Order Number 8824578
Properties of synthetic goethites and their effect on sulfate adsorption
Munoz, Miguel A., Ph.D.
The Ohio State University, 1988
Copyright ©1988by Munoz, Miguel A. All rights reserved.
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16. Other PROPERTIES OF SYNTHETIC GOETHITES AND THEIR EFFECT ON
SULFATE ADSORPTION
DISSERTATION
Presented in Partial Fulfillment of the Requirements for
the Degree Doctor of Philosophy in the Graduate
School of the Ohio State University
By
Miguel A. Munoz, B.S., M.S.
The Ohio State University
1988
Dissertation Committee: Approved by
Dr. J .M . Bigham
Dr. S.J. Traina ^ ' ~ ^ c / Adviser Dr. F.L. Himes Department of Agronomy
D r . R .T . Tettenhors t Copyright by Miguel A. Munoz 1988 DEDICATION
To all of you who build dreams and are strong enough to wake up and make those dreams a reality.
To my Homeland, my Country, ray People.
DEDICATORIA
A todos aquellos forjadores de suenos que tienen la suficiente fortaleza para despertar y hacer esos suenos una realidad.
A mi Patria, mi Pueblo, mi Gente. ACKNOWLEDGMENTS
I express my most sincere appreciation to Dr. Jerry
M. Bigham, Dr. Sam J. Traina and Dr. Trevor G. Arscott for their valuable help, guidance and insight throughout the research. I am also grateful to Dr. Frank L. Himes and
Dr. Rodney T. Tettenhorst for their suggestions and comments as members of the reading committee.
I am very grateful to the Department of Agronomy and the Agriculture Experiment Station of the University of
Puerto Rico for their financial support while pursuing this degree.
A special appreciation is extended to Dr. Clifford T.
Johnston at the University of Florida for his valuable help with the FTIR study. I would also like to thanks Dr.
William F. Jaynes who very kindly help me with the XRD analysis and the TEM study.
I wish to specially acknowledge the love and support of my wife, Fuly. We have shared the good and the bad times and always managed to keep a smile at the end of the road.
iii VITA
April 14, 1956 ...... Born - Aguada, Puerto Rico
1978 ...... B.S. in Agronomy University of Puerto Rico Mayaguez, Campus
1983 ...... M.S. in Soil Fertility Thesis Topic: Phosphorus Absorption by Corn at Different Moisture Levels The Ohio State.University
1976-1981 ...... Research Assistant Sugarcane Breeding Program, Agricultural Experimental Station University of Puerto Rico
1984-1986 ...... Graduate Teaching Associate, Romance Languages Department The Ohio State University
FIELDS OF STUDY
Major field: Agronomy TABLE OF CONTENTS
Page
DEDICATION...... ii
ACKNOWLEDGMENTS...... iii
VITA...... iv
LIST OF TABLES...... ix
LIST OF FIGURES...... xii
INTRODUCTION...... 1
CHAPTER
I. A REVIEW OF IRON OXIDES INSOILS ...... 7
1.1 Hematite (a-Fe203)...... 10 1.2 Lepidocrocite (8-FeOOH)...... 11 1.3 Ferrihydrite (5Fe2Q3 . 9H20)...... 13 1.4 Feroxyhite (6'-FeOOH)...... 15 1.5 Maghemite (tf-Fe203)...... 17 1.6 Magnetite (FesOO ...... 19 1.7 Goethite (a-FeOOH)...... 19
II. PROPERTIES AND SYNTHESIS OF GOETHITE...... 22
2.1 Structure of Goethite...... 22 2.2 Composition of Goethite...... 26 2.3 Goethite Morphology...... 29 2.4 Thermal Properties of Goethite...... 37 2.5 Synthesis and Formation of Goethite...... 39
2.5.1 Synthesis of Goethites...... 39 2.5.2 Factors Affecting the Formation of Goethite...... 40
v 2.6 Materials and Methods...... 48
2.6.1 Sample Preparation...... 48
2.6.1.1 Al-substituted Goethites...... 48 2.6.1.2 Hydrothermally Treated Goethites...... 51 2.6.1.3 Nipe Soil Clay...... 51
2.6.2 Characterization of the Synthetic Oxides and Nipe Soil Clay...... 53
2.6.2.1 X-ray Powder Diffraction...... 53 2.6.2.2 Calculation of Unit Cell Dimensions and Mean Crystallite Dimensions (MCD).... 54 2.6.2.3 Transmission Electron Microscopy...... 57 2.6.2.4 Differential Scanning Calorimetry (DSC)...... 59 2.6.2.5 Surface Area...... 59 2.6.2.6 Determination of Total Iron, Total Aluminum and Oxalate Soluble Iron...... 59 2.6.2.7 Reductant-soluble Iron Content of Nipe Clay...... 60
2.7 Results and Discussion...... 61
2.7.1 Synthesis of Al-substituted Goethites.... 61 2.7.2 X-ray Diffraction Studies...... 64
2.7.2.1 Effect of A1 substitution on lattice spacing, unit cell dimensions and mean crystallite dimensions (MCD)...... 66
2.7.3 Surface Area of Al-goethites...... 73 2.7.4 Transmission Electron Microscopy (TEM)... 76 2.7.5 Dehydroxylation Properties of Al- goethites...... 80 2.7.6 Conclusions of the Synthesis and Characterization of Al-substituted Goethites...... 84 2.7.7 Synthesis and Characterization of Hydrothermally-treated Goethites...... 85
vi 2.7.8 Conclusions of the Synthesis and Characterization of Hydrothermally- treated Goethites...... 96
2.7.9 Characterization of Nipe Clay...... 96
III. INFRARED SPECTROSCOPY STUDY...... 101
3.1 Infrared (IR) Characterization of Goethite: a Review...... 103 3.2 Infrared Spectroscopy Studies of Anion Adsorption...... 108 3.3 Infrared Spectroscopy of Sulfate and Sulfate Complexes...... 110 3.4 Materials and Methods...... 113
3.4.1 FTIR Characterization of the Adsorbents.. 114 3.4.2 FTIR Study of Sulfate Adsorption on Goethite...... 116
3.5 Results and Discussion...... 119
3.5.1 Characterization of the Adsorbents...... 119 3.5.2 FTIR Study of Sulfate Adsorption on Goethite...... 134 3.5.3 Conclusions from the IR Study...... 156
IV. ADSORPTION ISOTHERMS...... 157
4.1 Surface Charge of Oxides as Related to Sulfate Adsorption...... 157 4.2 Adsorption Isotherms...... 163 4.3 Materials and Methods...... 168
4.3.1 Preparation of Sulfate Adsorptio Isotherms...... 168
4.4 Results and Discussion...... 171
4.4.1 Al-substituted Goethites...... 171 4.4.2 Hydrothermally-treated Goethites and the Nipe Clay...... 180 4.4.3 Sulfate Adsorption and OH Release...... 187 4.4.4 Conclusions...... 193
vii V. SUMMARY...... 196
BIBLIOGRAPHY...... 199
APPENDIX
A. SULFATE ADSORPTION ISOTHERMS DATA...... 215
viii LIST OF TABLES
Table No. Page
1.1 Some crystallographic data for the major Fe oxide minerals occurring in soils...... 9
2.1 Aluminum substitution in goethites...... 27
2.2 Chemical composition of soil goethites.... 30
2.3 Solution volumes for the preparation of the alurainum-goethite series...... 50
2.4 Some physical and chemical properties of Al-substituted goethites...... 63
2.5 Unit cell dimensions and mean cystallite dimensions (MCD) of Al- substituted goethites...... 68
2.6 Average widths and lengths of particles from Al-substituted goethites as measured from electron micrographs and ratios of widths from XRD(MCDb) and TEM... 78
2.7 Oxalate soluble Fe content, surface area and mineralogy of hydrothermally-treated goethites...... 86
2.8 Unit cell dimensions and mean crystallite dimensions (MCD) of hydrothermally-treated goethites...... 89
ix Average widths and lengths of particles from hydrothermally-treated goethites as measured from electron micrographs and ratios of widths from XRD (MCDb) and TEM.. 92
Some chemical and physical properties of Nipe clay (< 2um)...... 99
Infrared band positions of sulfate in selected complexes...... 112
IR absorption frequencies of Al- goethites ...... 124
IR absorption frequencies of hydrothermally-treated goethites...... 131
Sulfate adsorption capacity of various oxides...... 165
Sulfate adsorbed by Al-substituted goethites at the highest SCU2- equilibrium concentration...... 175
Variation in S042- adsorbed with aluminum substitution...... 178
Sulfate adsorbed by hydrothermally- treated goethites at the highest S04"2 equilibrium concentration...... 182
Sulfate adsorption isotherms data (mraol/Kg) for samples S-0, S-10, S-30 and S-80...... 216
Sulfate adsorption isotherms data (mmol/Kg) for samples S-120, S-150A, and S-150B...... 217
Sulfate adsorption isotherms data (mmol/m2 x 10-6) for samples S-0, S-10, S-30, and S-80...... 218
x A.4 Sulfate adsorption isotherms data (mmol/m2 x 10-3) for samples S-120, S-150A, and S-150B...... 219
A.5 Sulfate adsorption isotherms data (mmol/Kg) for samples S-OC, S-0C(50 ®C), S-0C(100 *C), and S-0C(150 °C)...... 220
A.6 Sulfate adsorption isotherms data (mmol/Kg) for samples S-OC, S-0C(200 °C) and Nipe clay...... 221
A.7 Sulfate adsorption isotherms data (mmol/ra2 x 10-3) for samples S-OC, S-0C(50 °C), S-0C(100 °C) and S-0C( 150 “C ) ...... 222
A.8 Sulfate adsorption isotherms data (mmol/m2 x 10-3) for samples S-0C(200 °C) and Nipe clay...... 223
A.9 OH released (mmol/m2 x 10“3) for samples S-0, S-10, S-30 and S-80...... 224
A.10 OH released (mmol/m2 x 10“3) for samples S-120, S-150A and S-150B...... 225
A.11 OH released (mmol/ra2 x 10-3) for samples S-OC, S-0C(50 °C), S-0C(100 °C) and S-0C( 150 °C)...... 226
A.12 OH released (mmol/ra2 x 10-3) for samples S-0C(200 °C) and Nipe clay...... 227
xi LIST OF FIGURES
Figure No. Page
2.1 Pictorial representation of goethite structure, with the hydrogen bonds shown as tubes...... 23
2.2 Structure of the (100) face of goethite (a-FeOOH)...... 25
2.3 Schematic representation of acicular goethite showing the three major crystallographic faces...... 33
2.4 Schematic representation of a multidomainic crystal of goethite showing domain boundaries...... 35
2.5 Schematic representation of the competitive process of goethite and hematite formation...... 41
2.6 Acid Digestion Bomb Parr Instrument Co.... 52
2.7 Expanded plot of 110 peak of synthetic goethite used for MCD determination...... 56
2.8 (IkO) planes used to calculate MCDa...... 58
2.9 X-ray diffraction pattern of synthetic goethite showing peak assignments for the different (hkl) planes...... 65
2.10 Shift in position of the 111 peak of goethite with aluminum substitution...... 67
xii 2.11 Unit cell dimensions of synthetic goethites as a function of aluminum substitution...... 70
2.12 Mean crystallite dimensions (MCD) of synthetic goethites as function of aluminum substitution...... 74
2.13 Surface area of synthetic goethites as a function of A1 substitution...... 75
2.14 Electron micrographs of Al-substituted goethites...... 77
2.15 Differential Scanning Calorimetry (DSC) pattern of Al-substituted goethites...... 81
2.16 Variation in the width of the 110 peak of goethite with crystallinity...... 87
2.17 Electron micrographs of hydrothermally- treated goethites...... 90
2.18 Electron micrographs of single crystals of hydrothermally-treated goethites...... 94
2.19 Differential Scanning Calorimetry (DSC) pattern of hydrothermally-treated goethites...... 95
2.20 X-ray diffraction pattern of Nipe clay.... 97
2.21 Differential Scanning Calorimetry (DSC) pattern of Nipe clay...... 100
3.1 Infrared spectra of goethite showing surface OH bands, which exchange with D 2O to give OD bands...... 106
3.2 Metal-sulfate complexes...... Ill
3.3 Schematic diagram of the optical system of the diffuse reflectance accessory 115
xiii 3.4 Schematic representation of the cylindrical internal reflectance cell (CIRCLE)...... 117
3.5 FTIR spectra of synthetic goethite using KBr pellets...... 120
3.6 FTIR spectra of goethite (Diffuse Reflectance)...... 121
3.7 FTIR spectra of goethite showing surface and bulk hydroxyl bands...... 122
3.8 Shift in the 6-OH and &-0H bending frequencies of goethite with Al- substitution...... 125
3.9 Shift in the V-OH bending vibrations of goethite with aluminum substitution.... 127
3.10 Shift in the 6-OH bending vibrations of goethite with aluminum substitution.... 128
3.11 Shift in the IR frequencies of goethite with hydrothermal treatment...... 132
3.12 Shift in the 6-OH bending vibrations of goethite with hydrothermal treatment... 133
3.13 Shift in the vOH (stretching) frequencies of goethite with hydrothermal treatment... 135
3.14 FTIR spectra of Nipe clay (Diffuse Reflectance)...... 136
3.15 FTIR spectra of 0.1 M sodium sulfate (CIRCLE Cell)...... 139
3.16 FTIR spectra of sulfate adsorbed on synthetic goethite (S-0) (CIRCLE Cell).... 140
3.17 Schematic representation of S04-oxide complexes...... 141
xiv 3.18 FTIR spectra of surface and bulk hydroxyl region of synthetic goethite (S-0) (CIRCLE Cell)...... 144
3.19 FTIR subtraction spectra of goethite (S-0) showing negative surface hydroxyl bands after S04 adsorption (CIRCLE Cell).. 146
3.20 FTIR spectra of sulfate adsorbed on synthetic goethite (S-150B) (CIRCLE Cell)...... 147
3.21 FTIR subtraction spectra of goethite (S-150B) showing negative surface hydroxyl bands after SO4 adsorption (CIRCLE Cell)...... 149
3.22 Effect of hydrothermal treatment and Al-substitution on crystal morphology 150
3.23 FTIR spectra of sulfate adsorbed on synthetic goethite (S-OC) (CIRCLE Cell)... 152
3.24 FTIR subtraction spectra of goethite (S-OC) showing negative surface hydroxyl bands after S04 adsorption (CIRCLE Cell).. 153
3.25 FTIR spectra of sulfate adsorbed on synthetic goethite (S-0C(20O 0C)) (CIRCLE Cell)...... 154
3.26 FTIR subtraction spectra of goethite (S-OC (200 °C)) showing negative surface bands after S04 adsorption (CIRCLE Cell).. 155
4.1 Induction of positive charge onto the surface brough about by adsorption of sulfate...... 162
4.2 Sulfate adsorption isotherms on Al- substituted goethites on a mass basis 172
4.3 Sulfate adsorption isotherms on Al- substituted goethites on a surface area basis ...... 173
xv 4.4 Sulfate adsorption isotherms on Al- substituted goethites showing the increase in sulfate adsorption at intermediate sulfate equilibrium concentration...... 177
4.5 Sulfate adsorption isotherms on hydrothermally-treated goethites on a mass basis...... 181
4.6 Sulfate adsorption isotherms on hydrothermally-treated goethites on a surface area basis...... 184
4.7 Sulfate adsorption isotherms on synthetic goethites and Nipe clay on a mass basis...... 185
4.8 Sulfate adsorption isotherms on synthetic goethites and Nipe clay on a surface area basis...... 186
4.9 OH released vs. sulfate equilibrium concentration (Al-substituted goethites).. 188
4.10 OH released vs. sulfate equilibrium concentration (Hydrothermally-treated goethites and Nipe clay)...... 189
4.11 OH released/S04 adsorbed ratio as a function of A1 substitution...... 191
4.12 OH released/S04 adsorbed ratio as a function of hydrothermal treatment...... 192
xvi INTRODUCTION
The adsorption of anions at mineral surfaces is very
important in soils because of the limitations that this
process may impose on the availability of plant nutrients
occuring naturally or in fertilizers. Common inorganic
soil anions include Cl", HC03", C032-, N03", S042-, HP042 >
H 2PO 4 ", OH", and F". In addition, some micronutrients
exist as anions (such as H 2BO3 - and M0 O 4 2"). In contrast
to cationic species that are usually attracted to mineral
surfaces by simple electrostatic attractions, most anions
are retained by more complex mechanisms.
