511 °C would be expected to be larger, however the mass loss
of the 458 °C peak is around 3 times that of the 511 °C. This is due to the
decomposition of Mg-calcite (Mg0.1Ca0.9CO3
). There is considerably more Mg-
calcite in the precipitate, thus causing a larger mass loss. The decomposition of
Mg-calcite is proposed to be as follows:
8. 458 °C: 10(Mg0.1Ca0.9CO3)(s) → MgO(s) 9CaCO3 + CO
9. 612 °C: CaCO2(g)
3(s) → CaO(s) + CO
2(g)
The formation of hydromagnesite is limited by the concentration of carbonate in
solution. Thus, the formation of hydromagnesite is dependent on CO2 dissolution.
This limitation in hydromagnesite formation can be seen by the slow reduction in
the magnesium concentration over time (Fig. 5.6). The calcium concentration in
solution also steadily declines over time, 11 % of remaining calcium, due to the
formation of CaCO3. The bulk of aragonite is formed in the initial 2 minutes of
the SWN process, however, the continual decrease in calcium ions in solution
suggests calcium carbonate species continue to form. The formation of these
species is limited by the concentration of carbonate in solution, as previously
mentioned for hydromagnesite. Thus, the formation of these CaCO3
species is
dependent on the carbonate concentration and the formation of hydromagnesite.
The synthetic seawater used in this investigation had a sulfate concentration of
900 ppm and a sulfate concentration of 750 ppm in the neutralised solution. ICP
results indicate that around 50 ppm of sulfate is removed from solution within the
first 5 minutes, while a further 150 ppm of sulfate anions are removed over a 2
hour period (Fig. 5.9). The initial removal of sulfate anions is believed to be
through the intercalation of sulfate anions into the hydrotalcite interlayer. Both
carbonate and sulfate have a high affinity for the hydrotalcite interlayer, therefore,
it is though both anionic species are intercalated during the formation of Bayer
hydrotalcite. The further removal of sulfate is proposed to be due to the
precipitation of sodium sulfate (Na2SO4). It is also thought that adsorption and
intercalation of sulfate anions onto/into hydrotalcite continues to occur until the
hydrotalcite layers are perfectly aligned.
182
Figure 5.10: Combined pH plots for the SWN of synthetic
SW and SNL with varying concentrations of Ca(OH)2
.
Figure 5.11: Concentration of magnesium cations in solution for
varying concentrations of Ca(OH)2
in SWN-SNL over 2 hours.
183
2.2. Synthetic SNL with calcium hydroxide (Ca(OH)2
)
2.2.1. pH
The presence of Ca(OH)2 caused the pH of the synthetic SNL to increase after
neutralisation (Fig. 5.10). Therefore, Ca(OH)2 is a contributor to pH reversion. In
the presence of Ca(OH)2, reversion occurs almost instantaneously after the final
volume of seawater is added to SNL. At low concentrations (0.05 and 0.10M) an
increase of less than 0.5 pH units occurs. It is also apparent that an increase in
Ca(OH)2 increases the neutralisation point (minimum pH reached after the
addition of seawater). This increase in neutralisation point is due to the additional
dissociation of Ca(OH)2, which releases a greater amount of OH- ions into
solution. The primary mechanism for the removal of OH-
ions from solution is
through the formation of hydrotalcite-like structures (mixture of 3:1 and 4:1
structures – Chapter 4).
High concentrations of Ca(OH)2 (0.30M and greater) prevented a reduction in pH,
with the final pH remaining at values greater than 11. The neutralisation of high
Ca(OH)2 suspensions still showed a small reduction in pH before an increase
occurs. This reduction is due to the formation of hydrotalcite, hydrocalumite
(Ca2Al(OH)6Cl·2H2O), and brucite (Mg(OH)2). The formation of hydrotalcite
and brucite can be observed by the decrease in Mg2+ concentration (Fig. 5.11).
Hydrotalcite forms over a large pH range (Mg:Al ratio of 2:1 is favoured at high
pH), while hydrocalumite and brucite formation is favoured at high pH. A
significant reduction in pH is not observed due to an influx of OH- ions caused by
the continual dissociation of Ca(OH)2 as OH- ions are used up in the formation of
these three phases. The pH increases once there is a shortage of Mg2+ ions in
solution until a state of equilibrium for Ca(OH)2 is reached (between pH 11 and
11.5). Minimal changes in pH are observed for concentrations greater than 0.40M.
The Mg2+ ion concentration remains constant, therefore increasing the
concentration of Ca(OH)2 has no effect on pH as the dissociation of Ca(OH)2 is
dependent on the concentration of Mg2+ in solution. When all Mg2+ is removed
from solution, the dissociation of Ca(OH)2 ceases once equilibrium pH is reached.
184
The greatest percentage increase is observed for 0.30M Ca(OH)2 (Table 5.3),
which is believed to be the concentration whereby excess Ca(OH)2
in solution
could no longer be neutralised by the addition of this volume of seawater.
The dissociation of Ca(OH)2 releases 2 moles of OH-
ions into solution, which
causes the pH to rise.
10. Ca(OH)2(s) Ca2+(aq) + 2OH-
(aq)
Table 5.3: Initial and final pH of solution and the percentage increased over
an hour period.
Concentration
(Ca(OH)20.05M
) 0.10M 0.30M 0.40M 0.50M 1.00M
Neutralisation
point 8.51 8.66 10.55 11.07 11.12 11.55
2 hours 8.70 8.85 11.44 11.29 11.28 11.25
Difference 0.19 0.19 3.51 0.22 0.16 -
% increase 2.2 % 2.2 % 7.78 % 2.0 % 1.4 % -
There are three shifts in equilibrium observed for the dissociation of Ca(OH)2
during the seawater neutralisation process:
i) At low pH, the equilibrium reaction shifts to the right (dissociation of
Ca(OH)2) due to the use of OH-
ii) The use of Ca
ions in the formation of HT. 2+ in the formation of CaCO3 and CaCl2
iii) At high concentrations of Ca(OH)
, which also shifts the
equilibrium to the right.
2, the release of OH- ions are consumed in
the formation of brucite and hydrotalcite until all Mg2+ ions are removed from
solution. OH-
ions are also consumed in the formation of hydrocalumite.
The solubility of Ca(OH)2 is 1.26g/L at 50 ºC. [5] It should be noted that the
solubility of Ca(OH)2 increases at lower temperatures. Therefore, the presence of
Ca(OH)2 in the residue (disposed of in tailings dam) will show a continual pH
185
increase as the residue cools. The concentration of Ca(OH)2 used for each test and
the theoretical concentration of solid Ca(OH)2 left in solution is provided in
Table 5.4. For 0.05 and 0.10M the complete dissolution of Ca(OH)2
occurs.
At 0.05 and 0.10M, it is proposed that the dissolution of hydrocalumite also
occurs. An increase in Al3+ concentration is also expected to occur, however, due
to excess Mg2+ ions in solution hydrotalcite forms immediately removing Al3+
from solution. Hydrocalumite appears to be stable at high pH. A full discussion on
the effect of hydrocalumite in SNL will be discussed in section 2.3.
It is proposed that along with the dissolution of hydrocalumite and Ca(OH)2
, the
following reactions may also contribute to the rise in pH, albeit at a much slower
rate:
11. CaCO3(s) + 2NaCl(aq) CaCl2(aq) + Na2CO3(aq)
12. Ca(OH)
[5]
2(aq) + Na2CO3(aq) CaCO3(s) + 2NaOH(aq)
An inverse relationship exists between NaCl and Ca(OH)
[6]
2, as the concentration of
Ca(OH)2 increases the amount of solid NaCl in the precipitate decreases. This is
shown in the XRD patterns (Fig. 5.12).
186
Figure 5.12: XRD patterns of calcium aluminate species tested
as triggers and the corresponding reference patterns.
187
Table 5.4: Concentration of Ca(OH)2 in g/L and the concentration of solid
Ca(OH)2
left in SNL and SWN-SNL, if no reactions with the
dissolution products occur.
0.05M 0.10M 0.30M 0.40M 0.50M 1.00M
g/L Ca(OH)
in SNL 2
3.71 g/L 7.41 g/L 22.23 g/L 29.64 g/L 37.05 g/L 74.12 g/L
Soluble
Ca(OH)2 1.26 g/L in
SWN-SNL
1.26 g/L 1.26 g/L 1.26 g/L 1.26 g/L 1.26 g/L
Remaining
Ca(OH)2 2.45 g/L in
SNL
6.15 g/L 20.97 g/L 28.38 g/L 35.79 g/L 72.86 g/L
0.05M 0.10M 0.30M 0.40M 0.50M 1.00M
g/L Ca(OH)
in SWN-
SNL
2
0.67 g/L 1.35 g/L 4.04 g/L 5.39 g/L 6.70 g/L 13.50 g/L
Soluble
Ca(OH)2 1.26 g/L in
SWN-SNL
1.26 g/L 1.26 g/L 1.26 g/L 1.26 g/L 1.26 g/L
Remaining
Ca(OH)2 0 g/L in
SWN-SNL
0.09 g/L 2.78 g/L 4.13 g/L 5.44 g/L 12.24 g/L
188
Figure 5.13: Concentration of calcium cations in solution for varying
concentrations of Ca(OH)2 in SWN-SNL over 2 hours.
