stability of hydrotalcites formed from bayer … · 2012. 2. 15. · • tricalcium aluminate...

$: STABILITY OF HYDROTALCITES FORMED FROM BAYER … · 2012. 2. 15. · • Tricalcium aluminate hexahydrate • Vanadate • X-ray diffraction . ix LIST OF ABBREVIATIONS. AsO. 4. 3-BHT$
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


CHAPTER 3

Synthesis and characterisation of synthetic hydrotalcites


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



xiv

CHAPTER 4

Synthesis and characterisation of Bayer hydrotalcites


2. Identification of hydrotalcite formation in seawater neutralised

red mud … 131


2.2. Thermal analysis … 133

3. Bayer hydrotalcites formed during the seawater neutralisation of

bauxite refinery residues … 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


… 159


… 161



xv

CHAPTER 5

Reversion


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



CHAPTER 6

Thermally activated seawater neutralised red mud used for the

removal of arsenate, vanadate and molybdate from aqueous

solutions


2. Effect of Mg:Al cationic ratio on anion removal for mixed

anion solutions … 225

2.1. 2:1 synthetic hydrotalcite … 228



2.4. Bayer hydrotalcite … 229

3. Red mud and seawater neutralised red mud … 231



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


Figure 3.10: Raman spectra of the cation deformation modes of molybdate


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-


of HT(CO

).

32-,AsO4

3-



).

32-,AsO4

3-


of HT(AsO

).

43-


decomposition of HT(AsO

).

43-


of HT(CO

).

3,VO43-



).

3,VO43-


of HT(VO

)

43-


decomposition of HT(VO

).

43-


of HT(CO

).

3,MoO42-



).

3,MoO42-


of HT(MoO

).

42-


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)


2


varying concentrations of hydrocalumite.

before and after SWN.

Figure 5.16: Aluminium concentration in solution after the SWN of SNL with


Figure 5.17: Magnesium concentration in solution after the SWN of SNL with


Figure 5.18: Calcium concentration in solution after the SWN of 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.


varying concentrations of Bppt.



Figure 5.25: Magnesium concentration in solution after the SWN of SNL with



varying concentrations of whewellite.


varying concentrations of whewellite.


varying concentrations of sodalite.


varying concentrations of sodalite.


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


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


Figure 6.3: Mixed solution removal capacity of thermally activated 4:1


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.


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-


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.


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


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


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


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


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


thermogravimetric analysis of HT(AsO43-

).


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


thermogravimetric analysis of HT(CO3,VO43-

).


gases in the thermal decomposition of HT(CO3,VO43-).

115


thermogravimetric analysis of HT(VO43-

).


gases in the thermal decomposition of HT(VO43-).

116


thermogravimetric analysis of HT(CO3,MoO42-

).


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


thermogravimetric analysis of HT(MoO42-).


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.


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


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


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)


2-)(s) → MgAl2O4(s) + 5MgO(s) + (CO2,SO2) (g) + 8H2O(g) +

O2(g)

Mg

4:1 hydrotalcite:


2-)·xH2O(s) Mg8Al2(OH)18(CO32-,SO4

2-)(s) + xH2O

Mg

(g)


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.


Mg 10.60 2+ 10.39 11.18 10.72 3.34

Al 3.15 3+ 3.06 3.41 3.21


Mg 12.43 2+ 11.27 12.32 12.01 3.77

Al 3.30 3+ 2.91 3.35 3.19


Mg 10.08 2+ 11.23 11.74 11.02 3.44

Al 2.81 3+ 3.25 3.56 3.21


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


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


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)


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)


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


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,


Figure 5.7: Calcium concentration after neutralisation,


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


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


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



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


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


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


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


SWN of SNL with varying concentrations of whewellite.


SW and SNL with varying concentrations of sodalite.

209


SWN of SNL with varying concentrations of sodalite.


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


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


Figure 5.33: Magnesium concentration for SWN synthetic SNL containing

0.10M TCA with additional MgCl2·6H2


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.


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.


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,


[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.


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.



227




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.


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.


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.

238

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.

241

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

stability of hydrotalcites formed from bayer … · 2012. 2. 15. · • tricalcium aluminate...

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