An anion approaching a charged surface may be subject
to attraction or repulsion depending upon the net charge
on the surface. Since layer silicates in the clay
fraction of soils are normally negatively charged, anions
tend to be repelled from these mineral surfaces. However,
some soil constituents like the iron and aluminum oxides develop positive surface charge under certain conditions
1 2 and may play an important role in anion adsorption. These soil constituents, although present under a whole range of climatic conditions, are most abundant in the tropics where heavy rainfall and high temperatures promote rapid alteration of the soil parent material resulting in heavy loss of silicon and alkaline earth and alkali metals and accumulation of iron and aluminum oxides (Hasan et al.,
1970; Blair et al., 1980). Soils that are products of such weathering regimes have little permanent negative charge but exhibit considerable pH dependent charge.
Under acid conditions the net surface charge is neutral or positive and most cations leach readily (Mahilum et al.,
1970); however, 3ome oxyanions such as sulfate and phosphate are strongly retained (Fox, 1969; Hingston et al., 1972; Atkinson et al., 1974; El-Swaify and Sayegh,
1975; Rajan, 1978).
The surface charge of soil oxides is created by the unequal adsorption of potential determining ions (H30+ and
OH-) from aqueous solution (Parks and Bruyn, 1962;
Atkinson, 1967). The contribution made to surface charge by transfer of metal hydroxo complex ions is usually considered to be very small in relation to H30+ and OH- ion transfer. Based on this fact, surface charge measurements may be carried out by simple potentiometric titrations (Breeuwsraa and Lyklema, 1971). 3
The titratable surface charge is defined through:
a = F(Th + - rOH-)
where F is the Faraday constant and Th + and To h - are analytical surface excesses for H+ and OH" ions, respectively (Breewsma and Lyklema, 1973).
Potential determining ions (H30+, OH") are components of both the solvent and the solid phase and occupy the inner Helmholtz plane with respect to the metal ions of the oxide surface (Hingston, 1981). When the adsorbed anions are located either in the outer Helmholtz layer or the diffuse double layer the adsorption is termed non specific. Adsorption of these anions can be explained by electrostatic interactions alone. Exchange equations similar to those developed for cation exchange describe such reactions because non-specifically adsorbed anions are in the solution adjacent to the solid surface and are readily exchanged with the bulk solution. Most singly charged anions (e.g. N03" and Cl") are considered to be non-specifically adsorbed (Hingston et al., 1972; Harrison and Berkheiser, 1982). These anions are adsorbed in simple proportion to their concentration in solution. 4
Ions that are not constituents of the solid structure but that are adsorbed into the inner Helmholtz plane have been termed specifically adsorbed or ligand exchanged ions
(Hingston et. al., 1967; Berube and Bryun, 1968).
Specific adsorption can occur on surfaces initially carrying a net negative, positive or neutral charge. The classical example of specifically adsorbed anions is phosphate (White, 1980). Fluoride adsorption also conforms to the ligand exohange theory and is probably favored by the close similarity in size of F~ and OH" ions
(Bohn et al., 1979). Iron and aluminum oxides possess the ability to specifically interact with many anions and this characteristic gives oxidic soils an adsorption capacity for some anions that is much greater than predicted from electroneutrality alone.
In adsorption studies, the dream of any researcher would be to use natural materials as experimental media and to succesfully interpret the reactions observed; however, the use of natural materials is subject to many problems because of the heterogeneity of the solid phase.
A soil is a mixture of mineral and organic adsorbent surfaces surrounded by an aqueous phase of variable composition. Although net adsorption is readily estimated 5 in such a mixture, attempts to interpret the mechanisms involved and the relative contributions of various mineral components have not been very succesful due to the difficulty of separating the reactions (Hingston, 1981).
For this reason, model or synthetic mineral systems have been widely used in adsorption studies and have provided much information about adsorption reactions.
Model adsorbents employed in previous investigations have included such common soil oxides as gibbsite, goethite, hematite, and lepidocrocite (Atkinson et al., 1967;
Hingston et al., 1967; Breeuwsma and Lyklema, 1973;
Hingston et al., 1974; Russell et al., 1974; Parfitt et al., 1977a,b; Parfitt and Russell, 1977; Harrison and
Berkheiser, 1982; Ainsworth et al., 1985; Tejedor-Tejedor and Anderson, 1986; Martin and Smart, 1987).
Unfortunately, most studies conducted to date have lacked a detailed charaterization of the absorbent and have rarely evaluated the effect of systematic variations of absorbent properties like crystallinity, surface area and ion substitution on adsorption. Such properties should influence tho capacity of the absorbent to react with anions in solution. 6
Goethite (a-FeOOH) is the most abundant and
widespread iron oxide in soils and normally occurs in
solid solution with Al. In fact, Al may substitute for
over 30 mole % of the structural Fe in natural specimens.
The present study was conducted to test the hypothesis
that Al-substitution in goethite should alter the
physicochemical properties of the oxide and therefore
influence anion adsorption. Sulfate was selected as the
adsorbate of interest because of the controversy
surrounding its mode of bonding to oxide surfaces.
Contradictory results have been reported suggesting the
formation of both outer sphere (Yates and Healy, 1975)
and inner sphere complexes (Harrison and Berkheiser, 1982;
Martin and Smart, 1987).
The specific objectives of the study were as follows:
(1) to synthesize a series of goethites with systematic variations in Al-substitution;
(2) to correlate Al-substitution with the capacity of the oxide to adsorb sulfate;
(3) to compare the adsorption capacities of the synthetic oxides with that of a soil clay rich in goethite; and
(4) to examine the mechanisms of sulfate adsorption at a molecular level. CHAPTER I
A REVIEW OF IRON OXIDES IN SOILS
Weathering of silicate minerals in soils results in the release of iron that is predominantly bound in bivalent form. Hydrolytic and oxidative decomposition contribute to the release through a reaction of extremely low reversibility:
Fe(II)-0-Si + HzO ■JiP.y -Fe(III)0H + -SiOH + e-
Most of the Fe(III) liberated is precipitated as oxides or oxyhydroxides. A very small portion may be incorporated in secondary layer silicate clay minerals or complexed by organic materials. Based on these facts, the relative degree of weathering (i.e., age) of soils have often been characterized by measuring the ratio of reductant-soluble iron to total iron (Alexander, 1974; Torrent et al.,
1980; Arduino et al., 1984; Alexander, 1985).
7 8
The iron oxides formed during weathering can be driven into solution again by the combined action of the indigenous microbial population and organic matter. An excess of water in the soil forces the microbial population to utilize Fe+3 as the final electron acceptor in order to accomplish the oxidative decomposition of organic matter. The reduced ferrous iron is soluble and is readily complexed with organic ligands. Eventually, both Fe+2 and Fe-organic complexes go through reoxidation and hydrolytic decomposition, respectively, to again form
Fe(III) oxides (Schwertmann, 1985). Thus, iron oxides may be both formed and transformed within the soil environment.
The structure of most iron oxides consists of close- packed oxygens in octahedral coordination around Fe+3 and/or Fe+2 ions. The ideal chemical composition of iron oxides is also rather simple, consisting only of Fe, 0, and OH. Nevertheless, they differ considerably in their crystal structure and chemical behavior (Schwertmann and
Taylor, 1977; Taylor, 1980; Taylor, 1987).
Common iron oxides in soils include goethite, hematite, lepidocrocite, ferrihydrite, feroxyhite, maghemite, and magnetite. Some crystallographic properties of these oxides are presented in Table 1.1. 9
Table 1.1. Some crystallographic data for the major Fe oxide minerals occurring in soils.
Crystal Unit Cell „ Oxide Formula System Dimensions (A)
Goethite a-FeOOH orthorhombic a=4.608, b=9.956 c=3.022
Hematite a-Fe203 hexagonal a=5.034, c=13.752
Lepidocrocite> ^-FeOOH orthorhombic a=3.88, b=12.54 c=3.07
Ferrihydrite 5Fe203.9H20 hexagonal a=5.08, c=9.40
Feroxyhite 6'-FeOOH hexagonal a=2.93, c=4.60
Magheroite # -Fe203 cubic or a=8.34 tetragonal a=8.338, c=25.01
Magnetite Fe304 cubic a=8.391
(Murad and Johnston, 1987) 1 10
1.1 Hematite (a-Fe203)
The basic structure of hematite is a hexagonal close-
packed array of oxygens in which 2/3 of the octahedral
sites in every layer are occupied by Fe3+. In x-ray
diffraction analysis, hematite produces a line doublet at 9 O 2.70 A and 2.52 A that are highly diagnostic. Infrared
spectra of hematite yield three prominent bands around
337, 468, and 546 cra'l (Fysh and Fredericks, 1983). The mineral is thermally stable.
Hematite is one of the most common soil iron oxides and is especially abundant in soils of dry and/or tropical areas. Hematite particles in soils are usually isodimensional, whereas synthetic samples show a more defined hexagonal shape (Torrent et al., 1980; Barron et al., 1984). Hematite is characterized by a reddish coloration, and even in very low concentrations has a profound effect on soil coloration. Childs et al. (1979) found that reddening occurs even when the hematite;total
Fe-oxide ratio is as low as 0.3. Recently, there is a growing interest in obtaining correlations between soil or sediment color and the content of hematite (Torrent et al., 1980; Torrent et al., 1983). Although the 11 pigmenting power of hematite may vary from one region to another, coloration can be a useful tool to predict hematite content in soils from the same geographical region.
1.2 Lepidocrocite (#-FeOOH)
Lepidocrocite is an orange-colored form of FeOOH which forms lath-shaped crystals ranging from 0.1 to 0.7
M.m in length and of variable thickness (Schwertmann,
1973). The structure is based on a cubic close-packed array of oxygen ions with Fe3+ in octahedral coordination.
X- ray diffraction analysis yields intense peaks at 6.27 O and 3.29 A. Lepidocrocite is less common in soils than goethite and hematite and is usually restricted to hydromorphic soils, particularly those high in clay, in which ferrous iron forms due to an anaerobic environment
(Schwertmann and Fitzpatrick, 1977; Schwertmann and
Taylor, 1977; Ross et al., 1979; Chen et al., 1980). In such soils, lepidocrocite usually occurs as bright orange mottles or as bands. Tarzi and Protz (1978) also found lepidocrocite associated with mica particles in two well drained Ontario soils developed on granite and 12 granite-gneiss and Ross and Wang (1982) likewise reported lepidocrocite in calcareous, well drained soils from
Quebec. Both groups of authors suggested that Fe(II)- containing silicates were the most likely source of Fe2 + for the formation of lepidocrocite.