189
2.2.2. ICP-OES
The Mg2+ ion concentration decreases significantly with increased Ca(OH)2
concentrations, (Fig. 5.11). The increase in OH- ions in solution, due to a greater
amount of Ca(OH)2 dissociating, results in the additional formation of
hydrotalcite, hydrocalumite, and in particular brucite (Mg(OH)2). The small
amount of hydrocalumite which forms, is stable at high pH (greater than 10.5). [7]
At 0.05 and 0.10M the pH of solution is below pH 9, which is not favourable for
brucite formation, and thus the removal of Mg2+ is dependent on the concentration
of Al3+ ions that remain in solution. At these concentrations all Al3+ ions have
been removed from solution, and therefore the formation of hydrotalcite and
hydrocalumite can not occur. This means there is no mechanism for the removal
of OH- ions from solution, due to the dissociation of Ca(OH)2, and therefore the
solution pH increases. ICP results have shown a significant increase in calcium in
solution as the concentration of Ca(OH)2 increases (Fig. 5.13), but this
relationship is not linear. Therefore, the increase in Ca2+ is believed to be due to
soluble calcium salt CaCl2 (Eq. 11), the dissolution of hydrocalumite (Eq. 5), and
Ca(OH)2
(Eq. 10).
2.2.3. XRD
XRD identified six mineralogical components in the precipitate: hydrotalcite,
calcium hydroxide, hydrocalumite, calcite, aragonite, and sodium chloride
(Fig. 5.12). Ca(OH)2 and hydrocalumite are stable at high pH, whilst hydrotalcite,
calcite, and aragonite are stable in all alkaline solutions. The XRD pattern
highlights the absence of Ca(OH)2 peaks for 0.05 and 0.10M, confirming the
complete dissolution of Ca(OH)2. Very small peaks are observed for 0.50M
indicating that only a small portion of Ca(OH)2 is present in the solid phase,
whilst a large quantity of Ca(OH)2 is still present in the precipitate for 1.00M
Ca(OH)2. It is observed that an inverse relationship exists between the amount of
NaCl and Ca(OH)2 in the precipitate. It is believed equations 11 and 12 describe
this relationship, which contributes to pH reversion. XRD confirmed that
hydrotalcite forms at all concentrations of Ca(OH)2, however, different Mg,Al
ratios are predicted to form due to the range of pH values that these structures
190
Figure 5.14: DTG curves of 1.00M Ca(OH)2 before and after SWN.
191
form in. A high concentration of Ca(OH)2 increases the crystallinity of Bayer
hydrotalcite, due to the increased pH. This is clearly observed through an increase
in the sharpness of the d(003) plane peak. It is also observed that calcite (CaCO3)
predominately forms, with small amounts of aragonite forming at lower pH
(smaller concentrations of Ca(OH)2). The absence of CaCl2 in the precipitate
indicates that the formation of calcium carbonate species is the predominate
mechanism for the removal of Ca2+ ions. CaCO3 is a more stable structure than
CaCl2
, and hence its formation is favoured.
2.2.4. TGA
DTG curves of the solid from SWN of SNL-1.00M Ca(OH)2 confirmed the
formation of hydrocalumite, observed as a shoulder (285 °C) on the Bayer
hydrotalcite peak (300 ºC) (Fig. 5.14). Hydrocalumite is stable at this
concentration of Ca(OH)2 due to high solution pH, however, Bayer hydrotalcite
remains the predominant species. The decomposition of the precipitate, after
seawater neutralisation containing excess Ca(OH)2
, occurs in four steps:
i) the removal of adsorbed water from the external surface of the precipitate
(70 °C),
ii) the dehydroxylation and decarbonation of the brucite-like hydroxyl layers of
Bayer hydrotalcite and hydrocalumite (270-300 °C),
iii) the dehydroxylation of calcium hydroxide (360-380 °C),
iv) decarbonation of calcium carbonate (600-620 °C), and
The interpretation of the DTG curves are based on previous work done on
synthetic and Bayer hydrotalcites (Chapters 3 and 4).
192
Figure 5.15: Combined pH plots for the SWN of synthetic SW and
SNL with varying concentrations of hydrocalumite.
193
2.3. Synthetic SNL with hydrocalumite (Ca2Al(OH)6Cl·2H2
O)
2.3.1. pH
Reversion (pH) is observed when the concentration of hydrocalumite in SNL is
greater than 0.10M (Fig. 5.15). At concentrations below this, the final pH remains
between 8.0 and 8.5. At low concentrations (0.10M and less), the release of OH-
ions from the dissolution of hydrocalumite is used up by the formation of
hydrotalcite, thus preventing the pH from increasing. Reversion only occurs when
the concentration of Mg2+ ions in solution is insignificant, and hydrotalcite is
unable to form. The rate at which reversion occurs is dependent on the
concentration of hydrocalumite in solution (Table 5.5). The increase in pH is due
to the release of OH- ions into solution, caused by the dissolution of
hydrocalumite (Eq. 13). Increasing the concentration of hydrocalumite releases
more OH- ions into solution, therefore causing the pH to rise at a faster rate until a
state of equilibrium is reached. At lower concentrations, the released OH-
ions are
consumed by hydrotalcite formation, so reversion is not observed.
Hydrocalumite (2Ca2Al(OH)6Cl·2H2O) can be re-written in the oxide phases that
it is formed from: 3CaO·Al2O3·CaCl2·10H2O. The synthesis of hydrocalumite in
a sulfate rich environment would have resulted in the formation of ettringite
(3CaO·Al2O3·CaSO4·10H2
O), common in the cement industry. [7]
Hydrocalumite and ettringite are chemically similar, therefore possessing similar
stabilities and reactivity. It has been reported that the stability of ettringite
decreases in solutions with a pH below 10.5. [7] As the pH falls below 10.5, the
dissolution of ettringite occurs. Therefore, the same is proposed to be true for
hydrocalumite. The following equilibrium reaction is proposed:
13. Ca2Al(OH)6Cl·2H2O(s) + NaCl
NaAl(OH)4(aq) + Ca(OH)2(aq) + CaCl2(aq) + 2H2O(l)
In solution NaAl(OH)4 dissociates into Na+ and Al(OH)4- (increase in aluminium
concentration), while Ca(OH)2 dissociates releasing OH-
ions (increase in pH).
194
Figure 5.16: Aluminium concentration in solution after the
SWN of SNL with varying concentrations of hydrocalumite.
Figure 5.17: Magnesium concentration in solution after the
SWN of SNL with varying concentrations of hydrocalumite.
195
As the pH drops below 10.5 during neutralisation, the dissolution of
hydrocalumite becomes favoured, releasing aluminate and hydroxyl ions into
solution (pH and aluminium reversion). The rate of pH increase is determined by
the initial pH and the amount of hydrocalumite in solution. It is observed that the
neutralisation point increases with elevated hydrocalumite concentrations. The
final solution pH, for all concentrations of hydrocalumite that showed pH
reversion, is between pH 10 to 10.5. In this pH range, the remaining
hydrocalumite in solution is in equilibrium, and therefore, the dissolution of
hydrocalumite is no longer favoured.
Table 5.5: Summary of pH results for hydrocalumite concentrations that
showed pH reversion.
Concentration Neutralisation
point
Delay time
(mins) Final pH
% increase
(pH)
0.10M 8.20 N/A 8.04 N/A
0.20M 8.32 32.5 10.23 23.0%
0.30M 8.46 15.0 10.23 20.9%
0.40M 8.53 5.5 10.17 19.2%
0.50M 10.17 2.75 10.38 2.10%
2.3.2. ICP-OES
ICP confirmed that aluminium reversion occurred for hydrocalumite when present
in concentrations above 0.10M in SNL (Fig. 5.16). The reversion of aluminium
appeared to correspond well with pH reversion, therefore, it is believed pH and
aluminium reversion are directly related. This observation reinforces the complete
dissolution of hydrocalumite (Eq. 13). There is an inverse relationship between
the aluminium (Fig. 5.16) and magnesium (Fig. 5.17) concentrations in solution. It
can be clearly seen from these charts that aluminium reversion became prevalent
when magnesium is absent from solution, shown for 0.50M hydrocalumite. The
concentration of magnesium steadily decreases after the SWN process due to the
simultaneous dissolution of hydrocalumite and hydrotalcite formation. The
formation of additional hydrotalcite, stimulated by the release of aluminate ions,
196
Figure 5.18: Calcium concentration in solution after the
SWN of SNL with varying concentrations of hydrocalumite.
Figure 5.19: Combined pH plots for the SWN of synthetic
SW and SNL with varying concentrations of TCA.
197
causes both the magnesium concentration and pH to decrease. The aluminium
concentration in solution does not appear to increase significantly after 30
minutes, because the pH of solution has reached equilibrium within this time
period. Therefore, the dissolution of hydrocalumite no longer occurs.
The calcium concentration varied significantly with the concentration of
hydrocalumite (Fig. 5.18). Increasing the concentration of hydrocalumite in
solution increased the concentration of calcium in solution, as expected. However,
once the concentration of hydrocalumite reached 0.50M, the calcium levels in
solution decreased significantly. The calcium concentration is dependent on the
concentration of hydrocalumite and pH. At lower hydrocalumite concentrations,
0.01M, the amount of calcium in solution from hydrocalumite dissolution is
negligible, since the calcium ions are immediately consumed via calcium
carbonate formation. However, as the concentration of hydrocalumite increased
(0.02, 0.03, and 0.04M), the dissolution of hydrocalumite became more
noticeable. The excess calcium remained dissolved because the carbonate anions
had been depleted. The same trend is observed for hydrocalumite concentrations
of 0.20, 0.30, and 0.40M. However, when the hydrocalumite levels are above
0.5M, the corresponding pH increase appears to cause a secondary increase in
carbonate levels, mainly due to an elevated rate of CO2 absorption (and so CO32-
formation). [8] The pH for 0.50M hydrocalumite remained between pH 10 and
10.5. As a result, more CaCO3 precipitated out of solution, thus decreasing
dissolved Ca2+ levels when normal reaction kinetics suggested it should increase.
It is also proposed a small amount of Ca(OH)2
formed.
2.4. Synthetic SNL with tricalcium aluminate hexahydrate
2.4.1. pH
It is observed that the neutralisation of SNL containing TCA causes an increase in
pH (reversion) after all the seawater has been added (Fig. 5.19). The extent of
reversion is dependent on the concentration of TCA, where reversion occurs for
concentrations above 0.03M. The rate of reversion is also concentration dependent
198
Figure 5.20: Concentration of aluminium in solution
after the SWN process, using ICP-OES.
Figure 5.21: Concentration of magnesium in solution
after the SWN process, using ICP-OES.
199
- the percentage pH increase for 0.05 and 0.10M are 3 and 15 %, respectively,
after 2 hours.
Small quantities of TCA, up to 0.05M, cause the neutralisation point to decrease.
The slight dissolution of TCA increases the concentration of soluble aluminium in
solution, which allows for a larger concentration of Bayer hydrotalcite to form.
The formation of the extra Bayer hydrotalcite removes a larger concentration of
OH-
ions from solution, thus causing the neutralisation point to decrease.
However, the pH increases for higher concentrations of TCA due to the continued
dissolution of TCA, which exceeds the neutralisation capacity of the seawater,
causing the uncontrolled release of hydroxyl ions, calcium and aluminium ions
into solution.
2.4.2. ICP-OES
The increase in aluminium ions, after the seawater neutralisation point, is only
observed at concentrations of TCA greater than 0.05M (9 ppm in the first hour). It
is proposed that a higher concentration of aluminium is released into solution,
however, the continual formation of Bayer hydrotalcite removes this soluble
aluminium from solution. It is not until all magnesium is removed from solution
that the dissolution of TCA is clearly observed (Fig. 5.20). This is seen for 0.10M
TCA, where the concentration of magnesium is below detection limit (Fig. 5.21),
whilst the aluminium concentration has risen to 180 ppm. Therefore, once all
available magnesium is removed from solution (via hydrotalcite formation), there
does not appear to be another mechanism for aluminium removal.
2.4.3. Mechanism for TCA reversion
Based on the works by Whittington et al., [9] Blenkinsop et al., [10] and Alekson,
[11, 12] two predominant reactions are involved in pH and aluminium reversion.
It is thought a combination of the following reactions, involving NaOH and
Na2CO3, participate in forming soluble Al(OH)4- ions from the dissolution of
TCA (C3AH6). It is proposed that the initial dissolution step of TCA involves a
combination of these reactions occurring simultaneously. Eq. 14 is believed to
200
Figure 5.22: Flow chart of the reactions involved in the dissolution of hydrogarnet and TCA.
201
occur initially, which releases NaOH into solution, until the Na2CO3
concentration is depleted. The release of 4 moles of NaOH therefore increases the
pH of solution. However, the consumption of 2 moles of NaOH in Eq. 15 reduces
the rate at which the pH increases.
14. C3AH6(s) + 3Na2CO3(aq)
→ 3CaCO
3(s) + 2Na[Al(OH)4](aq) + 4NaOH
(aq)
15. C3AH6(s) + 2NaOH(aq) 3Ca(OH)2(s) + 2Na[Al(OH)4]
(aq)
The dissolution of Ca(OH)2, formed in Eq. 15, causes an increase in OH- ions.
This increased OH- concentration is then believed to cause the further dissolution
of TCA, removing OH- ions by reforming Ca(OH)2. It is not until the majority of
TCA is removed from solution that the dissociation of Ca(OH)2 becomes apparent
and continues to dissociate until equilibrium is reached.. With no reactions
removing OH- ions from solution, the pH increases significantly. Other reactions
involved occur at a much slower rate, which contribute to reversion either directly
or indirectly. Those that contribute directly are equations 10 and 12, while
equations 12 and 11 fuel other dissolution reactions of TCA. The carbonate
concentration is depleted after the SWN process, therefore equations 14 and 12
are limited by the dissolution of CO2
.
The reactions involved in the dissolution of TCA are summarised in Fig. 5.22.
202
Figure 5.23: Combined pH plots for the SWN of synthetic
SW and SNL with varying concentrations of Bppt.
203
3. Triggers NOT causing reversion
3.1. Synthetic SNL with Bayer precipitate
The term Bayer precipitate refers to the precipitate that forms from the
neutralisation of SNL and synthetic seawater. The precipitate predominately
contains Bayer hydrotalcite, calcium carbonate species, and residual salts (see
Chapter 4). The following results focus on the effects of Bayer hydrotalcite in the
precipitate and how it affects the neutralisation process.
Previous theories on the cause of pH and aluminium reversion have focused on
Bayer hydrotalcite that forms during the SWN process. The dissolution of metals
back into solution is suggested to be caused by either ion exchange reactions,
adsorption/desorption reactions, or a combination of both. However, this is less
predominating than originally believed.
3.1.1. pH
The addition of dried Bayer precipitate did not cause an increase in pH after the
neutralisation point (Fig. 5.23), and therefore is not thought to cause pH reversion.
The final pH of the 0.005M sample (lowest concentration tested) gave a final pH
of 8.10 after 2 hours, in comparison to 8.35 observed for the other samples. This
is a decrease of 0.25 pH units, and is believed due to the adsorption of OH- ions
on the external surfaces of the Bayer precipitate. It is thought that at higher
concentrations, more OH- anions are removed from the hydrotalcite interlayer
(exchanged for carbonate and/or sulfate anions) compared to the concentration of
OH- ions that are adsorbed to the surface of the precipitate. Therefore, more OH-
ions are released into solution causing the pH to increase. The amount of available
carbonate/sulfate for these exchange reactions is assumed to be relatively
constant, therefore it is proposed that the same concentration of OH- ions are
released for all precipitate concentrations above 0.005M. At 0.005M, it is believed
a smaller number of OH- ions in the hydrotalcite interlayer are available for these
exchange reactions. This pH decrease is due to less OH- ions being released into
solution compared with the number of OH- ions adsorbed to the external surfaces.
204
Figure 5.24: Aluminium concentration in solution after the
SWN of SNL with varying concentrations of Bppt.
Figure 5.25: Magnesium concentration in solution after the
SWN of SNL with varying concentrations of Bppt.
205
3.1.2. ICP-OES
ICP results confirmed Bayer hydrotalcite does not contribute to aluminium
reversion, observed in SWN-RM (Fig. 5.24). The magnesium concentration also
remains relatively constant for all concentrations of Bayer precipitate (Fig. 5.25).
Therefore, the amount of precipitate added to solution does not appear to
influence the neutralisation process. The complete removal of aluminium is
observed within the first 2 minutes of the neutralisation process.
3.2. Synthetic SNL with whewellite (CaC2O4·H2
O)
3.2.1. pH
The pH curves for the whewellite samples do not show any increase in pH,
therefore whewellite does not contribute to pH reversion (Fig. 5.26). The
relatively steep increase in pH before the pH decreases is due to the dissolution of
hydrocalumite that forms during the initial neutralisation of SNL. Increasing the
concentration of whewellite in solution causes a reduction in the final pH of the
neutralised solution. As whewellite appears to be stable, increased concentrations
of whewellite in the red mud residue should not affect residue stability.
The variation in pH observed for each of the whewellite concentrations are given
in Table 5.6, and show a linear relationship. Increasing the concentration of
whewellite by a factor of 5 decreased the average pH by 0.12 pH units, and
increasing the whewellite concentration by a factor of 10 reduced the pH by 0.25
units. Therefore, the reduction in pH is proportional to the total surface area of
whewellite in solution. This relationship is proposed to be due to the adsorption of
OH- ions onto the external surface of whewellite.