Thermodynamically, lepidocrocite should not persist in soils but should transform to hematite or goethite.
Laboratory experiments have shown that lepidocrocite can be easily converted to goethite in the presence of alkali hydroxide or in ferrous sulfate systems (Oosterhout, 1967;
Schwertmann and Taylor, 1972a,b). The conversion to goethite starts with the dissolution of lepidocrocite as expressed by the following reaction:
FeOOH + OH- + H 20-*Fe(0 H)«-.
From the dissolution products nucleation and crystal growth of goethite takes place.
In spite of its lower stability relative to hematite and goethite lepidocrocite persists in soils and sediments for long periods of time. This fact suggests that some soil constituents are capable of retarding or even inhibiting its transformation. For example, 13
Schwertmann and Taylor (1972b) found that the presence of
Si in solution inhibited the conversion of lepidocrocite
to goethite by retarding the nucleation of the latter.
1.3 Ferrihydrite (5 Fe2C>3 • 9H 2 0 )
Ferrihydrite is the most recently reported iron oxide
in nature, and it is often the first to form during the
hydrolysis and precipitation of dissolved Fe from
solution. Due to its "amorphous" nature there is still controversy as to whether or not this yellow-brown to dark brown material can be considered as another Fe mineral. A diffraction pattern was first successfully obtained for
synthetic "ferric hydroxide" by Van der Giessen (1966) and
independently by Towe and Bradley (1967). They suggested a structure similar to hematite, based on hexagonal close- packing of oxygens, even though the two diagnostic o diffraction lines of hematite (3.67 and 2.69 A) were missing.
Chukhrov et al. (1973) proposed the name ferrihydrite for a similar naturally occurring compound with composition 5Fe203-9H20. This name has been accepted by the nomenclature commission of the International 14
Mineralogical Association. More recently, Russell (1979) has shown that structural OH groups are present and has suggested that the formula be amended to
Fe2 0 3 •2Fe00H-2.6H 2O in order to indicate their presence.
The suggested chemical formula of ferrihydrite indicates its similarity to hematite. In contrast to hematite, the
Fe positions are partly vacant, and the z-periodicity of the octahedral sheets is 4 rather than 6 . Both minerals o show a similar a-dimension (hematite:a = 5.08 A, o ferrihydrite:a = 5.04 A) but the c-dimension of ferrihydrite is two-thirds that of hematite (9.4 and 13.8
0 A, respectively).
Ferrihydrite occurs in nature in close association with goethite or lepidrocrocite (Schwertmann, 1985). It has frequently been identified in brown, gel-like precipitates from Fe-bearing waters (Schwertmann and
Fischer, 1973; Henmi et al., 1980; Carlson and
Schwertmann, 1981). Under these conditions, Fe is quickly oxidized and precipitated when the water is exposed to earth surface conditions. Ferrihydrite formed in this manner can be recognized easily by routine x-ray diffraction analysis since it is reasonably pure. To identify the oxide in soils where concentrations and 15 degree of purity are low, other techniques like dissolution kinetics, differential x-ray diffraction, and
Mossbauer spectroscopy are very useful (Schwertmann et al,
1982).
Ferrihydrite is thermodynamically unstable and with time converts to more stable Fe-oxides, usually goethite under temperate or cool, humid climates or hematite under dry and warmer conditions. Although it is considered a precursor of both goethite and hematite, it has not been positively identified together with the latter under natural conditions. A possible explanation for this might be that under conditions favorable for hematite formation, the conversion of ferrihydrite to hematite is much faster than the formation of ferrihydrite itself.
1.4 Feroxyhite (6 '-FeOOH)
This polymorph of FeOOH was described by Chukhrov et al. (1977) from deep sea nodules and gley soils in the
USSR. Electron diffraction studies showed that the structure of feroxyhite was related to the ferromagnetic
6 -FeOOH described by Bernal et al. (1959) and Dasgupta
(1961). Bernal et al. (1959) proposed a structural model 16
for 6 -FeOOH consisting of a two layer hexagonal packing of
0 and OH in which all the octahedra are half populated by atoms of Fe. They also reported a and c hexagonal cell o parameters of 2.94 and 4.49 A, respectively. Dasgupta
(1961) showed that the structure of 6 -FeOOH is intermediate between those of Fe(0 H )2 and Fe2 0 a.
Like ferrihydrite, the basic structure of feroxyhite
(6 '-FeOOH) is similar to hematite insofar as it consists of hexagonally close-packed oxygen atoms with Fe ions in the octahedral and tetrahedral interstices. Oxygen is partly replaced by OH and OH2 , and the degree of occupancy by Fe is less than the 2/3's typical of hematite. The periodicity of the octahedral sheets along the z direction o is 2 for feroxyhite (c = 4.60 A) as compared to 4 for
0 0 ferrihydrite (c = 9.40 A) and 6 for hematite (c = 13.77 A)
(Carlson and Schwertmann, 1980). X-ray powder diffractograms for feroxyhite and ferrihydrite are very similar, but the higher periodicity of the octahedral sheet of ferrihydrite results in additional diffraction o lines. A line at 1.97 A is characteristic of ferrihydrite and is the diagnostic feature to differentiate it from feroxyhite. Although lacking some spacings of ferrihydrite, feroxyhite seems to be more crystalline
(Carlson and Schwertmann, 1980). 17
Feroxyhite (8 '-FeOOH) is thought to form topotactically by the oxidation of Fe(0H)2. Under electron microscopy the particles show a platy form with irregular outline, but are rarely hexagonal. Electron diffraction patterns of 8 'FeOOH are typical of a poly crystal. In certain cases 8 '-FeOOH shows magnetic properties like 8 -FeOOH but to a lesser extent. Chukhrov et al. (1977) suggested that the cause of such magnetism might be the result of an admixture of maghemite.
1.5 Maghemite (8-Fe20a)
Maghemite is the cubic, ferromagnetic form of Fe2 0 a and it occurs most frequently in subtropical and tropical soils (Taylor and Schwertmann, 1974; Coventry et al.,
1983). Abreu and Robert (1985) also found maghemite in the B horizons of two Alfisols and one Ultisol from
Portugal under a climate that can be considered temperate.
The chemical composition of maghemite is ideally the same as hematite but usually includes some Fe2 +. Maghemite
(unlike hematite) is strongly magnetic and has a crystal structure similar to magnetite. It is frequently found in association with hematite in soils. 18
Maghemite can be formed by the low-temperature oxidation of magnetite via a topotactic process, so called because the oxygen framework is left basically unchanged.
The Fe2+ ions in magnetite diffuse to the surface of a grain and oxidize to Fe3 + leaving structural vacancies.
If nothing impedes this diffusion process, one raicrocrystal of magnetite can be completely transformed to maghemite without an additional new crystal of maghemite being formed (Murray, 1979).
Although a realistic possibility for maghemite formation, the oxidation of magnetite does not explain the occurrence of maghemite in soils whose parent materials are low in magnetite. Synthesis studies suggest the formation of an Fe2+ Fe3+ hydroxy compound as a necessary precursor of maghemite. The simultaneous rates of oxidation and dehydration determine whether lepidocrocite or maghemite is formed (Taylor and Schwertmann, 1974).
The frequent occurrence of maghemite in tropical areas supports the theory that a dehydration process is involved. Also, the presence of silicates high in Fe2+ appears to favor its formation (Schwertmann and Taylor,
1977). 19
1.6 Magnetite (FesCM)
Magnetite (Fe304) is a member of the spinel group and is very common in nature. The structure of magnetite is that of an inverse spinel; that is, one third of the iron as Fe(III) occupies all 8 tetrahedral sites, one third as
Fe(III) occupies half of the 16 octahedral sites, and one- third as Fe(II) occupies the other half. The x-ray diffraction pattern of magnetite shows two strong lines at o 2.53 and 2.97 A. Upon heating, magnetite converts to hematite.
1.7 Goethite (a-FeOOH)
Goethite (a-FeOOH) is the most important and widespread iron oxide in soils. It occurs in almost every soil type and climatic region and is responsible for the yellowish to yellowish brown color of many soils. In many dry and/or tropical environments it is found in association with hematite (cc-Fe2Q 3 ). The relative abundance of each oxide is closely related to the climatic conditions of the region with wetter and cooler climates favoring the presence of goethite. Hematite in sediments 20 and paleosols exposed to cooler and wetter climates frequently transforms to goethite via solution as
\. evidenced by a yellow color penetration from the top into the redder subsoils or substrata (Schwertmann, 1971).
Apparently, hematite is dissolved by reduction and/or organic complex formation, and subsequent oxidation and precipitation under the new environmental conditions causes goethite to form. Organic matter plays a very active role in these processes. This is demonstrated by the occurrence of yellow zones adjacent to plants roots in red soils where, under the influence of organic compounds supplied from the roots, hematite is transformed to goethite.
The stability of goethite and hematite is represented by the reaction: goethite = hematite + water. Berner
(1969) found that fine grained goethite is unstable relative to hematite under most geological conditions.
Langmuir (1971) reported similar results but, in addition to a particle size effect, he included as well temperature and pressure effects. Langmuir concluded that goethite cubes less than about 0.01 um on an edge are thermodynamically less stable than hematite under geological conditions. He suggested that goethite may 21 precipitate directly at low temperatures and pressures, but that hematite usually forms by long term aging or by slow dehydration of amorphous ferric hydroxide. Yapp
(1983) disagreed in part with Langmuir and suggested that
Langmuir's data may be correct only for pure goethite. He found that the solid solution of A100H in goethite can increase the thermodynamic stability of goethite with respect to hematite, allowing even very fine-grained Al- bearing goethite particles to be thermodynamically stable.
Additional information about factors affecting goethite/hematite formation will be discussed in a subsequent chapter dealing with the properties and synthesis of goethite. CHAPTER II
PROPERTIES AND SYNTHESIS OF GOETHITE
2.1 Structure of Goethite
The structure of goethite consists of hexagonally
close-packed planes of 0 atoms in octahedral coordination with Fe. Double chains of linked [Fe(0,0H)e] octahedra running parallel to the c axis are linked by sharing opposite edges (Fig. 2.1). One of the chains is displaced by c/2 with respect to the other so that an octahedra from one chain shares an edge with two octahedron from the other chain. The double chains of Fe-O-OH octahedra are o also linked by O-H-O bonds approximately 2.65 A in length.
There are two kinds of 0 sites in goethite: Oi sites where 0 atoms are shared between octahedra in two different chains, and Oii sites where 0 is shared by octahedra in the same chain. The H atom does not lie on a direct line joining the Oi-Oii atoms, but occurs on a line
22 Fig. 2.1. Pictorial representation of goethite structure, with the hydrogen bonds shown as tubes CEwing. 1935).
to to 24 making an angle of about 11.6 degrees with the Oi-Oii vector. The H atoms are more closely associated with the
On atoms so that the On sites can be considered the hydroxyl sites. In addition to these structural hydroxyl groups, the presence of three types of surface hydroxyls has been suggested (Parfitt, 1978). These hydroxyls have been designated as A (coordinated to one Fe atom), B
(coordinated to three Fe atoms), and C (coordinated to two
Fe atoms) (Fig. 2.2). Their distribution is not even over the different crystal faces. All three types occur on the
(1 0 0 ) face, singly and doubly coordinated on the (0 1 0 ), and only singly coordinated on the (001) face (Cornell et al., 1974).
The presence of goethite in soils can be recognized « by x-ray diffraction lines at 4.18, 2.69, and 2.44 A. The presence of quartz and hematite may impede identification o since the 4.18 and the 2.69 A lines of goethite may be o difficult to resolve from the 4.26 A line of quartz and 0 the 2.70 A line of hematite, respectively. However, in 0 soil materials with high amounts of goethite the 2.44 A line is also visible (Schwertmann and Taylor, 1977). 25
Fig. 2.2. Structure of the (100) face of goethite (^-FeOOH) (Parfitt, 1978). 26
2.2 Composition of Goethite
Goethite normally forms via solution so that ions
other than Fe may also be incorporated into the
structure. Aluminum is the most common impurity in soil
goethites and can be present in higher concentrations than
any other foreign ion. The extent of substitution in
nature commonly lies between 10 to 20 mole %; however,
values up to 36 mole percent have been reported (Table
2.1). Al-substituted goethites have also been obtained by
laboratory synthesis and substitutions up to 33 mole %
have been achieved (Thiel, 1963; Golden, 1978).
The Al®+ ion is slightly smaller than the Fe®+ ion » s (0.53 A vs. 0.65 A); thus, when Al for Fe substitution
takes place in the goethite structure, the average size of
the unit cell decreases. A progressive increase in Al
substitution will be reflected in a progressive shift of goethite x-ray diffraction lines to higher degrees 29,
i.e., in a shift to smaller d values. Thiel (1963) found that d(lll) and the unit cell dimensions decreased linearly with Al substitution. Similar results were obtained by Jonas and Solymar (1970), Golden (1978), and
Fey and Dixon (1981). 27
Table 2.1. Aluminum substitution in goethites.
Soil Order Location Mole % Al Reference
Alfisols South Africa 7-12 Fitzpatrick and Schwertmann (1982)
Ultisols Southeast USA 11-32 Bigham et al., (1978)
Oxisols Brazil 24-36 Curi and Franzmeir (1984) 28
In contrast, Schulze (1984) showed that while there is a general decrease in the unit cell dimensions, this decrease is not uniform. His results indicate that only the c-dimension decreases linearly with increasing Al substitution, whereas the a-dimension is quite variable for samples with the same degree of Al substitution.