206
Figure 5.26: Combined pH plots for the SWN of synthetic
SW and SNL with varying concentrations of whewellite.
Table 5.6: Comparison of pH and the concentration of whewellite in SNL.
Concentration of whewellite pH
30 mins 2 hours
0.01 8.52 8.45
0.05 8.42 8.31
0.1 8.32 8.18
0.5 8.13 8.09
Concentrations of whewellite Δ in concentration Δ in pH
0.01-0.10 x10 0.27
0.05-0.50 x10 0.22
Δ in concentration Δ in pH
0.01-0.05 x5 0.14
0.05-0.10 x5 0.13
0.10-0.50 x5 0.09
207
3.2.2. ICP-OES
Whewellite does not cause the reversion of aluminium (Fig. 5.27). Results show
100 % removal. No significant changes are observed for the magnesium and
sulfate concentrations in SWN-SNL.
3.3. Synthetic SNL with sodalite (Na8(AlSiO4)6
Cl)
3.3.1. pH
Sodalite does not cause an increase in pH after the SWN of synthetic SNL, and
therefore is not considered to be a contributor to pH reversion (Fig. 5.28). An
average final pH of 8.35 is obtained for each concentration of sodalite tested,
which strongly resembles the final pH of the blank test. Therefore, sodalite does
not undergo any structural changes or reactions that result in the release or
adsorption of OH-
ions.
3.3.2. ICP-OES
Essentially all aluminium is removed from solution via the SWN process
(Fig. 5.29). Therefore, the sodalite scale remains stable under SWN conditions,
and does not contribute to aluminium reversion. The concentrations of sulfate,
magnesium and calcium all share the same trend as the blank sample. Slight
deviations in the rate of removal are observed, however, these are attributed to
experimental variation.
3.3.3. EDX
The elemental composition of the sodalite scale shows that the elemental ratios of
Na:Al, Na:Si, and Al:Si agree quite well with the theoretical formula
(Na8(AlSiO4)6Cl). The elemental ratios are given Table 5.7.
208
Figure 5.27: Aluminium concentration in solution after the
SWN of SNL with varying concentrations of whewellite.
Figure 5.28: Combined pH plots for the SWN of synthetic
SW and SNL with varying concentrations of sodalite.
209
Figure 5.29: Aluminium concentration in solution after the
SWN of SNL with varying concentrations of sodalite.
Figure 5.30: Combined pH plots for the SWN of synthetic
SW and SNL with varying concentrations of Na2CO3.
210
Table 5.7: Elemental ratio of sodalite scale.
Theoretical Experimental
Na:Al 1.33 1.20
Na:Si 1.33 1.24
Al:Si 1 1.03
3.4. Synthetic SNL with Na2CO
3
3.4.1. pH
It does not appear that Na2CO3
causes the gradual increase in pH that resembles
that of reversion (Fig. 5.30). The SWN-SNL with 0.10 and 0.50M were conducted
over 4 hours and no pH increase was observed. The pH curve resembled that
observed for the blank and whewellite tests. These results indicate that carbonate
does not contribute to pH reversion. The pH for 0.50M solution was monitored
over a 20 hour period, and the pH did not alter above experimental variation.
However, it can be seen that increasing the carbonate concentration delays the
neutralisation-induced reduction in pH by up to 20 minutes. It is proposed that
carbonate acts as a buffering agent before aragonite or magnesian calcite forms.
Therefore, the decrease in pH is dependent on the carbonate concentration in
solution. At concentrations below 0.01M Na2CO3
, the final pH of solution
remains the same as the blank sample. However, at higher concentrations the final
pH of solution increases significantly. This increase in final pH is due to the
buffering effects of the carbonate anions. The primary removal of carbonate from
solution is through the formation of aragonite and Mg-calcite. However, the
concentration of calcium in solution after the SWN process is too low to facilitate
aragonite formation in the presence of this additional carbonate.
211
4. Minimising reversion
The formation of Bayer hydrotalcite has a dual purpose in the neutralisation of
Bayer liquors: 1) reducing OH- ions and thus the pH of solution, and 2) the
removal of aluminium from solution, Eq. 7. The seawater neutralisation of Bayer
liquors results in the formation of hydrotalcite and hydrocalumite, both of which
remove soluble aluminium and OH-
ions from solution. However, hydrotalcite is
stable under seawater neutralisation conditions, whereas, hydrocalumite is
unstable in solution below pH 10.5 and will re-dissolve back into solution. [7]
It has been established that the primary mechanism for aluminium removal from
solution is the formation of Mg,Al hydrotalcite. Therefore, aluminium and OH-
ions that are released back into solution by TCA and hydrocalumite dissolution
can be removed from solution through the formation of hydrotalcite, provided
excess magnesium ions are in solution. The formation of hydrotalcite is only
limited by the concentration of magnesium. Hydrotalcite can still form in low
carbonate concentrations as any anion that meets the intercalation criteria can be
intercalated into the hydrotalcite structure.
4.1. Neutralisation ratio
Increasing the seawater neutralisation ratio has been shown to minimise the extent
of reversion, section 1.1.1. The primary mechanism for aluminium removal from
solution is believed to be the formation of Mg,Al hydrotalcite. This is facilitated
by MgCl2 in seawater, and aluminium and OH- ions from the dissolution of TCA
and hydrocalumite in RMS. The addition of larger volumes of seawater may cause
a dilution effect, where the effective concentration of aluminium in solution
decreased. Therefore, the addition of different concentrations of magnesium
chloride to the same volume of seawater should verify if the formation of Mg,Al
hydrotalcite is the mechanism of aluminium removal and reduction in pH, or if
dilution is the primary cause. These results are discussed in section 4.2.
212
Figure 5.31: pH curves for the addition of MgCl2·6H2
seawater and the SWN of synthetic SNL containing 0.10M TCA.
O to
213
4.2. Addition of MgCl2·6H2
O to synthetic supernatant liquor
Red mud slurry contains a considerable amount of TCA with respect to the
concentration of TCA required to cause an increase in pH after neutralisation
(8 g/L in Bayer liquor causes pH and aluminium reversion). As TCA is a known
trigger for reversion, synthetic liquors containing TCA will be tested to determine
if increasing the magnesium content in seawater will reduce reversion.
Magnesium chloride (MgCl2·6H2O or MgCl2
for short) will be used as a source
of excess magnesium ions.
The addition of increasing levels of magnesium chloride to seawater improved the
reduction in pH and aluminium after neutralisation of synthetic SNL solutions
containing TCA (Fig. 5.31). Using normal seawater, the pH increased by 1 pH
unit, while the aluminium concentration increased by 25 ppm after neutralisation.
The addition of 100 ppm MgCl2 had a minimal effect on reducing the pH and
aluminium concentration of the neutralised liquor (Fig. 5.31 and 5.32). However,
when the concentration of MgCl2 is 200 ppm a slight reduction in pH and
aluminium concentration are observed. The continued increase in MgCl2 in
seawater, up to 500 ppm, showed a continual reduction in both pH and aluminium
reversion. Using 500 ppm MgCl2
in seawater resulted in a pH rise of less than 0.1
units, while the aluminium concentration remained below 3 ppm.
The magnesium concentrations used for these investigations are given in
Fig. 5.33. Increasing the concentration of MgCl2 in seawater causes the initial
magnesium concentration in the system to be higher. At low concentrations
(0-200 ppm), the magnesium concentration in the neutralised solution is very low
(0-5 ppm left in solution after 1 hour). This large reduction in magnesium
concentration is due to the formation of hydrotalcite. It is not until the aluminium
concentration in solution is depleted (no more dissolution reactions involving
TCA and hydrocalumite) that an increase in magnesium concentrations can be
detected (no more hydrotalcite formation). This is clearly observed for 500 ppm
MgCl2, where the magnesium concentration after 75 mins is approximately
200 ppm, whilst the aluminium concentration is approximately 3 ppm.
214
Figure 5.32: Aluminium concentration for SWN synthetic SNL containing
0.10M TCA with additional MgCl2·6H2
O added to seawater.
Figure 5.33: Magnesium concentration for SWN synthetic SNL containing
0.10M TCA with additional MgCl2·6H2
O added to seawater.
215
Increasing the magnesium concentration in SWN solutions removes OH- and
aluminium ions released into solution by the dissolution of TCA and
hydrocalumite. Bayer hydrotalcite is stable in alkaline solutions, and therefore the
OH- and aluminium ions are permanently removed by this process. [13]
Therefore, these results indicate that an increase in seawater volume would also
provide additional magnesium to the system, and so facilitate formation of
additional hydrotalcite. Importantly, this represents a chemical decrease in
aluminium and OH-
concentration, and not simply a dilution effect.
Comparison of figures 5.32 and 5.33 clearly showed an inverse relationship
between magnesium and aluminium. As the concentration of magnesium
increases in solution, the concentration of aluminium decreased. The increase in
aluminium at low concentrations of MgCl2·6H2
O is due to the continual
dissolution of TCA and hydrocalumite. It is not until magnesium is in excess that
the concentration of aluminium in solution is essentially 0 ppm, and therefore
reversion is prevented.
4.3. Addition of MgCl2·6H2
O to red mud slurry
To ensure red mud components do not interfere with the addition of MgCl2 in the
reduction of reversion, 1000 ppm MgCl2 has been added to SWN-RMS
(Fig. 5.34). SWN-RMS and SWN-RMS (1000 ppm MgCl2) both use the same
volume of seawater for the neutralisation process, the only difference is that
SWN-RMS (1000 ppm MgCl2) had 1000 ppm of MgCl2·6H2O added to seawater
before neutralisation. Comparison of the pH curves of neutralised slurry with and
without this additional MgCl2 clearly shows the minimisation of reversion using
excess magnesium (Fig. 5.34). There is a reduction in final pH of around 1.5 pH
units when 1000 ppm of MgCl2 is added to seawater. The pH of the SWN-RMS
with additional MgCl2·6H2O was monitored over a 24 hour period, and no
increase (above pH drifting errors) was observed. Therefore, the formation of
additional Bayer hydrotalcite in the slurry permanently removed both OH and
aluminium ions from solution. Both the synthetic SNL and RMS trials prove that
reversion can be essentially eliminated with the addition of MgCl2·6H2O.
216
Figure 5.34: pH curves for the addition of MgCl2·6H2
seawater and the SWN of red mud slurry.
O to
Figure 5.35: Thermal analysis of seawater neutralised red mud
slurry with an additional 1000 ppm of magnesium chloride.
217
4.4. Confirmation of hydrotalcite formation
Comparison of figures 5.34 and 5.35 clearly shows that the addition of 1000 ppm
of MgCl2·6H2O results in the formation of additional hydrotalcite. This is
observed as an increase in the mass loss percentage, in the region 230 to 400 °C,
of 1.02 %. The mass loss in the absence of 1000 ppm MgCl2 is 3.49 %, compared
to 4.51 % in the presence of 1000 ppm MgCl2. Comparison of the two DTG
curves also shows that the addition of magnesium chloride reduces the formation
of CaCO3 (absence of a peak at 500 °C). The reduction in CaCO3 formation is
due to a lack of carbonate anions. Carbonate anions have a high affinity for the
hydrotalcite interlayer and are thus consumed in the formation of the additional
hydrotalcite.
218
5. Chapter summary
The pH plot of seawater neutralised red mud slurries show a common trend: 1)
rapid decrease in pH, 2) similar neutralisation points, 3) a slow increase in pH
after neutralisation, and 4) plateau of the final pH. The initial pH decrease
represents the rapid formation of Mg, Ca, and Al hydroxycarbonates, in particular
hydrotalcite and hydrocalumite. The slower increase in pH subsequent to
neutralisation is due to the dissolution of Ca(OH)2
, hydrocalumite (formed during
the initial decrease in pH), and TCA in the red mud. The average neutralisation
point for SWN-RMS at 55 °C is pH 9.40, with an average increase of 1.1 pH units
giving a final pH of 10.53 ± 0.03. The pH of the neutralised slurry does not
continue to increase if a pH > 10.5 is achieved, because reactions involving the
three trigger compounds reach equilibrium in this pH range.
It has been shown that reversion is prevalent for volumetric neutralisation ratios
less than 5, and is absent for ratios greater than 8. Increasing the neutralisation
temperature increases the rate of reversion, which agrees with the Arrhenius
equation. Increasing the neutralisation temperature causes all reaction rates to
increase, thus causing the pH to plateau in a shorter amount of time. At increased
neutralisation temperatures, the neutralisation point decreases due to the increased
formation and dissolution rates, and the formation of hydromagnesite at 75 °C.
However, the formation of hydromagnesite removes magnesium ions from
solution, thus reducing neutralisation efficiency and causing higher aluminium
concentrations to remain in solution. Combining high neutralisation temperatures
with additional MgCl2·6H2
O should prevent reversion, and also reduce the
neutralisation point.
This investigation has shown that the presence of calcium hydroxide in
supernatant liquor resulted in a pH rise after neutralisation. It was believed this pH
increase was due to:
1) the dissolution of Ca(OH)2
2) the dissolution of hydrocalumite.
and,
219
A high concentration of Ca(OH)2 facilitates the formation of hydrocalumite
(increase in calcium ions). Therefore, aluminium and pH reversion increases if the
solution remains below 10.5. However, at very high concentrations of Ca(OH)2
the dissolution of Ca(OH)2 results in pH values greater than 10.5, and thus
hydrocalumite is stable. The solubility of Ca(OH)2
is sufficient that the pH is
above 11 at equilibrium. The presence of carbonate also promoted calcium
hydroxide dissolution, through the precipitation of calcite, resulting in additional
release of hydroxide ions into solution. The dissolution of hydrocalumite not only
caused pH reversion, but aluminium reversion as well. The concentration of
calcium hydroxide needs to be minimised in bauxite refinery residues to ensure
the neutralisation process is efficient.
Tricalcium aluminate hexahydrate (TCA) has also been identified as a trigger
responsible for pH and aluminium reversion. A number of reactions are proposed
to be involved in the dissolution of TCA, including: 1) reaction with sodium
carbonate to form sodium hydroxide, sodium aluminate ions, and calcium
carbonate, and 2) with sodium hydroxide to form sodium aluminate ions and
calcium hydroxide.
It has been proven that Bayer hydrotalcite does not cause reversion, as previously
speculated. An increased amount of whewellite in solution has been found to
reduce the pH of solution. The adsorption of OH-
ions on the surface of
whewellite causes a reduction in pH. Therefore, the reduction in pH is dependent
on the surface area of whewellite. Increased concentrations of sodalite and
gibbsite had no effect on pH after the neutralisation of synthetic liquors. The only
component that had an adverse effect on the pH, but did not contribute to
reversion, was sodium carbonate. The dissolution of sodium carbonate caused an
overall increase in pH due to a buffering effect caused by the high carbonate
concentration in solution.
This investigation has identified two methods for the minimisation of pH and
aluminium reversion: 1) increasing the seawater neutralisation volumetric ratio,
and 2) addition of magnesium chloride to seawater. Increasing the volume of
seawater showed a significant reduction in reversion. The addition of
220
MgCl2·6H2O to seawater confirmed that the decrease in aluminium and hydroxyl
ions can be attributed to the formation of hydrotalcite and not simply a dilution
effect (previously thought to occur for increase seawater volumes). Therefore,
increasing the magnesium concentration, either by additional seawater or
MgCl2·6H2
O, can be used to minimise reversion and reduce the final pH of the
slurry for safe disposal. Both methods utilise the formation of hydrotalcite to
remove both hydroxide and aluminium ions from solution permanently. By
ensuring that there is an excess of magnesium ions in solution, reversion can be
prevented.
The final chapter looks at the use of seawater neutralised bauxite refinery residues
and synthetic hydrotalcites as adsorbents for the removal of oxy-anions in
solutions. There are a variety of oxy-anion species within the refinery residues,
and therefore the use of materials already on site would prove beneficial to the
industry.
221
6. References
[1] D.J. Glenister, M.R. Thornberg, Alkalinity of red mud and its application for the
management of acid wastes, Chemica. 85 (1985) 100-113.
[2] C. Hanahan, D. McConchie, J. Pohl, R. Creelman, M. Clark, C. Stocksiek, Chemistry of
Seawater Neutralization of Bauxite Refinery Residues (Red Mud), Environmental
Engineering Science. 21 (2004) 125-138.
[3] Y. Sawada, J. Yamaguchi, O. Sakurai, K. Uematsu, N. Mizutani, M. Kato,
Thermogravimetric study on the decomposition of hydromagnesite 4
MgCO3.Mg(OH)2.4H2
[4] V. Vagvolgyi, R.L. Frost, M. Hales, A. Locke, J. Kristof, E. Horvath, Controlled rate
thermal analysis of hydromagnesite, Journal of Thermal Analysis and Calorimetry. 92
(2008) 893-897.
O, Thermochimica Acta. 33 (1979) 127-140.
[5] P.J. Durrant, General and Inorganic Chemistry, Longmans, London, UK, 1960.
[6] R.B. Heslop, P.L. Robinson, Inorganic Chemistry, Elsevier, London, UK, 1961.
[7] M. Chrysochoou, D. Dermatas, Evaluation of ettringite and hydrocalumite formation for
heavy metal immobilization: Literature review and experimental study, Journal of
Hazardous Materials. 136 (2006) 20-33.
[8] H.D. Smith, G.M. Parkinson, Seawater Neutralisation: Factors affecting adsorption of
anionic chemical species, 7th International Alumina Quality Workshop, Perth, Australia,
2005.
[9] B.I. Whittington, The chemistry of CaO and Ca(OH)2
[10] R.D. Blenkinsop, B.R. Currell, H.G. Midgley, J.R. Parsonage, The carbonation of high
alumina cement, Part 1, Cement and Concrete Research. 15 (1985) 276-284.
relating to the Bayer process,
Hydrometallurgy. 43 (1996) 13-35.