The existence of the isostructural Mn mineral groutite (cc-MnOOH) suggests that Mn3 + may also replace Fe in the goethite structure; however, Mn-substituted goethites have not yet been found in soils. Synthesis experiments have demonstrated that a limited Mn-Fe solid solution is possible. Stiers and Schwertmann (1985), and
Cornell and Giovanoli (1987) were able to synthesize goethites containing up to 15 mole % Mn and observed changes in unit cell dimensions like those obtained with
Al-substituted goethites.
Impurities of Si02 are also found in goethite.
Silica not only is adsorbed on the surface of the oxide, but seems to be incorporated into the structure. Some evidence supporting the incorporation of silica into the estructure of goethite was presented by Schwertmann and
Taylor (1972a) who were unable to remove freshly precipitated silica with KOH without dissolution of 29 co-forraed goethite. A factor that may enhance incorporation or occlusion of silica into the structure of goethite is the presence of raultidomainic crystals. Smith and Eggleton (1983) found that some natural goethites containing silica showed a microstructure with clear boundaries or interdomainic areas.
Other ions that can be found in small concentrations in goethite are Mg2+, Cr3+, Ni2 + and Co2+ (Table 2.2).
Carbon can also occur as a minor impurity in goethite and appears to exist in two principal forms: CO2 trapped in the mineral structure and organic matter (Yapp and Poths,
1986).
2.3 Goethite Morphology
Synthetic goethite is usually acicular with preferential growth in the direction of the z axis.
Atkinson et al. (1968) found that two types of crystals may form depending mainly on the hydrolysis rate. Small acicular crystals were nucleated by very small goethite crystals in the parent solution at slow hydrolysis rate, whereas twinned crystals occurred when hydrolysis was rapid. Twin crystal growth arose from a small nucleus of 30
Table 2.2. Chemical composition of some soil goethites.
Concentration %
Soil sample number *
1 2 3 4 5
Fe203 72.7 74.8 69.9 71.7 59.2
AI2O 3 4.17 2.60 3.71 3.61 11.0
Si02 3.27 2.40 1.99 2.69 4.88
MgO 0.63 0.40 0.81 0.43 1.42
Cr203 2.93 2.92 3.51 3.60 3.31
MnO 0.65 0.33 0.85 1.53 1.15
NiO 1.27 1.01 1.92 1.49 1.86
CoO 0.11 N.D. N.D. 0.03 0.11
* Samples 1-3 from New Caledonia, samples 4 and 5 from Philippines. (Taylor, 1987) 31
hematite. Formation of twinned crystals was also reported
by Cornell and Giovanoli (1985), but they found that a
nucleus of hematite was not necessary for twin formation.
Twins without hematite were nucleated from ferrihydrite
and their formation was enhanced by increasing suspension
concentration and ionic strength.
The morphology of goethite in soils is variable in
contrast to the predominant needle-shape of synthetic
samples. Acicular goethite, while found in soils, is not very common and occurs primarily in voids where formation takes place isolated to a certain degree from the
immediate soil environment. Solid solution and
interferences by foreign ions may restrict the development of needle-shaped goethite in soils. Nakai and Yoshinaga
(1980) identified fibrous goethite in some Japanese and
Scottish soils, and attributed this morphology to structural incorporation of Al, Si and other ions. A similar material identified as Al-substituted goethite was found in a podzol from Tasmania by Fordham et al. (1984).
At high magnification the fibers were seen to consist of very small (5 nm) particles aligned along the length of each fiber. 32
Cornell et al. (1974) reported that only three faces of synthetic goethite, the (1 0 0 ), (0 1 0 ), and (0 0 1 ), develop to a significant extent (Fig. 2.3). The (100) face comprises 50 to 60 % of the total surface area and the (010) face about 31 %. A decrease in crystal size will enhance the contribution of the (0 1 0 ) and (0 0 1 ) faces to the total surface area.
More recently, Schwertmann (1984), and Cornell and
Giovanoli (1986) have shown that the (110) face tends to dominate on some goethite crystals. Cross sections taken almost at right angles to the z axis showed crystals surrounded by (1 1 0 ) faces with very limited development of the (1 0 0 ) and (0 1 0 ) surfaces.
A morphological feature of goethite crystals that has received much attention recently is the presence of multiple domains or intergrowths (Cornell and Mann, 1983;
Schwertmann, 1984; Schwertmann et al., 1985; Cornell and
Giovanoli, 1985; Mann et al., 1985; Cornell and Giovanoli,
1986; Cornell et al., 1987; Schulze and Schwertmann,
1987). Multidomainic crystals are common in synthetic goethites and have been observed also in natural specimens (Smith and Eggleton, 1983). The domains or intergrowths run parallel to the z axis of the crystal Fig. 2.3. Schematic representation of acicutar goethite showing the three major crystallographic faces CCornel I et al.. 19743. w LO 34 and are usually of unequal lengths. They seem to nucleate somewhere in the center of the crystal and then grow along the z axis (Fig. 2.4). Electron diffraction analysis of these composite crystals reveals in many cases that the material is monocrystalline with no long raisorientation at the domain boundaries (Cornell et al., 1983; Mann et al.,
1985).
The presence of domains accelerates the desintegration of the crystals in acid media. After acid attack, parallel fissures following the domain boundaries develop on the (100) face (Cornell et al., 1974;
Schwertmann, 1984). The domain boundaries seem to be areas of imperfect crystallization and consequently are more susceptible to weathering.
The formation of multidomainic goethites appears to be the result of rapid growth (Cornell and Giovanoli,
1986). Crystal growth of goethite is thought to occur via addition of monovalent Fe complexes (Fe(0 H) 4 _ or Fe(0 H)2+) to the nucleus. At high pH the rate of formation is so rapid that growth units do not have time to order completely and a stacked structure with irregular ends develops. The presence of Al in solution retards the rate of crystal growth and reduces the number of domains per Fig. 2.4. Schematic representation of a muItidomainic crystal of goethite showing domain boundaries. to Ln 36 crystal (Schwertmann, 1985; Mann et al., 1985). The retarding effect of Al can be attributed to the competition for growth sites on the crystal between the divalent Al(0H)s2- ion, which is the predominant Al species in solution above pH 11, and monovalent Fe(0 H)4 -.
Monovalent Fe(0 H) 4 “ is a more suitable unit for crystal growth. Since these ions must be neutralized at the crystal surface before being built into the crystal, this neutralization is probably easier for monovalent than for divalent ions.
High levels of NaNOa in the system also retard crystal growth but at the same time enhance intergrowth formation (Cornell and Giovanoli, 1986). The authors suggest that the formation of ion pairs between FeO~ and
Na+ may interfere with the surface mobility of the adsorbed growth units. Blocking of the "appropriate" growth pathways results in intergrowths. Foreign ion interference is probably responsible for the formation of intergrowths in natural goethites. 37
2.4 Thermal Properties of Goethite
Goethite exhibits an endothermic dehydroxylation peak between 280 and 400 °C. The dehydroxylation temperature increases with Al substitution and improved crystallinity and decreases with increasing surface area (Jonas and
Solymar, 1970; Fey and Dixon, 1981; Schulze and
Schwertmann, 1984; Schwertmann, 1984). Quite often a double dehydroxylation peak has been observed in both natural and synthetic specimens (Jonas and Solymar, 1970;
Creer et al., 1971; Derie et al., 1976; Murad, 1979; Fey and Dixon, 1981; Schwertmann, 1984). Derie et al. (1976) proposed that in well crystallized goethites of low surface area a hematite shell is formed at the beginning of the dehydration process. This hematite shell surrounds the remaining goethite and retards further dehydration.
On the other hand, Schwertmann (1984) found that a double dehydroxylation peak was most common in samples with surface areas in the range of 40 to 60 m 2/g, whereas below and above this range a single peak occurred.
Murad (1979) produced a double dehydroxylation peak by grinding a natural sample of goethite. He suggested that the release of loosely bound water was facilitated by 38 grinding and that vaporization of this water was the cause of the low temperature dehydroxylation peak. Jonas and
Solymar (1970) attributed the double dehydroxylation peak to gibbsite impurities in the samples whereas Fey and
Dixon (1981) attributed it to the higher dehydroxylation temperature of A1-0H compared to Fe-OH groups in Al- substituted goethites. These explanations cannot account for the occurrence of a double dehydroxylation peak in pure (unsubstituted) goethites. In fact, Schulze and
Schwertmann (1984) observed a more pronounced doublet in unsubstituted goethite, and noted that it disappeared as
Al substitution increased.
Schwertmann (1984) found that for goethites of high crystallinity a significant contraction in the (001) direction of the unit cell occurs during heating forming what he calls a "high temperature goethite". This contraction raises the dehydroxylation temperature of residual goethite and thereby produces the high temperature peak. The double peak is not evident in goethite of low crystallinity because it transforms to hematite at a temperature too low to produce a change in the unit cell size. 39
2.5 Synthesis and Formation of Goethite
2.5.1 Synthesis of Goethites
Goethite (a-FeOOH) can be easily synthesized under laboratory conditions by aging freshly precipitated hydroxy gels. The precipitation process involves hydrolysis of the Fe(III) hexaquo complex ion, which is promoted by the addition of a base (NaOH or KOH) or other reagents to remove H+ ions (Towe and Bradley, 1967;
Atkinson et al., 1968; Lewis and Schwertmann, 1979;
Schulze, 1984). Aluminum-substituted goethites can be prepared by a similar process. These goethites are usually synthesized by co-precipitating a solution containing an Fe salt and Al salt, and aging the hydroxide gel thus formed (Jonas and Solymir, 1970; Golden, 1978;
Taylor and Schwertmann, 1978, Lewis and Schwertmann,
1979a,b; Fey and Dixon, 1981; Schulze, 1984).
Ferrous salts can also be used to synthesize goethite but require oxidation of Fe2 + by bubbling Oz or C02 through them. The system is usually kept above pH 7
(Taylor and Schwertmann, 1978; Fey and Dixon, 1981;
Schulze, 1984). In general, goethites formed from Fe2 + 40 systems are not as well crystallized as goethites from
Fe3+ systems which show sharper x-ray diffraction lines and smaller surface areas.
Hematite (a-Fe203) or a mixture of goethite/hematite can also form from the aging of these hydroxy gels
(Gastuche et al., 1964; Lahan, 1976; Fischer and
Schwertmann, 1975; Lewis and Schwertmann, 1979a,b,c). The competitive formation of hematite and goethite under both laboratory and natural conditions is one of the most interesting aspects of the iron oxide system. This factor must be taken into consideration when attempting to synthesize pure samples of either goethite or hematite.
2.5.2 Factors Affecting the Formation of Goethite
The formation of goethite and its properties can be affected by various factors such as foreign ions in the growth medium, organic matter, temperature, moisture and pH. The small arrows in Fig. 2.5 indicate that increasing or decreasing expression of a given factor favors hematite formation over goethite. Fischer and
Schwertmann (1975) have theorized that goethite and hematite form from ferrihydrite through two competitive 41
Factors Process
Fe Q ions /rale of Fe release' ✓organic matter ✓pH {pH 3 -8 )
Ferrihydrife
o € ✓soil lemperolure ✓soil moisture 5 ?
Hematite Goethite
Fig. 2.5. Schematic representation of the competitive process of goethite and hematite formation (Schwertmann. 1085). 42 processes. Whereas goethite develops through dissolution of ferrihydrite and crystallization from solution, hematite formation is preceded by an aggregation of the ferrihydrite and nucleation within these aggregates.
Goethite and ferrihydrite both form from Fe(III) ions depending mainly on whether or not the solubility product of the mineral is exceeded. Ferrihydrite has a higher solubility product (10-38) than goethite (10~42) and will require a higher concentration of Fe(III) in solution for its formation. Since ferrihydrite appears to be a necessary precursor for hematite formation, any factor inhibiting its stability will favor the development of goethite over hematite.
Aluminum is the foreign ion that most commonly influences the formation of goethite in soils. The incorporation of Al in the crystal structure results in shorter but thicker crystals with a more homogeneous appearance and mostly monodomainic character (Schulze and
Schwertmann, 1987). These changes in morphology also result in changes in the surface area of goethite, an effect that will vary depending on the amount of Al incorporated. Golden (1978) and Ainsworth (1985) reported an increase in surface area with Al substitution, but 43
Schulze and Schwertmann (1987) found a decrease in surface area with Al substitution up to about 12 mole %, after which an increase in surface area took place.
The presence of Al in solution also appears to favor the development of hematite over goethite (Lewis and
Schwertmann, 1979b). Aluminum probably retards the nucleation of goethite, thereby indirectly favoring hematite formation. On the other hand, Taylor and
Schwertmann (1978) found that Al inhibited the development of lepidocrocite and maghemite in Fe2+ systems and thereby favored goethite formation.
The presence of silicate in solution retards the transformation of ferrihydrite to either goethite or hematite but favors the latter over the former. It has been suggested that silicate species stabilize ferrihydrite by adsorbing on the particles surfaces and linking them in an immobile network (Cornell et al.,
1987). These authors also found that silicate modifies the morphology of goethite yielding pseudohexagonal crystals rather than the common acicular type.