[11] A.I. Alekseev, Calcium Hydroaluminates and Hydrogarnets: Synthesis, Properties, and
Application, LGU, Leningrad, USSR, 1985.
[12] A.I. Alekseev, L.D. Barinova, N.P. Rogacheva, O.V. Kulinich, Thermodynamic and
experimental analysis of equilibriums in the sodium oxide-calcium oxide-carbon dioxide-
water system, Zhurnal Prikladnoi Khimii, 57 (1984) 1256-1261.
[13] V. Rives, Layered Double Hydroxides: Present and Future, Nova Science, New York,
2001.
222
CHAPTER 6
Thermally activated seawater neutralised red
mud used for the removal of arsenate,
vanadate and molybdate from aqueous
solutions.
223
1. Introduction
Due to the vast quantity of bauxite refinery residues that are produced each year,
there is keen interest in developing alternate uses for this material or its
derivatives. This study looks at using thermally activated Seawater Neutralised
Red Mud Slurry (SWN-RMS) for water purification. The neutralisation of bauxite
residues results in the formation of hydrotalcite. Hydrotalcite has been shown to
be a useful adsorbent material after thermal activation, [1-4] while a number of
studies have also involved the use of treated red mud for the removal of heavy
metals and arsenic from solutions. [5-10]
Bayer liquor contains high levels of hydroxide, carbonate, aluminate, chloride,
oxy-anions of transition metals (such as arsenate and vanadate) and sulfate (as
sodium co-anions), in addition to sodium oxalate and organic acid anions. [11]
Therefore, these liquors need to be treated before they can be safely disposed and
stored. The formation of hydrotalcite in-situ is one method for the removal of
these anionic species, however, this method is limited by the relative affinities of
all anions in solution (carbonate prevents other anions to be intercalated). The
thermal activation of hydrotalcite removes water and carbonate, therefore creating
a chemically more reactive structure (unstable when dehydrated and therefore
readily wants to react with water and anions to reform a hydrotalcite structure),
which allows for a larger concentration of anions to be removed from solution.
The thermally activated hydrotalcite is also more susceptible to removing larger
anionic species (organics) due to this increase in reactivity.
The thermal activation of hydrotalcite materials dehydrates the structure,
removing water and other volatile anions from the interlayer region. [12] Re-
hydration of the thermally activated hydrotalcite with an aqueous solution returns
the hydrotalcite to its original structure. Therefore, any anion present in the re-
hydration solution has the potential to be adsorbed into the thermally activated
hydrotalcite. The removal of anions from solution is dependent on their relative
affinities for the hydrotalcite, since those with higher affinities will be removed
preferentially.
224
The thermal activation of red mud removes adsorbed water as well as dehydrating
the oxide species found within red mud. [12] Therefore, the addition of thermally
activated red mud to solutions containing oxy-anions of transition metals will
remove these anionic species via adsorption on the red mud particles. However,
thermal activation of SWN-RMS will result in the dehydration of the red mud
oxides and the newly formed hydrotalcite. Therefore, thermally activated
SWN-RMS will remove anionic species via adsorption (red mud particles and on
the external surface of hydrotalcite) and intercalation (hydrotalcite). The
combination of these mechanisms will remove a larger quantity of anionic species
from solution, and therefore thermally activated SWN-RMS is a more efficient
adsorbent material.
This chapter investigates the viability of hydrotalcites, red mud, and SWN-RMS
for the removal of arsenate and vanadate, commonly found in Bayer liquors. The
removal of arsenate and vanadate from Bayer liquor allows for the residues safe
disposal and storage. Thermally activated hydrotalcite or SWN-RMS (0.5 g) was
mixed with each anionic solution (10 mL) and then analysed by ICP-OES to
determine the uptake capacity of each material.
225
2. Effect of Mg:Al cationic ratio on anion removal for mixed anion solutions
Synthetic Mg,Al hydrotalcites with Mg:Al ratios of 2:1, 3:1, and 4:1 were
prepared and thermally activated. Five solutions (ultra pure water with pH 8)
containing different concentrations of arsenate, vanadate, and molybdate (in the
same solution) were treated by the three thermally activated hydrotalcites. The
effectiveness of the thermal activation of the three synthetic hydrotalcites for the
removal of arsenate, vanadate, and molybdate are summarised in Figures 6.1-6.3.
Bayer hydrotalcite was also thermally activated and used to treat the same
solutions as the three synthetic hydrotalcites (Fig. 6.4). Comparison of the uptake
capacity between thermally activated Bayer hydrotalcite and synthetic
hydrotalcites will provide an indication of the Mg:Al ratio of Bayer hydrotalcite.
The results clearly indicate that the order of affinity for these particular anions is
arsenate, vanadate, and molybdate, independent of the Mg:Al ratio. However,
increasing the Mg:Al ratio to 4:1 improves the anion uptake capacity from
solution. The removal of molybdate is much less effective than arsenate or
vanadate, especially at higher concentrations. The molybdate oxy-anion is too
large to be effectively intercalated, so is primarily adsorbed on the external
surface of hydrotalcite. There are more intercalation sites than adsorption sites in
hydrotalcite, meaning smaller concentrations are removed from solution. The
complete removal of arsenate and vanadate can be obtained for solutions
containing up to 25 ppm for 2:1 and 3:1 hydrotalcites, and greater than 100 ppm
for 4:1 hydrotalcites and Bayer hydrotalcite.
XRD of the synthetic carbonate hydrotalcites has shown that the 4:1 hydrotalcite
has a larger d(003) spacing, which is proposed to allow many anions to be
intercalated (Fig. 3.17). The 2:1 and 3:1 hydrotalcite had the same d(003) spacing,
and therefore the same percentage of anions are removed for both. Bayer
hydrotalcite also had a higher d(003) spacing (7.93 Å) than the 2:1 and 3:1
synthetic hydrotalcites (Fig. 4.11). The d(003) spacing of Bayer hydrotalcite and
4:1 synthetic hydrotalcite is 7.93 Å, compared to 7.67 Å for 2:1 and 3:1 synthetic
hydrotalcites. Therefore, it is not surprising that the 4:1 hydrotalcite and Bayer
226
hydrotalcite both show similar uptake capacities for arsenate, vanadate, and
molybdate from
Figure 6.1: Mixed solution removal capacity of thermally
activated 2:1 synthetic hydrotalcite.
Figure 6.2: Mixed solution removal capacity of thermally
activated 3:1 synthetic hydrotalcite.
227
Figure 6.3: Mixed solution removal capacity of thermally
activated 4:1 synthetic hydrotalcite.
Figure 6.4: Mixed solution removal capacity of thermally
activated Bayer hydrotalcite.
228
solution. This indicates that Bayer hydrotalcite has a Mg:Al ratio closer to 4:1
than 3:1.
For each anionic species, the percentage uptake is seen to decrease as the
concentration in solution increases. This is because a limited number of
intercalation sites exist within the hydrotalcite interlayer. It is proposed anions of
high affinity are intercalated rapidly, whilst lower affinity anions are intercalated
as space in the interlayer region becomes available (when the interlayer region
rearranges to form more aligned structures). The intercalation of anions is also
limited by the overall charge of the hydrotalcite, where totally neutralised
structures do not remove anions from solution. Once electrostatic neutrality is
reached, anions can only be removed from solution if they have a higher affinity
for the interlayer, and are exchanged for lower affinity anions (anion exchange).
2.1. 2:1 synthetic hydrotalcite
The uptake percentages of arsenate and vanadate are relatively similar for all
initial concentrations tested (Fig. 6.1). Percentage uptake values for molybdate are
significantly lower than those of arsenate and vanadate. At 5 ppm, 100 % of
arsenate and vanadate are removed. As the initial concentration of all anions in
solution increases from 25 to 50 ppm, the uptake capacity ability of the 2:1
hydrotalcite decreases. At 100 ppm, the percentage uptake of arsenate and
vanadate decreased to 60 and 50 %, respectively. The use of thermally activated
2:1 hydrotalcite for the uptake of molybdate is not an efficient removal technique.
In the presence of other anions, molybdate anions are hardly removed. The high
competition of arsenate and vanadate for sites in the interlayer and the external
surfaces of the hydrotalcite structure limited the uptake of molybdate. The lower
affinity of molybdate, compared to arsenate and vanadate, is due to its larger
anionic radius and smaller charge density.
2.2. 3:1 synthetic hydrotalcite
The percentage removal of the three anions from solutions with increasing anion
concentrations shows a similar overall trend to that of the thermally activated 2:1
229
hydrotalcite structures (Fig 6.2). The 3:1 thermally activated hydrotalcite shows
higher percentage uptake values than the 2:1 thermally activated hydrotalcite. The
lowest concentration of anions in solution, 5 ppm, showed almost 100 % removals
for all three anions. Arsenate also showed 100 % removal for 25 ppm solutions,
whereas, a slight decrease in vanadate uptake is observed and a significant
decrease in molybdate uptake is observed. The order of affinity is arsenate,
vanadate, and molybdate. Therefore, higher percentage removals for arsenate are
expected. As the initial concentration of all three anions in solution increased to
50 ppm, significant decreases in removal ability are observed. This reduction is
due to a limited number of intercalation sites in the hydrotalcite interlayer.