Manganese ions may also influence the transformation of ferrihydrite to goethite. Cornell and Giovanoli (1987) showed that Mn(II) suppressed the formation of hematite 44 and promoted twinning of goethite crystals. Incorporation of Mn into the goethite structure caused the oxide to darken with the color changing from olive to gray as Mn content increased.
Organic acids retard or inhibit the crystallization of other Fe oxides from ferrihydrite in soils.
Schwertmann (1969) showed that the behavior of organic acids falls into two categories: (1) those that completely inhibit crystallization at pH 9-10 (hydroxy carboxylic acids), and (2) those that only retard crystallization and also alter the ratio of goethite/hematite formed
(carboxylic acids).
The effect of organic acids on ferrihydrite crystallization seems to be through adsorption mechanisms.
Fischer and Schwertmann (1975) proposed three ways in which organic molecules may influence the crystallization/condensation process: (1) by adsorption on ferrihydrite, thereby preventing JLts dissolution to form goethite or condensation to form hematite; (2) by association of the organic acid ligand with Fe in solution
(complexation) and/or at the surface of the goethite nuclei and crystals, thus inhibiting nucleation and/or crystal growth; and (3) by acting as a template and thereby favoring the formation of hematite. 45
The effect of organic acids that weakly retard crystallization (like lactic acid) can be overcome just by adding seeds of goethite to the system containing ferrihydrite. In contrast, in the presence of an acid that strongly inhibits crystallization (like tartaric acid), addition of goethite or hematite does not encourage further formation of either oxide (Cornell and
Schwertmann, 1979).
Hematite formation can be prevented at low concentrations of organic acid if the acid is capable of linking ferrihydrite particles to form a relative immobile network in which the area of direct contact between the particles is reduced and their internal nucleation to hematite becomes impossible. A piece of evidence supporting this was obtained by Schwertmann and Fischer
(1973) who found an unexpected low surface area for natural and synthetic ferrihydrites containing organic compounds.
The capacity of an organic acid to stabilize ferrihydrite will depend on whether it can adsorb on the oxide surface in the pH range considered and on the nature of the ligands involved. Linking between an acid and ferrihydrite particles is strengthened by COOH/OH pairs. 46
Therefore, hydroxy-di (tri) carboxylic acids such as citric and tartaric are particularly strong inhibitors, whereas dicarboxylic acids are only effective at high concentrations.
The transformation of ferrihydrite to goethite and hematite at high pH (9-13) is also retarded by the presence of simple sugars (Cornell, 1985). Like organic acids, the retarding effect of a sugar will depend on its capacity to adsorb on ferrihydrite. Maltose and glucose adsorb strongly and inhibit the transformation by preventing aggregation or dissolution of ferrihydrite.
Sucrose, which adsorbs to a much lesser extent, seems to hinder the nucleation and growth of goethite in solution.
The effect of temperature and moisture on the formation of goethite and hematite is evident from the zonal distribution of soils with predominance of either oxide. Cool humid areas are usually free of hematitic soils whereas in warmer areas hematite tends to be much more abundant (Kampf and Schwertmann, 1983; Schwertmann et al., 1982). The effect of temperature has also been shown in synthesis experiments where an increase in temperature favored the formation of hematite (Lewis and Schwertmann,
1979c). Pena and Torrent (1984) studied the 47
goethite/hematite contents of well and poorly drained
Alfisols from Spain. Goethite was present in all soils
but hematite was only present in the well drained soils
and in A horizons of the poorly drained soils.
The effect of relative humidity on the transformation
of ferrihydrite to goethite and hematite was studied by
Torrent et al (1982). A decrease in relative humidity
decreased the rate of transformation of ferrihydrite to
either oxide, but an increase in hematite formation
relative to goethite was observed.
The importance of pH in goethite and hematite
formation is very clear from synthesis experiments.
Schwertmann and Murad (1983) found that a maximum amount
of hematite was formed from ferrihydrite in the pH range
of 7 to 8. Maximum goethite formation occurred at pH 4
and 12. Their results confirmed the existence of two
competitive processes of formation: goethite forming via
solution and hematite forming by internal dehydration and
rearrangement of ferrihydrite aggregates. Goethite crystal growth appears to occur by addition of monovalent
Fe(III) ions to nuclei. The Fe(0H)2+ ion predominates at pHs around 4 whereas, the Fe(0H)4~ ion is abundant at pHs above 11. At pHs near neutrality both species occur at 48
very low concentrations; hence, hematite formation is favored. The importance of pH as a determining factor in the formation of goethite and hematite under natural conditions is not well documented, but a study conducted by Kampf and Schwertmann (1983) showed that goethite/hematite ratios in some Brazilian soils increased with decreasing soil pH within the range of 4.0-5.6.
2.6 Materials and Methods
2.6.1 Sample Preparation
Al-substituted Goethites
A series of aluminum-substituted goethites was prepared following the method of Schulze (1984). This method involved the precipitation of a mixture of ferric nitrate (Fe(N03)3) and aluminum nitrate (A1(N03)3) using
KOH. All solutions containing KOH were handled in polypropylene bottles and were not allowed to come in contact with glass to avoid silica contamination. The solutions were measured to the accuracy attainable with polypropylene graduated cylinders. A solution of Q.5M A1(N03)3 (500 ml) was slowly
poured into 300 ral of 5M KOH. Some local precipitation
(presumably Al-hydroxide or oxyhydroxide) occurred after
each small addition of the A1(N03)3 solution, but the precipitate quickly dissolved with stirring of the
solution. The required amount of this solution was poured
into a 2 1 polypropylene bottle, followed by the required amounts of 5M KOH and M Fe(N03)3 solutions in that order
(Table 2.3). The coding of the samples (S-0---S-200) indicates the volume of aluminate solution added. After the addition of the Fe(N03)3 solution, the contents were thoroughly mixed and quickly diluted to 2 1 with deionized water.
The bottles were capped and shaken, then placed in an oven already set at 70 *C and kept there for 14 days.
During that period of time the bottles were shaken once a day. At the end of this aging period, most of the supernatant solution was syphoned off with a plastic tube and the precipitate was washed twice with M KOH to remove any Al not incorporated into the structure of the oxide.
The pH of the precipitate was adjusted to 7.5 with M HC1, then washed six times with deionized water to remove excess salts. The oxide was then dried at 70 °C, gently crushed in an agate mortar, and stored for subsequent use. 50
Table 2.3. Solution volumes for the preparation of the aluminum-goethite series.
Sample Aluminate 5M KOH 1M Fe(N0s)3
ml
S-0 0 180 100 S-10 10 179 100 S-30 30 176 100 S-80 80 170 100 S-120 120 165 100 S-150A 150 161 100 S-150B 150 161 100 S-200 200 155 100 51
2.6.1.2 Hvdrothermallv-treated Goethites
Subsamples of pure goethite (no Al), prepared as
described in 2.6.1.1, were hydrothermally treated at 50,
100, 150, and 200 *C using a pressure bomb. The bomb used was the model 4748 Acid Digestion Bomb manufactured by
Parr Instrument Co. It consists of a 125 ml teflon holder sealed in a stainless steel body (Fig. 2.6). This bomb can withstand temperatures up to 250 °C and pressures of
1900 psig. Approximately 4 g of sample with 70 ml of deionized water were placed in the bomb which was then kept in a temperature controlled furnace for seven days at the desired temperature. After aging the samples were collected by centrifugation, dried at' 70 °C, gently crushed in an agate mortar and stored for subsequent use.
This series of goethites was designated as S-0C (starting material), S-0C (50 'C), S-0C (100 °C), S-0C (150 °C), and
S-0C (200 ‘C).
2.6.1.3 . .Nipe^Soil -Clay.
A clay extracted from the Nipe soil (clayey, oxidic, isohyperthermic, Typic Acrorthox) of Puerto Rico was also used in the study. This soil contains considerable amounts of aluminum and iron oxides and it has shown 52
Compression Screw
Pressure Plote | to e HA Cm Scr. M l AC Screw Cap •MAC
Spring _U )30IA C Rupture Disc "-Compression 911AC Ring I07AC Corrosion Disc-^1 310AC
"-Cover— I! MSAC Liner A s s y - A305AC ^-Cup- 1 i? 104AC Bomb Bod M3AC
W a
Bottom Plate MIAC Fig. 2.6. 4748 flcid Digestion Bomb Parr Instrument Co. 53
strong capacity for phosphate and sulfate retention (Fox,
1982). Soil material from the surface horizon of the Nipe soil was sodium saturated and fractionated into sand (2 mm-50jj.m), silt (50um-2M.ra) and clay (<2p.ra) fractions using conventional centrifugation and gravitational techniques
(Jackson, 197b). The sand fraction was separated from the silt and clay fractions by sieving. Separation of silt and clay fractions was achieved by repetitive stirring- syphoning cycles using an automatic clay separator
(Rutledge et al., 1967). The syphoned clay was flocculated with IN MgCl2 solution to facilitate the decanting of the supernatant liquid. The clay was then repetitively washed with deionized H 2O to remove excess salt, frozen, freeze dried, and stored for subsequent use.
2.6.2 Characterization of the Synthetic Oxides and Nipe Soil Clay
2.6.2,1 X-rav Powder Diffraction
The oxides and clay samples were gently ground in an agate mortar to break up large aggregates and then back filled into an aluminum sample holder (15x20 mm sample area). The material was gently pressed against a filter 54
paper to minimize preferred orientation. The random powder mounts were irradiated with Cu Ka radiation (35 kV,
20 raA) on a Philips PW 1316/90 wide-range goniometer fitted with a theta compensating slit, a 0.2 mm receiving slit, and a diffracted beam graphite monochromator. The samples were step scanned from 10* to 65*28 at 0.1*28 increments using a counting time of 40 s/increment. The digital data were stored on floppy discs using an IBM personal computer, re-formatted for compatibility with the
Lotus 1-2-3 graphics package, and plotted using an IBM XY
749 plotter.
2...6-,.2_„2 Calculation of Unit Cell Dimensions and Mean Crystallite Dimensions (MCD1
The average unit cell dimensions of goethite
(orthorhombic structure) were calculated using the d values of the (020), (110), (120), (130), (021), (040),
(111), (121), (140), (131), (041), (221), (240), (060), and (061) lines. The d value for a given (hkl) line is related to the unit cell dimensions a,b,c by the equation:
d(hki) = [(h/a)2 + (k/b)2 + (l/c)2]-*. 55
The Mean Crystallite Dimensions (MCD) of the goethites were also calculated from the x-ray powder data using the Scherrer formula:
MCD = 0.9 X/Q cos 9
where, A= 1.5418 A (wavelength for Cu Kcc);
(3 = pure breadth of a powder reflection in radians
(after correction for instrumental broadening);
9 = angle in radians corresponding to the peak
maximum.
The experimental widths of the goethite lines were corrected for instrumental broadening using nearby lines of a well-crystallized synthetic hematite. This hematite was obtained by hydrothermally treating a pure sample of goethite in the pressure bomb (see section 2.6.1.2) for seven days at a temperature of 250 “C. Expanded plots of the step-scanned peaks were made using an Energraphics 2.1 graph program (Enertronics Research Inc.). These expanded plots facilitated the measurement of peak width at half height (Fig. 2.7). 2000
1500 TJ
O 1000
U 500
20.0 20.5 21.0 21.5 22.0 22.5 Degrees 2 Theto Fig. 2.7. Expanded plot of M O peak of synthetic goethite used for MCD determination. 57
The MCDa (thickness in the a direction) was estimated from the average MCD of the (110), (120), (130), and (140) lines. The (110) and (120) lines of goethite were corrected for instrumental broadening using the (012) line of hematite, and the (130) and (140) using the (110) and
(113) lines, respectively. Since all these lines result from planes parallel to c and forming an angle a with the
(100) plane (Fig. 2.8), the thickness along the a axis of goethite is given by MCD(ikO) times cos a. The MCDb
(thickness in the b direction) was estimated from the MCD of the (020) line. This line was corrected for instrumental broadening using the (012) line of hematite.
2.6.2.3 Transmission Electron Microscopy
Transmission electron micrographs were obtained using a Philips EM-300 transmission electron microscope operated at 60 kv. Samples were prepared by dispersing 1 rag of precipitate or clay in 50 ml of distilled water with an ultrasonic probe. A drop of each sample suspension was then air-dried on a 200-mesh copper grid coated with polyvinyl formvar. The samples were coated with platinum metal in a vacuum evaporator before examination with the transmission electron microscope. The approximate shadowing angle was 26.5“. (140) /1130)
/ / 1120] 11101 >Y
--I100] "T (010) Fig. 2. 8 . ClkO) planes used to calculate MCDa. MCDa = M C D |1K0| cos 1,2 ,3 ,4 . 59
2.6.2.4 Differential Scanning Calorimetry (DSC)
Approximately 10 rag of each goethite was heated from ambient temperature to 600 °C at 10"/rain under a stream of nitrogen using a Dupont 990 Thermal Analyzer and a DSC cell. An empty aluminum pan was used as a reference to measure calorimetric changes in the samples.
2..6...2.J) Surlage.-Area
The surface areas of the oxides and the Nipe clay were determined using the BET equation (Brunauer et al.,
1938). A Quantachrome Quantasorb Multipoint Instrument was used to measure N2 adsorption on samples cooled to liquid N 2 temperature. Samples of about 70-80 rag were outgassed under N 2 gas flow for an hour at 110 °C.
Surface area determinations were then made using N 2/He mixtures of 10.1, 20.5, and 30.5% N 2 (triple point) or
30.5% N 2 only (single point). Gas adsorption was converted to surface area assuming that each molecule of
N 2 covers 16 x 10~2 0 m2 of sample surface.
2.6.2.6 Determination of Total Iron. Total Aluminum and Oxalate Soluble Iron
Total Fe and Al contents of the oxides were determined by dissolving duplicate 10 mg samples in 3 ml of 6M HC1. The samples plus acid were placed in plastic 60
centrifuge tubes and heated for 15 rain at 110 “C to speed the dissolution process. After cooling, the solutions were transferred to 25-ml volumetric flasks and diluted to volume with distilled water. For Fe determinations, 5 ml of the samples were diluted to 100 ml. Aluminum was determined directly from the 25 ml dilution. Both elements were determined by atomic absorption spectroscopy using a Varian AA-6 spectrophotometer.
Oxalate soluble Fe was determined by shaking subsamples of the oxide (100 rag) with 10 ml of ammonium oxalate at pH 3 for 2 hr in the dark (Schwertmann, 1964).
The samples were then centrifuged and 1/25 dilutions of the supernatants were prepared for Fe determination by atomic absorption.
2.6.2.7 Reductant-soluble Iron Content,_of Niue Clay
The reductant-soluble Fe content of the Nipe clay was determined by the sodium dithionite citrate bicarbonate
(DCB) method (Mehra and Jackson, 1960). Forty ml of citrate buffer (pH 7.3) were added to duplicate samples of
1 g clay, and the samples were placed in a water bath at
75 to 80 °C. Two g of sodium dithionite were added and the samples were stirred occasionally for 5 min. 61
Following centrifugation, each supernatant was decanted
into a 250-ml volumetric flask. This sequence of steps
was repeated until the reddish coloration of the sample
disappeared, an indication of complete removal of Fe20s.
All extracts were combined into the 250-ml volumetric
flask and brought to volume. A 1/100 dilution of the
original extract was then analyzed for Fe using a Varian
AA-6 atomic absorption spectrophotometer.
2.7 Results and Discussion
2.7.1 Synthesis of Al-substituted Goethites
A total of eight different mixtures of Fe(N03)3 and
A1(N03)3 solutions were employed in the preparation of the
Al-substituted goethites (see Table 2.3). Enough 5 M KOH was added to keep the KOH concentration at approximately
0.5 M with a final volume of 2000 ml. The Fe concentration was kept constant at 0.05 M whereas the Al concentration was varied from 0 to 0.03 M. This range of
Al was selected to produce a set of goethites closely spaced in terms of Al-substitution. 62
At the beginning of the experiment the sample with the highest concentration of alurainate (S-2Q0) showed the largest volume of precipitate. The volume decreased sharply in the following days and at the end of the equilibration period was the lowest among all samples.
The pigmentation of the samples changed around the third day from an initially dark brown to a yellowish brown color. Only sample S-200 maintained a reddish coloration indicating that a concomitant formation of goethite and hematite was taking place. The synthesis procedure yielded approximately 9 g of oxide. The range of Al substitution was from 0 to 9.33 mole percent (Table 2.4).
A similar range of Al substitution was obtained by Schulze and Schwertmann (1986) following the same procedure.
All synthetic goethites were well crystallized as indicated by low oxalate-soluble/total Fe ratios
(Feox/Fet) (Table 2.4). Oxalate-soluble Fe gives an indication of the amount of "amorphous" iron oxide in the samples. Although the amount of oxalate-soluble Fe increased slightly with Al-substitution (S-150A, S-150B,
S-200), the content was very low indicating that the bulk of the Fe oxide was well crystallized. 63
Table 2.4. Some physical and chemical properties of Al-substituted goethites.
Mole Surface % Al Area Sample AA* Feox/Fet (m2/g) Mineralogy
S-0 0 4.30x10-3 42.0 Goethite S-10 1.07 5.00x10-9 41.0 Goethite S-30 1.65 4.00x10-3 40.0 Goethite S-80 4.75 2.90x10-3 36.0 Goethite S-120 6.54 2.90x10-3 28.0 Goethite S-150A 7.92 1.48x10-2 30.0 Goethite S-150B 7.98 1.05x10-2 32.0 Goethite S-200 9.33 3.73x10-2 Goethite (51%) Hematite (49%)
* Determined by atomic absorption spectrometry. 2.7.2 X-ray Diffraction Studies
X-ray diffraction (XRD) results not only serve as a fingerprinting technique to identify a mineral but also provide information about its crystallinity and particle size. X-ray diffraction patterns for all goethites displayed sharp peaks superimposed on a very straight baseline indicating once again the presence of well crystallized materials. Goethite was the only product in all samples except S-200 which also contained 49% hematite. The hematite/goethite ratio was taken from the relative intensities of the (012) line of hematite
(multiplied by 3.5) and the (110) line of goethite
(Schwertmann, 1988). A step scanned XRD pattern of pure goethite is shown in Fig. 2.9. The computer collected data was very useful to measure peak position, peak intensity, and peak width. All peaks were clearly resolved, even those of very low intensity like the (141) and (060). 65 55 45 35 25 15 Degrees 2-0 (Cu Ka) .9. X-ray diffraction pattern of synthetic goethite showing peak assignments for the different (hkl) planes. 66
2.-7 , 2 rJ. Effect of Al-substitution on lattice spacings. unit cell dimensions and mean crystallite dimensions (MCD)..
The Al3+ ion is slightly smaller than the Fe3 + ion o o (0.53 A vs 0.65 A); thus, substitution of Al for Fe should result in a decrease of the unit cell dimensions of the crystal and a contraction of the spacings between sets of parallel planes. For example, Fig. 2.10 shows a progressive shift of the (111) line of goethite to higher degrees 28 as Al content increases from 0 to 9.33 mole percent. This shift corresponds to a decrease in d- a spacing from 2.4532 to 2.4402 A. Also, a narrowing of the peaks is observed which indicates some improvement in crystallinity of the oxide with Al substitution.
The unit cell dimensions of goethite were calculated using the standard equation for relating d-spacings to the unit cell dimensions of an orthorhombic crystal (see section 2.6.2.2). The a and b unit cell dimensions shown in Table 2.5 represent the average of seven different values calculated using the d-spacings of fifteen goethite peaks. For the c-dimension, six different values were averaged. goethite with aluminum ShiftFig. in 2.10. substitution. the position of the 111 peak Counts/40 seconds 3000 2000 2500 1000 1500 500 5536.0 35.5 S0 CNo fll3 S-0 o ere 2-Theta Degrees 653. 7538.0 37.5 37.0 36.5
68
Table 2.5. Unit cell dimensions and mean crystallite dimensions (MCD) of Al-substituted goethites.
Unit Cell ..Dimensions Mean Crystallite Mole (A) Dimensions (nm) % Al Sample XRD* a b c MCDa MCDb
S-0 0 4.611 10.071 3.029 21.7 57.4 S-10 1.71 4.611 10.058 3.026 21.4 57.3 S-30 2.86 4.627 9.981 3.024 24.7 80.4 S-80 5.14 4.618 9.995 3.020 32.4 57.4 S-120 8.00 4.598 9.958 3.015 33.0 114.9 S-150A 9.14 4.601 9.950 3.013 38.0 73.1 S-150B 8.00 4.600 9.952 3.015 40.8 114.9 5-200 12.57 4.641 9.904 3.007 42.8 89.4
* Mole % Al = 1730-572.0c, where c is thei unit cell dimension in A (Schulze, 1984). 69
An example showing how a value of the a-dimension was obtained from d(120) and d(020) follows:
(1) di20 = [1/a2 + 4/b2]-H
1/(di2 o)2 = l/a2 + 4/b2
l/a2 = 1/(di2o)2 - 4/b2
(2) a = [1/(di2o)2 - 4/b2]-H
Since,
(3) d020 = [4/b2]-H,
(4) 4/b2 = 1/(do 2 o )2 and equation (2) becomes
(5) a = [1/(di2o)2 - l/(d020)2]-*.
Similar relationships can be developed to calculate the c-dimension from, for example, the d(020) and d(021) lines. The b-dimension was calculated from d(020), d(040) and d(060) using a and c values calculated from other peaks.
The b- and c-dimensions of goethite decrease linearly with Al substitution (Table 2.5, Fig. 2.11) showing negative r-values of -0.913 and -0.982, respectively. The a-dimension shows no correlation with Al substitution.
The lack of correlation of the a-dimension of goethite 70
4.63
4.50
4.57
4.55 0.0 2.0 4.0 6.0 8.0 10.0 Hole Z a I uni run eubetltution 10.10
10.06
0.00
0.85 0.0 2.0 4.0 6.0 8.0 10.0 Mole Z alunlnun sUratl tut Ion
3.03.
3.02 t d u 3.01
Cr—0.982 D
3.00 0.0 2.0 4.0 6.0 8.0 10.0 Hole Z aluninum substitution
Fig. 2.11. Unit cell dimensions of synthetic goethites as a function of Al substitution. 71 with Al substitution has been attributed to structural defects of the crystals which are more pronounced along the a-axis. Schulze (1984) suggested that the stacking of the double chains of [Fe,Al(0,0H)s] octahedra along the a- axis could be easily disrupted during crystal growth since the H-bonds that link the chains together are the weakest bonds in the structure. Such a disruption of the structure would not take place along the b- and c-axes where covalent bonds predominate.
The strong linear dependency of the c-dimension on Al substitution should make it an accurate means of estimating Al substitution in goethites. Schulze (1984) developed the following equation relating the c-dimension with mole % Al substitution:
Mole % Al = 1730 - 572.0c
This equation shows that an unsubstituted goethite should « have a c-dimension of 3.024 A. Since the calculated c-dimension for the unsubstituted sample in our study was O larger (3.029 A) the data was normalized by multiplying the c-dimensions by (3.024/3.029). The equation still overestimates Al substitution in all samples (Table 2.5) 72
relative to data acquired by total chemical analysis
(Table 2.4). Schulze and Schwertraann (1986) reported that
the equation predicted Al substitution within ±2.6 mole %
at the 95% confidence level. In the present study, mole %
Al estimated from the equation lies within this range for
all samples except S-200 which shows a difference of 3.24
mole % Al. This sample also contains 49% hematite. The
incorporation of Al into the structure of hematite is
always lower than in goethite (Fysh and Fredericks, 1983).
Thus, it is likely that the chemical analysis
underestimates the mole % Al in goethite due to the
presence of hematite.
X-ray diffraction also provides information about the
average size of the crystals, or more exactly, that of the
coherently scattering domains within the crystals (Schulze
and Schwertmann, 1984). The Scherrer formula (see
section 2.6.2.2) gives the mean crystallite dimension
(MCD) perpendicular to a given plane and also allows for
the calculation of MCD along the crystallographic axes.
The MCD's parallel to "a" and "b" give measures of the
thickness and the width of the goethite crystals,
respectively. 73
Incorporation of Al into the structure of goethite results in the formation of thicker crystals as indicated by a progressive increase in MCDa (Table 2.5, Fig. 2.12).
The thickness of the crystals increases from 21.7 to 42.8 o A as Al substitution increases from 0 to 9.33 mole percent. Crystal width (MCDb) also increases with Al substitution but the correlation (0.633) is not nearly as good as that between Al substitution and crystal thickness
(0.984) (Fig. 2.12). This suggests more variability in crystal width than in crystal thickness.
2.7.3 Surface Area of Al-goethites
Aluminum substitution in the range from 0 to 9 mole % resulted in a decrease in the specific surface area of goethite (Table 2.4, Fig. 2.13). This decrease is in agreement with XRD data which indicates that Al substitution increases crystallinity and results in the formation of larger crystals. Schulze and Schwertraann
(1986) also reported a decrease in surface area with Al substitution from 0 to 11.6 mole %; however, further incorporation of Al into the goethite structure resulted in an increase in surface area (Schulze and Schwertmann, 74 45.0
40.0 n | 35.0 W a 8x . o
20.0 0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 0.0 10.0 Mole Z Rluminum Substitution 125.0
110.0 r'l 05.0
80.0
65.0
50.0 0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 0.0 10.0 Mole Z Rluminun Substitution
Fig. 2.12. Mean crystallite dimensions (MCD) of synthetic goethites as a function of Al substitution. Surface Rrea Cm2/g oxide) a function of fll SurfaceFig. 2.13.area substitution.of synthetic goethites as 45.0 40.0 35.0 5 2 30.0 00 . 20 . 40 . 60 . 80 . 10.0 9.0 8.0 7.0 6.0 5.0 4.0 3.0 2.0 1.0 *0.0 q ------1 ------1 Mo Ie ------Z 1 ------flI urn i inum Subs it tut on 1 ------1 ------1 ------■ BET Method ■BET C r=-0.941) f ------1 ------1 ------
76
1984; Schulze and Schwertraann, 1986; Golden, 1978) due to sharp decreases in particle size. As Al concentration in the synthesis solutions increases, the possibility of obtaining more "amorphous" (oxalate-extractable) material also increases (Goodman and Lewis, 1981).
2.7.4 Transmission Electron Microscopy
Transmission electron microscopy (TEM) provides information about particle size and shape that complements
XRD and surface area data. Crystals of goethites with no or low Al substitution consist of several domains parallel to the c-axis giving the crystals a stacked appearance
(Fig. 2.14). The domains are of unequal length (see arrow on S-0) which results in crystals with stepped ends.
These crystals also appear to contain small pores running along the c-axis.
Aluminum substitution causes a clear decrease in the multidomainic morphology. The crystals become shorter with a more uniform appearance and smoother ends (Fig.