2.3. 4:1 synthetic hydrotalcite
The thermally activated 4:1 hydrotalcite is the most effective in the removal of
equal concentrations of arsenate, vanadate, and molybdate from contaminated
solutions (Fig. 6.3). The removal of all three anionic species is considerably
higher than the other two hydrotalcite ratios. Arsenate and vanadate are almost
completely removed from solution for all concentrations tested, whilst the
removal of molybdate is also significantly higher (minimum of 60 % removal for
the 100 ppm solution compared to 15-20 % for the 2:1 and 3:1 hydrotalcites). The
increased removal ability of the thermally activated 4:1 hydrotalcite is believed to
be due to the higher magnesium content resulting in a larger number of strong
chemical bonds between anions and magnesium cations, compared with weaker
bonding of anions with aluminium anions. This is caused by increased anionic
polarisation by the higher charge density of aluminium ions versus magnesium,
thus reducing the ionic character of the bonds.
2.4. Bayer hydrotalcite
Like the 4:1 thermally activated hydrotalcite, the thermally activated Bayer
hydrotalcite exhibited 100 % uptake of arsenate and vanadate, and relatively high
removal percentages for molybdate (Fig. 6.4). The similarity of the graphs
suggests that the chemical characteristics of the 4:1 hydrotalcite and Bayer
hydrotalcite interlayer’s are similar. ICP analysis of the Mg:Al ratio for the Bayer
230
Figure 6.5: Comparison of the removal abilities of thermally activated red mud
and seawater neutralised red mud for the removal of arsenate, vanadate, and
molybdate.
231
hydrotalcite supports that the structure will have similar characteristics to both the
3:1 and 4:1 samples, as its Mg:Al ratio falls between the two. XRD suggests that
Bayer hydrotalcite resembles the 4:1 hydrotalcite more. All three anions are
present in bauxite refinery residues in relatively low concentrations. The treatment
of residue liquor with thermally activated Bayer hydrotalcite should remove most
of these anions from the liquor, provided the concentrations of higher affinity
anions (carbonate and sulfate) are minimised.
Treating the residue liquor with thermally activated Bayer hydrotalcite could
potentially double the removal of anions from the residue, since the initial
formation of Bayer hydrotalcite during neutralisation would be supplemented with
this secondary treatment. Thermally activated Bayer hydrotalcite should also
remove a larger concentration of low affinity anions from the treated liquor, since
high affinity anions are removed in the initial treatments.
3. Red mud and seawater neutralised red mud
Raw red mud contains very small concentrations of hydrotalcite, therefore any
significant removal of anionic species from solution is due to adsorption onto the
external surfaces of the red mud particles and other Bayer-derived components.
The capacity for the removal of anionic species through intercalation is far greater
than that for adsorption (Fig. 6.5). Red mud is found to remove no more than
55 % for any concentration, while SWN-RMS removed almost
100 % of arsenate and vanadate up to 50 ppm. Therefore, the presence of
hydrotalcite significantly improves the removal of anions from contaminated
solutions. The use of thermally activated SWN-RMS for the removal of
molybdate is not favourable, with a maximum uptake capacity of 10 % being
obtained for 5 ppm solutions (Fig. 6.5). The capacity of red mud to remove anions
through adsorption is dependent on the initial concentration of the contaminated
solution and the surface area of red mud particles that adsorb anions. The presence
of thermally activated hydrotalcite in SWN-RMS significantly improves anion
removal by increasing the amount of intercalation reactions, as well as removing
anions by adsorption onto the external surfaces of the hydrotalcite structure.
232
4. Chapter summary
Solutions containing arsenate, vanadate, and molybdate have been treated with
thermally activated synthetic hydrotalcite, Bayer hydrotalcite, red mud, and
seawater neutralised red mud (SWN-RMS) to quantify their anion uptake
capacities. The order of affinity for all thermally activated hydrotalcite materials
is arsenate, vanadate, and molybdate. Significant removal values of arsenate and
vanadate, concentrations less than 100 ppm, can be achieved using 4:1 Mg,Al
hydrotalcite structures. The same results are observed for Bayer hydrotalcite. In
most cases the Bayer hydrotalcite performance was similar to the 4:1 synthetic
hydrotalcite structure. This indicates that Bayer hydrotalcite has similar metal
layer characteristics to this material. XRD also confirmed Bayer hydrotalcite has
the same interlayer distance as 4:1 hydrotalcite, indicating Bayer hydrotalcite has
a Mg:Al ratio closer to 4:1 than 3:1.
The seawater neutralisation of red mud vastly improves the uptake capacity of red
mud, with the percentage of anion removal almost doubling. The increased
removal ability is due to the formation of Bayer hydrotalcite during the seawater
neutralisation process. Thermally activated seawater neutralised red mud removes
anions from solution through: 1) the adsorption of anions onto the external
surfaces, and 2) the intercalation of anions into the hydrotalcites lamellar
structure. Anions are only removed by adsorption for thermally activated red mud.
It should be noted that if RMS is neutralised with seawater with additional
MgCl2·6H2
O, an increase in the quantity of hydrotalcite would be observed,
which would therefore increase the percentage of anions that can be removed
from solution.
233
5. References
[1] R.L. Frost, A.W. Musumeci, Journal of Colloid and Interface Science, 302 (2006) 203-206. [2] Y. Kiso, Y.J. Jung, T. Yamada, M. Nagai, K.S. Min, Water Science & Technology: Water
Supply, 5 (2005) 75-81. [3] H. Hirahara, S. Aisawa, H. Sato, S. Takahashi, Y. Umetsu, E. Narita, Nendo Kagaku, 45
(2005) 6-13. [4] N. Murayama, M. Tanabe, R. Shibata, H. Yamamoto, J. Shibata, Kagaku Kogaku
Ronbunshu, 31 (2005) 285-290. [5] V.K. Gupta, S. Sharma, Environmental Science and Technology, 36 (2002) 3612-3617. [6] H. Genc, J.C. Tjell, D. McConchie, O. Schuiling, Journal of Colloid and Interface Science,
264 (2003) 327-334. [7] H.S. Altundogan, S. Altundogan, F. Tumen, M.Bildik, Waste Management, 22 (2002) 357-
363. [8] D. McConchie, H. Genc, J.C. Tjell, Journal of Colloid and Interface Science, 271 (2004)
313-320. [9] N.W. Menzies, I.M. Fulton, W.J. Morrell, Journal of Environmental Quality, 33 (2004)
1877-1884. [10] M.S. Rahaman, A. Basu, M.R. Islam, Bioresource Technology, 99 (2008) 2815-2823. [11] S.C. Grocott, L.E. Jefferies, T. Bowser, J. Carnevale, P.E. Jackson, Journal of
Chromatography, 602 (1992) 257-264. [12] F. Malherbe, J.p. Besse, Investigating the Effects of Guest-Host Interactions on the
Properties of Anion-Exchanged Mg-Al Hydrotalcites, Journal of Solid State Chemistry. 155 (2000) 332-341.
234
CHAPTER 7
Conclusions and recommendations for future
work
235
1. Conclusions
Bauxite refinery residues are a highly complicated system, with 15 mineralogical
phases being detected in some cases. Therefore, utilising synthetic counterparts
simplified the system and allowed the identification of triggers that cause
reversion. Three components within bauxite refinery residues have been identified
as triggers that contribute to reversion: 1) tricalcium aluminate hexahydrate, 2)
hydrocalumite, and 3) calcium hydroxide. Other components within the residue
have been shown to have an impact on the pH of solution (increase or decrease),
however, these variations in pH are not associated with the phenomenon of
reversion. Hydrotalcite that forms from the SWN process has been found not to
contribute to reversion, and is stable throughout the neutralisation process.
High concentrations of calcium hydroxide in solution cause an increase in
hydrocalumite formation. Therefore, the presence of calcium hydroxide in the
residue contributes to pH reversion in two ways: 1) dissolution of calcium
hydroxide, and 2) the formation and subsequent dissolution of hydrocalumite. The
dissolution of calcium hydroxide contributes to pH reversion, while the
dissolution of hydrocalumite contributes to both pH and aluminium reversion.
Hydrocalumite formed in all synthetic liquors (reaction of calcium in seawater
with aluminate and hydroxide ions in residue liquor), and caused an increase in
calcium, aluminium, and hydroxide ion concentrations. However, formation of
hydrotalcite immediately precipitates these newly dissolved ions out of solution,
making it difficult to quantify the exact amount of hydrocalumite that formed.
However, the quantity of hydrocalumite that forms during the SWN process is
such that when the dissolution of hydrocalumite occurs, there is insufficient
magnesium left in solution to fully neutralise OH- and aluminium ions.
Hydrocalumite also becomes problematic when the residue contains high
Ca(OH)2 levels. The increase in calcium ions was successfully observed, as the
mechanism for the removal of calcium ions (CaCO3
) is considerably slower than
hydrotalcite formation.
The dissolution of TCA is proposed to involve a number of reactions occurring
simultaneously. The two major reactions involved in the dissolution of TCA are:
236
1) with sodium carbonate to form sodium hydroxide, sodium aluminate ions, and
calcium carbonate, and 2) with sodium hydroxide to form sodium aluminate and
calcium hydroxide. The reaction of sodium carbonate with TCA is believed to
occur initially, which releases NaOH into solution, until the Na2CO3
concentration is depleted. The release of 4 moles of NaOH during this process
therefore increases the pH of solution. However, this is partially off-set by the
consumption of 2 moles of NaOH in the other dissolution reaction. The
dissolution of calcium hydroxide formed by the reaction of TCA with sodium
hydroxide then causes an increase in OH- ions (pH). When there are no reactions
removing OH-
ions from solution, the pH increases significantly (reversion). TCA
has been found to be a major contributor to pH and aluminium reversion.
Therefore, the minimisation of TCA compounds in the residue should reduce the
extent of reversion.
This investigation has also identified optimisation processes that could be
employed to reduce the potential environmental risk of the residue. The first
utilises the formation of hydromagnesite at 75 °C during the neutralisation
process. The formation of hydromagnesite reduces the final pH of solution,
however, additional MgCl2·6H2O is required to ensure hydrotalcite can still form
to remove any aluminium that is re-dissolved back into solution. It has also been
shown that increased whewellite concentrations reduce the final pH of solution
via the adsorption of OH- ions onto the surface of whewellite particles. Therefore,
increasing the concentration of whewellite should reduce the pH of solution. As
previously mentioned, Ca(OH)2 concentrations need to be kept to a minimum to
prevent hydrocalumite formation and Ca(OH)2
dissolution. Finally, small
concentrations of TCA have been shown to reduce the final pH of solution by the
formation of additional hydrotalcite after TCA dissolution in magnesium rich
solutions.
This investigation has identified two methods for the minimisation of pH and
aluminium reversion: 1) increasing the seawater neutralisation ratio, and 2)
addition of magnesium chloride. Increasing the volume of seawater showed a
significant reduction in reversion. The addition of MgCl2·6H2O to seawater
confirmed that the decrease in aluminium and hydroxyl ions can be attributed to
237
the formation of hydrotalcite and not simply a dilution effect (previously thought
to occur for increased seawater volumes). Therefore, increasing the magnesium
concentration, either by additional seawater or MgCl2·6H2
O, can be used to
minimise reversion and reduce the final pH of the slurry for safe disposal. Both
methods utilise the formation of hydrotalcite to remove both hydroxide and
aluminium ions from solution permanently. By ensuring there is always an excess
of magnesium ions in solution, reversion can be prevented.
Initial theories into the cause for reversion believed hydrotalcite formed during
the neutralisation process was a major contributor, either through the partial or
complete dissolution of the structure. Synthetic hydrotalcites have been
successfully synthesised using the same conditions used in the neutralisation
process. Due to the large pH range over which neutralisation occurs, hydrotalcites
have been synthesised with variable cationic ratios and characterised. Bayer
hydrotalcites have then been prepared under the same conditions and compared to
the synthetic hydrotalcites. The results indicate that Bayer hydrotalcites
predominantly having Mg:Al ratios between 3 and 4 form. XRD and thermal
activation experiments have shown that Bayer hydrotalcite and 4:1 hydrotalcites
react similarly in the removal of anions. This suggests that Bayer hydrotalcite has
a Mg:Al ratio closer to 4:1 than 3:1. Bayer hydrotalcites are the closest
representation of the hydrotalcite that forms in bauxite refinery residues.
This investigation has shown that Bayer hydrotalcite is not a contributor to
reversion. Its presence in synthetic Bayer liquor and real Bayer liquor did not
affect the pH of the neutralised solution or the aluminium concentration. The
extensive study on the hydrotalcite structures has been completed to help
understand the mechanism of inclusion of arsenate, vanadate, and molybdate. The
combination of a number of instrumental techniques enabled the identification of
the primary mechanism in the removal of these anions. The predominant removal
mechanism from solution for these anionic species during the precipitation of
hydrotalcite is adsorption for 2:1 and 4:1 hydrotalcites, and intercalation for 3:1
hydrotalcites. Intercalation of molybdate anions is far less effective than for the
other two anions, and believed to be a 2-step mechanism.
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This investigation looked at using thermally activated hydrotalcite as a means for
the removal of arsenate, vanadate, and molybdate. The thermal activation process
removes anions from the hydrotalcite interlayer, thus creating an adsorbent
material that is considerably more chemically reactive. This increase in reactivity
is due to a need to neutralise the positive hydrotalcite layers. It has been found
that thermally activated 4:1 Mg,Al hydrotalcites remove at least 100 ppm of
arsenate and vanadate from solutions. The removal of molybdate has been found
to be minimal, with percentage removals less than half that of the other two
anionic species. Comparison of results for thermally activated synthetic and Bayer
hydrotalcite shows that the chemistry of the Bayer hydrotalcite most resembles
the 4:1 synthetic hydrotalcite. Thermally activated red mud and seawater
neutralised red mud (containing Bayer hydrotalcite) were also analysed and
showed that thermally activated seawater neutralised red mud significantly
improves the removal ability of this material. The significant increase is due to a
combination of intercalation and adsorption reactions compared to just the
adsorption of the anions on the dehydrated red mud particles. Therefore, this
residue material has the potential to be used for a water purification technique.
239
2. Recommendations
It is suggested that further investigations into using components already found
within bauxite refinery residues, as a means to reduce the final pH of the
neutralised product, should be considered. Such components include whewellite
(organic impurity) and tricalcium aluminate hexahydrate (TCA), both of which
have been shown to reduce the final pH of the residue. The use of higher
neutralisation temperatures in conjunction with MgCl2·6H2
O is also believed to
produce a more environmentally friendly residue. Other areas of future work
include using hydrotalcite, Bayer hydrotalcite, and seawater neutralised red mud
within the alumina industry for a water purification technique or for
causticisation. These methods need to be explored in full depth to ensure that the
material used in the method does not have an adverse effect on another area
within the refinery or during the disposal of the residue.
TCA has the potential to reduce the final pH of solution, as long as the
magnesium concentration in solution is greater than the resulting concentration of
aluminium ions in solution caused by the dissolution of TCA. The pH of solution
is reduced through the formation of additional hydrotalcite. By increasing the
concentration of TCA in solution, a greater amount of TCA is dissolved into
solution releasing aluminium and hydroxide ions. The increase in these ions
facilitates the formation of hydrotalcite, as long as there is an adequate
magnesium concentration. As soon as the magnesium concentration is depleted,
pH and dissolved aluminium levels will rise. Therefore, optimisation of TCA
addition into the residue would be required before this process is viable.
Whewellite shows the greatest potential to reduce the pH of the residue before
disposal. However, investigations into the effect of high organic content in the
previous stages of the Bayer process, particularly production of alumina, needs to
be undertaken. An investigation into the stability of whewellite, in regards to the
removal of hydroxide ions, also needs to be conducted. This is to ensure that the
hydroxide ions are not released back into the slurry after disposal into tailings
dams, thus causing pH reversion at a later stage.
240
The actual temperature that SWN occurs at should be investigated further. It has
been shown that at 75 °C a lower neutralisation point is obtained, due to the
formation of hydromagnesite. However, the pH after the neutralisation point rises
by 11 % compared to 10 % for 55 °C. This increase is due to a lack of magnesium
ions used to neutralise OH- and aluminium ions that are released by the
dissolution of TCA and hydrocalumite. Therefore, increasing the SWN
temperature is only viable if additional seawater or MgCl2·6H2
O is used. A full
investigation into the benefits and disadvantages of higher SWN temperatures
needs to be conducted.
Finally, an investigation into using Bayer hydrotalcites for causticisation would be
beneficial to the alumina industry, due to the abundance of this material at the
refinery.
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APPENDIX
A.1 Calculation of water in the carbonate hydrotalcite – Chapter 3.
Calculation of water content for carbonate intercalated hydrotalcite:
Composition: Mg6Al2(OH)16CO3 * xH2
Removing water up to 235°C: 25.3 mg that is 1.404 mmol
O
Remaining dehydrated mineral up to 235°C: 129.73 mg that is 0.244 mmol
Molar mass of dehydrated mineral: 531.99 g/mol
Calculation of x:
1 mol dehydrated mineral – x mol H2
0.244 mol dehydrated mineral – 1.404 mol H
O
2
O
x = 5.75 ~ 6 mol
Formula: Mg6Al2(OH)16CO3 * 6 H2
Steps of water liberation according to the decomposition steps up to 235°C:
O
1. step: 1.097 mol
2. step: 2.600 mol
3. step: 2.304 mol