2.14, Table 2.6). The crystals are assumed to be lying with their a-axis perpendicular to the grid surface.
Oriented this way, the width of the crystals measured from 77
S-30 1.7%AI
S-150B S-200 8.0% Al 1 9.3% Al
0.5 um
Aluminum-substituted Goethites
Fig. 2.14. Electron micrographs of Al-substituted goethites. 78
Table 2.6. Average widths and lengths of particles from Al-substituted goethites as measured from electron micrographs and ratios of widths from XRD(MCDb) and TEM.
Width (nm) Length (nm) MCDb
Sample Average S.D.* Average S.D.* n# Width TEM
S-0 99 26 1084 356 15 0.58 S-10 120 28 996 273 15 0.48 S-30 116 28 846 388 15 0.69 S-80 122 39 754 226 10 0.47 S-120 117 53 516 208 15 0.98 S-150A 106 34 564 179 15 0.70 S-150B 119 52 499 148 15 0.97 S-200 98 34 378 141 15 0.91
* S.D.=Standard Deviation # n=Number of particles measured 79 the micrographs represents the length along the b-axis, a value that should agree with the MCDb calculated from the
Scherrer formula. Crystal widths measured from the electron micrographs are summarized in Table 2.6. For samples with low Al substitution (S-0---S-80) the measured widths are much larger than MCDb determined from XRD as indicated by the low MCDb/width TEM ratios. As Al substitution increases the agreement between the two measurements becomes better (samples S-120, S-150B, and S-
200) thereby supporting an increase in monodomainic character. In contrast, the unsubstituted sample and those with low Al are multidomainic and the value of MCDb apparently reflects the size of those domains, not the size of the multidomainic crystal.
The presence of hematite in sample S-200 is also confirmed by TEM. The hematite particles show a rounded or pseudohexagonal shape (see arrow in Fig. 2.14) and are thicker than the goethite particles as indicated by the inability of the electron beam to penetrate the crystals. 80
2.7.5 Dehydroxylation Properties of Al-goethites
Aluminum substitution improved the thermal stability of goethite. Differential scanning calorimetry patterns
(Fig. 2.15) show an increase in the bulk dehydroxylation temperature from about 260 to 350 SC as Al substitution increases from 0 to 9.33 mole %. The dehydroxylation endotherm of most samples is split into two or more components. A splitting of the endotherm into two components has been reported previously (Creer et al.,
1971; Derie et al., 1976; Murad, 1979; Schwertmann, 1984).
The low Al-substituted samples (S-0, S-10, S-30, and
S-80) show a more pronounced splitting of the endotherm than samples S-120, S-150A, S-150B, and S-200. The former samples show a small but well resolved peak around 230 *C and a splitting of the very intense bulk dehydroxylation peak around 300 °C. For example, sample S-0 shows a splitting of the latter into two poorly resolved peaks at
275 and 310 °C. The proportion of the high temperature component increases with Al substitution and the double peak eventually evolves into a sharper and more symmetric single peak around 350 °C. MOLE % Al
S-0 1.07
S-10 1.65
S - 3 ll> 4.75
S-8Q. 6.54
S-120
7.92 S-150A 7.98 S-15
9.33 S-20
150 200 250 300 350 400 TEMPERATURE *C
Fig. 2.15. Differential Scanning Calorimetry CDSC) pattern of fll- substituted goethites. 82
The splitting of the dehydroxylation endotherm of goethite indicates a momentary decrease in the dehydroxylation rate which has been associated with the degree of crystallinity of the sample (Derie et al., 1976;
Schwertmann, 1984). Derie et al. (1976) suggested that during heating of well crystallized samples of low surface area (24 - 68 m2/g), a hematite shell forms at the beginning of the dehydration process which retards further dehydration of the remaining goethite, the result being the splitting of the bulk dehydroxylation peak. On the other hand, Schwertmann (1984) showed that in a series of goethites with surface area ranging form 12 to 156 ra2/g, the splitting of this peak was most pronounced in samples of intermediate (40 - 60 m2/g) surface area. Samples with surface area below and above this range showed a single dehydroxylation peak. Schwertmann (1984) attributed the splitting of the bulk dehydroxylation endotherm to a contraction of the c dimension during heating which led to the formation of an intermediate goethite with slightly higher dehydroxylation temperature. The absence of a double peak in samples of high surface area was attributed to a rapid and complete transformation to hematite which 83
did not allow any changes in the unit cell to occur; however, no explanation was given as to the absence of a double peak in samples with surface areas below 40 m2/g.
The small endothermic peak around 230 *C shows a general decrease in intensity as Al substitution increases and also shifts to slightly higher temperatures. Samples
S-0 (0 mole % Al), S-10 (1.07 mole % Al), and S-30
(1.65 mole % Al) also show a shoulder around 250 ®C which becomes less pronounced or dissapears at higher Al substitution.
The occurrence of endothermic peaks at these low temperatures(230 - 250 °C) is not well documented. Golden
(1978) observed three small peaks in this region that he attributed to the dehydroxylation of A, C, and B - type surface OH goups. However, attempts to remove specific kinds of hydroxyls by heating the samples to appropriate temperatures failed. Infrared spectra indicated the dissappearance of all surface hydroxyl bands after heating to a temperature where only A-type OH should have been removed.
The presence of these low temperature endothermic peaks, as well as the broader and more assymetric peaks of the least Al-substituted samples seems to be related to 84 both the morphology and crystallinity of the sample. The unsubstituted (S-0) and low Al-substituted goethites
(S-10, S-30, and S-80) are multidomainic in contrast to the more monodomainic nature of S-120, S-150A, S-150B, and
S-200. The interdomainic regions are regions of imperfect crystallization (Schwertmann, 1985) and any hydroxyls in those regions of the crystal should be removed with more ease upon heating than hydroxyls located within the domains. Also, the presence of pores in these samples may result in entrapment of water that presumably will be removed at lower temperatures than the bulk hydroxyls.
These two factors could result in changes in the dehydroxylation rate and thus the occurrence of the endothermic peak around 230 *C as well as the broader and more asymmetric peaks.
2.7.6 Conclusions from the Synthesis and Characterization of Al-substituted goethites
Incorporation of Al into the structure of goethite
(0-9.33 mole %) improved the crystallinity of the mineral, decreased the unit cell dimensions and surface area, and influenced markedly the morphology of the crystals. With 85
increasing Al substitution the crystals became shorter and
thicker and changed from a predominant multidomainic
morphology at the lower Al range (0-4.75 mole %) to a
raonodomainic morphology in the upper range (6.54-9.33
mole %).
The changes in morphology and surface area of the
oxide should influence its surface chemistry and could
confound any effects due purely to the substitution of Al
for Fe at surface adsorption sites. In an attempt to overcome this problem a second series of synthetic goethites of varying crystallinity and surface area but no
Al substitution was prepared by hydrothermally treating subsamples of unsubstituted goethites in a pressure bomb at different temperatures (see section 2.6.1.2).
2.7.7 Synthesis and Characterization of Hydrothermally- treated Goethites
Hydrothermal treatment improved the crystallinity of the goethite samples as indicated by a progressive decrease in Feox/Fet ratios with increasing temperature
(Table 2.7). Improved crystallinity is also evident from a narrowing of the XRD lines (Fig. 2.16). For example, the measured width at half height (WHH) of the (110) peak 86
Table 2.7. Oxalate soluble Fe content, surface area and mineralogy of hydrotherraally-treated goethites.
Surface Area Sample Feox/Fet (m2/g) Mineralogy
S-0 4.35x10-3 38.0 Goethite S-OC (50 °C) 3.77x10-3 35.0 Goethite S-OC (100 •C) 3.33x10-3 35.0 Goethite S-OC (150 °C) 2.19x10-3 23.0 Goethite S-OC (200 °C) 1.15x10-3 16.0 Goethite*
* Sample contained a trace of hematite. Counts/40 seconds 2000 1000 1500 500 20.0 i , .6 VariationFig, 2.16, in the width of the C1103 peakof goethite S-OC WHH=0.47 20.5 ere 2 Theta 2 Degrees 21.0 with hydrothermal treatment. 21.5 ■ S-OC (200 °CD (200 S-OC■ WHH=0.29 22.0
88
of goethite decreased from 0.47 degrees 29 for the
starting material (S-OC) to 0.29 degrees 29 for the sample
treated at 200 °C. The (110) peak also shows a shift to
higher degrees 29 with hydrothermal treatment. There is
the possibility that the untreated sample contains
adsorbed water between octahedral layers which may result
in a larger d-spacing. X-ray diffraction analysis also
indicated the presence of a trace amount of hematite in
sample S-OC (200 °C). Heating of a subsample to 250 °C
for 7 days resulted in a complete conversion of the goethite to hematite.
The surface areas of the samples decreased from 38.0 m2/g for S-OC to 16.0 m2/g for S-0C(200 °C) (Table 2.7) which is consistent with an increase in the mean crystallite dimensions (MCDa and MCDb) as determined by
XRD analysis (Table 2.8). The effects on MCD achieved with hydrothermal treatment are similar to those induced by Al substitution; however, no significant decrease in unit cell dimensions was observed in the hydrothermally- treated samples.
Transmission electron micrographs of this goethite series indicate that the crystals of the starting material
(S-OC) are multidomainic with step-shaped ends (Fig.
2.17). The multidomainic structure is still clearly 89
Table 2.8. Unit cell dimensions and mean crystallite dimensions (MOD) of hydrothermally-treated goethites.
Unit Cell Dimnesions Mean Crystallite Sample (A) Dimensions (nm)
a b c MCDa MCDb
S-OC 4.611 10.060 3.023 21.9 61.9 S-OC(50 ’C) 4.611 9.901 3.022 19.7 50.3 S-OC(100 •C) 4.621 10.022 3.025 22.0 80.4 S-OC(150 •C) 4.601 9.955 3.024 28.4 67.0 S-OCC200 •C) 4.603 9.901 3.020 40.6 100.5
i Hydrothermally-treated Goethites
Fig. 2.17. Electron micrographs of hydrothermally- treated goethites. 91
visible in samples treated at 100 °C but is very subdued
in samples S-QC(150 °C) and S-0C(200 °C). These crystals
show a more uniform appearance and have smooth, rounded
ends. This morphological change agrees with the drastic
decrease in surface area observed for these two samples
(Table 2.7).
The higher initial surface area can be attributed to
the step-shaped ends of the crystals and to the presence of micropores between domains of the unsubstituted sample.
The hydrothermal treatment apparently induces dissolution of Fe from the ends and interdomainic areas of the crystals. This Fe is then reprecipitated along domain boundaries sealing the raicropores and causing a progressive "healing" of the crystals.
The average widths and lengths of the crystals obtained from electron micrographs are shown in Table 2.9.
Both parameters increase with hydrothermal treatment; however, the low MCDb/Width TEM ratios indicates poor agreement between the two measures of crystal width.
Ratios less than unity suggest that the MCDb values do not represent the gross width of the crystals but rather the width of the domains within the crystals. This also suggests that while hydrothermal treatment improves 92
Table 2.9. Average widths and lengths of particles from hydrothermally-treated goethites as measured from electron micrographs and ratios of widths from XRD(MCDb) and TEM.
Width (nm) Length (nm) MCDb
Sample Average S.D.* Average S.D.* n# Width TEM
S-OC 97 24 924 353 15 0.63 S-0C(50 •C) 127 34 1121 283 15 0.39 S-0C(100 ’C) 141 43 1386 402 8 0.57 S-0C(150 •C) 150 34 1230 555 15 0.45 S-0C(200 *C) 150 42 1262 541 15 0.67
* S.D.^Standard Deviation # n=Nuraber of particles measured 93
crystallinity, decreases surface area and results in the
formation of thicker and more uniform crystals, the
multidomainic nature is preserved. Apparently, the
domains became somewhat thicker and the interdomainc
regions and micropores are sealed with reprecipitated Fe
making it more difficult to distinguish the domain
boundaries. Such an effect can be observed in Fig. 2.18.
The crystals of samples S-QC(150 *C) and S-0C(200 °C) are
thicker, wider, and have more uniform surfaces than
samples S-OC and S-0C(100 °C). Nevertheless, the domain
boundaries in the former can still be traced along the
needle axes.
The thermal stability of goethite improved with
hydrothermal treatment. The bulk dehydroxylation
temperature increased from 270 *C for the untreated sample
(S-OC) to 360 "C for sample S-0C(200 *C) (Fig. 2.19). The
splitting of the primary dehydroxylation endotherm is less pronounced in this series than in the Al-substituted
series but two low temperature endotherms between 200 and
270 °C persist throughout. Contraction of the c unit cell dimension upon heating as suggested by Schwertmann (1984) may be responsible for this stepwise lost of hydroxyls.
There is also a possibility that sealing of the micropores 94
s t s s i i
‘•tfy-wiss • f,"’- v,f;X'.v‘v W 8 M r i - ' ^ U 1
« ■
s-oc §^ S- f O C S i W S P 150 °C
m & & 9 « S ? : r/-
g&g O.ITjun Hydrothermally-treated Goethites
Fig. 2.18. Electron micrographs of single crystals of hydrothermally-treated goethites. 95
TEMPERATURE °C
untreated
50
100
150
200
150 200 250 300 350 400 TEMPERATURE «C
Fig. 2.19. Differential Scanning Calorimetry (DSC) pattern of hydrothermally-treated goethites. 96 through hydrothermal treatment resulted in water entrapment inside the crystals. This water would presumably be lost at a lower temperature(s) than the bulk hydroxyls.
2.7.8 Conclusions from the Synthesis and Characterization of Hydrothermally-treated Goethites
The changes in goethite properties induced by hydrothermal treatment were similar to those caused by Al substitution. The crystallinity of the samples improved, the crystals became thicker and wider, and their multidomainic character and surface area were reduced.
Thus, by comparing sulfate adsorption in these two series of goethites any effect on sulfate adsorption due purely to substitution of Al for Fe at surface adsorption sites should be easier to elucidate.
2.7.9 Characterization of the Nipe Clay
X-ray diffraction analysis of the Nipe clay confirms the presence of goethite and hematite in association with gibbsite and kaolinite (Fig. 2.20). The intensities of the goethite lines clearly indicate that it is one of the Counts /40 Seconds 0 0 4 800 600 200 Fig. K=kaollnlte. Gb=glbbslte. Gt=goethlte and Ht=hematIte. 65 2.20. 55 X-raydiffraction pattern ofNipe clay. Gt ere 26 (CuKal 2*6 Degrees 3545 Gt-Ht 25 Gt Gb 15 98 dominant minerals in the clay fraction. Using an equation developed by Schulze (1984) which relates the d(lll) value of the goethite to the mole Al content, it is estimated that the Nipe goethite contains approximately 19 mole % Al
(Table 2.10). The Nipe clay contains 26% (37% as Fe203) reductant soluble Fe (Fed). The low Feox/Fed ratio indicates that most of this Fe occurs as relatively well crystallized Fe oxides.
The surface area of the clay is larger than the surface areas of the synthetic goethites. This result is expected since soil minerals seldom achieve the degree of crystallinity and crystal size of synthetic samples. The
DSC pattern of Nipe clay (Fig. 2.21) shows an intense endotherm at 300 °C that is actually a doublet corresponding to the dehydroxylation of goethite and gibbsite present in the sample. Both minerals show a strong endothermic peak in the range of 270 to 300 ’C.
The endotherm centered at 525 ’C corresponds to the dehydroxylation of kaolinite. Table 2.10. Some chemical and physical properties of Nipe clay (< 2um).
Mole % Al Surface Area Fed Feox/Fed (Goethite)* («2/g)
25.9% 1.17x10-2 18.8 55.26
* Mole % A1=2O06-85O.7 d(lll) (Schulze, 1984) 100
0 100 150200 2 50 400300 4 5 0 5 0 0 too T lttPIM A T U M (C
Fig. 2.21. Differential Scanning Calorimetry (DSC) pattern of Nipe clay. CHAPTER III
INFRARED SPECTROSCOPY STUDY
Infrared (IR) spectroscopy deals with the interaction
of infrared radiation with matter. The IR spectrum of a compound can provide important information about its chemical nature and molecular structure. In adsorption studies IR spectroscopy can be used to study the chemical nature of the adsorbent surface, the adsorbate, and any new species that may form. Dispersive infrared spectroscopy has been used by several workers for the characterization of goethite and other iron oxides (Jonas and Solyroar, 1970; Russell, 1979; Fey and Dixon, 1981;
Schulze and Schwertmann, 1984; Cambier, 1986a,b,; Lewis and Farmer, 1986) and also to study the adsorption of anions on some of these oxides (Atkinson et al., 1974;
Parfitt et al., 1975; Parfitt et al., 1977; Parfitt and
Smart, 1977, 1978; Harrison and Berkheiser, 1982).
Although these studies have provided valuable information 102 about the chemical nature of the adsorbents and the mechanisms of anion adsorption, the quality of the IR spectra has not always been good. The intensity and resolution of the IR absorption bands, especially in anion adsorption experiments, has been rather poor.
In recent years Fourier Transform Infrared
Spectrophotometers (FTIR) have become more available to soil scientists and mineralogists. FTIR spectrophotometers present several advantages over dispersive instruments in terms of sensitivity, resolution, energy throughput and wavenumber accuracy
(Koenig, 1975). The FTIR can collect a whole spectrum in a matter of seconds, and with the use of computers many collected spectra can be coadded to improve the signal to noise ratio. In addition, collected spectra can be expanded or contracted, and specific regions can be selected for detailed analysis without the need of rescanning. Another advantage is the capability for difference spectroscopy. This method allows one to detect small changes in samples in terms of both composition and physical state. Common features in the spectra cancel, and bands which are recorded can be interpreted in terms of differences between samples. 103
The present study deals with the use of FTIR spectroscopy to characterize a series of unsubstituted and
Al-substituted goethites of varying crystallinity and surface area, and to study the mechanisms of sulfate adsorption on these oxides. The chemical nature of hydroxyl groups in goethite may change with Al substitution and crystallinity, which in turn may affect the capacity of the oxides to react with sulfate. FTIR spectroscopy can provide information about structural changes and can also be used to define the type of surface complexes that sulfate forms with the oxide.
3.1 Infrared (IR) Characterization of Goethite: a Review
The structure of goethite consists of a mixture of
OH" ions and water molecules coordinated to Fe3+ ions.
Water molecules are coordinated to metal ions exposed to the surface (Lewis acid sites) or can form H bonds to surface hydroxyls (Cornell et al. 1974, Parfitt, 1978).
Thus, the IR spectrum of the oxide is dominated by OH related IR active bands. The IR spectrum of goethite is characterized by a broad, intense absorption band around
3150 cm~i attributed to stretching of hydroxyls in the 104
bulk structure of the oxide and by two sharp bands around
900 and 800 cm-1 attributed to OH bending. These bands
are designated as vOH, 8OH and tfOH, respectively (Schulze
and Schwertmann, 1984; Cambier, 1986a). The 80H bending
band is due primarily to vibrations in the a-b plane and
#0H to vibrations out of the a-b plane (ie. along the c
axis). Two other sharp absorption bands are observed
around 600 and 400 cm-i which have been attributed to.Fe-0
stretching modes (Cambier, 1986a).
In addition to structural hydroxyls, three types of
surface hydroxyls reportedly occur in goethite. These
hydroxyls have been designated as A (coordinated to one Fe
atom), B (coordinated to three Fe atoms), and C
(coordinated to two Fe atoms) (see Fig 2.2). These
surface hydroxyls represent a minor portion of the total
hydroxyls in goethite, but they play a major role in
adsorption reactions. IR absorption bands corresponding
to stretching vibrations of the surface hydroxyls are
observed around 3600 (B and C type hydroxyls) and 3500 cm-1 (A type hydroxyls) (Russell et al., 1974; Parfitt and
Russell, 1977; Rochester and Topham, 1979). The low
intensity of these bands combined with their close proximity to the very intense vOH band makes their 105
detection very difficult. This may explain why most of
the IR characterization work on goethite has focused on
OH stretching (vOH) and OH bending (6H, tfOH) bands (Jonas
and Solymar, 1970; Fey and Dixon, 1981; Fysch and
Fredericks, 1983; Schulze and Schwertmann, 1984; Cambier,
1986a,b).
The removal of adsorbed water either by evacuation or
exchange with D 2O helps in the identification of OH
surface sites (Atkinson et al., 1974; Parfitt et al.,
1976; Parfitt, 1978). Exposure of goethite to deuterium
oxide vapor results in the disappearance of the surface OH
bands around 3600 and 3500 cnr* and the concomitant
appearance of OD bands around 2700 and 2580 cm-1 (Fig.
3.1). The OD bands occur in a region of the spectrum where the interference of the bulk hydroxyl band is
reduced which, in turn, facilitates their detection.
Incorporation of Al into the structure of goethite
results in shifting of the infrared bands. Jonas and
Solymar (1970), and Fey and Dixon (1981) reported an almost linear increase in the position of the OH bending
(60H, tfOH) bands with the 50H band (bending vibrations in
the a-b plane) being the most sensitive to Al substitution. The shifting of these absorption bands with 106
*BOO *400 9000 3600 9300 on*'
OH
Fig. 3.1. Infrared spectra of goethite showing surface OH bands, which exchange with D20 to give 0D bands (Parfitt. 1978). 107
A1 substitution has been attributed to enhanced strength of the structural H-bonds. The smaller A1 atom presumably causes the H-bonds to become shorter and consequently stronger. A stronger H-bond should also result in a shift of the vOH (stretching) bands to lower frequencies since weakening of the intraionic OH bond occur (Lutz et al.,
1987); however, this is not always the case. For example,
Mendelovici et al. (1979) reported a shifting of the vOH band to higher frequencies as Al substitution increased in some soil goethites. This observation led the authors to conclude that Al substitution in goethite actually weakens the H-bonds. They overlooked the fact that the OH bending bands shifted to higher frequencies which suggests the formation of a stronger H-bond.
These contradictory results suggest that factors other than Al substitution also contribute to shifting of the IR bands in goethite. Schul2 e and Schwertmann (1984) found that the OH bands were affected not only by Al substitution but also by structural defects, and that these two factors may counteract each other. Aluminum substitution led to a decrease of the OH stretching frequencies and to an increase in the OH bending frequencies whereas increases in structural defects had 108 the opposite effect. More recently, Cambier (1986b) found that a decrease in crystallinity of unsubstituted goethites resulted in an increase of the OH stretching frequencies and a decrease in the bending frequencies.
The author also suggested that some types of crystal disorder resulted in weaker H-bonds which, in turn, caused the observed shift in OH frequencies.
3.2 Infrared Spectroscopy Studies of Anion Adsorption
Most infrared studies of anion adsorption on iron oxides have focused on phosphate. Phosphate reacts strongly with oxide surfaces forming inner sphere complexes. Infrared spectroscopy studies have indicated that phosphate undergoes ligand exchange reactions with hydroxyl groups at the oxide surface (Atkinson et al.,
1974; Parfitt et al., 1975; Parfitt et al., 1976). The spectra of phosphated goethites have typically shown three bands at 1190, 1100 and 1000 cra-i which agree with reported band positions of a well characterized bidentate bridging phosphate complex (Kumamoto, 1965).
Previous studies have involved air dried or evacuated samples; therefore, bidentate bridging complex may not be the actual complex formed when the anions are in solution. 109
Tejedor-Tejedor and Anderson (1986) studied the formation of phosphate complexes in goethite suspensions using FTIR spectroscopy. Their results showed the presence of phosphate bands at 1225, 1151, 1095, and 1040 cm-i. They concluded that inner sphere phosphate complexes were formed but made no specific assignment of these bands to monodentate or bidentate complexes.
Parfitt and Russell (1977) investigated the anion complexes formed when acids (H3PO 4 , H 2SO4 , H2Se03 and
(C00H)2) were evaporated onto goethite. A deuterated goethite with 400 nmol g-* of surface OD groups was used in the study. At 200 nmol g -1 H 3PO 4 , H 2SO4 , and H2Se03 adsorbed, the A-type OD band at 2584 cm-1 was absent or very weak. This suggested that like phosphate, selenite and sulfate were adsorbed as binuclear species on the goethite surface. The spectrum of goethite with 200 umol g_1 oxalic acid showed a broad peak at 2580 cm-1 indicating that, unlike the other acids, oxalic acid did not completely replace A-type OH(OD) groups. Oxalate seems to form both monodentate and bidentate complexes
(Parfitt et al., 1977a). 110
3.3 Infrared Spectroscopy of Sulfate and Sulfate Complexes
The "free" sulfate ion belongs to the high symmetry point group Td. It has two infrared active fundamental bands in the S=0 stretching region; V 3 and V 4 (Nakaraoto,
1986). Complex formation of sulfate with metal cations lowers the symmetry of the anion (Fig. 3.2). When a decrease in symmetry occurs, forbidden vibrations of the
"free" ion are permitted, and splitting of the infrared active bands is observed.
The IR spectrum of the "free" sulfate ion shows a sharp peak around 1100 cm-1 corresponding to the V 3 infrared active fundamental band (Table 3.1). The formation of a monodentate metal-surface complex with a
C3v symmetry results in a splitting of the V 3 fundamental band into two components around 1038 and 1130 crtri. A further decrease in symmetry to C 2v (bidentate complex) causes an additional splitting and the spectrum includes three peaks in the v3 region. Parfitt and Smart (1977,
1978) reported similar results to those shown in Table 3.1 in a study of sulfate adsorption on goethite. These authors observed three sulfate adsorption bands at 1295, Ill
V M ~ \ J >
/ ■ X / A , o o o *o
Free ion Unidenfafe complex < V (fa.)
M O \) M— 0 0— M .s. X O O 0 o
Bidcntnte complex Bridg'd liidciitute (Cjy) complex(V2v)
Fig. 3.2. Metal-sulfate complexes (Nakamoto, 1986). 112
Table 3.1. Infrared band positions of sulfate in selected complexes.
Anion Complex Symetry vs
Free sulfate Td 1104 (1)
Monodentate C0 -SO4 complex C3V 1038, 1130 (1)
Bidentate C0 -SO4 complex C2V 1055, 1105, 1170 (1)
Bidentate Fe-S04 complex C2V 1050, 1125, 1170 (2)
(1) Nakamoto (1986) (2) Harrison and Berkheiser (1982) 113
1145 and 955 cm _1. They attributed the vibrations near
1295 and 1145 to the free -S02- group and the one near 955 to Fe-O-S groups. The differences in the vibration frequencies reported by Parfitt and Smart (1977) and
Harrison and Berkheiser (1982) may be the result of adsorbed water. Harrison and Berkheiser (1982) found that the peak around 1050 shifted to 1040 cm-* and the peak around 1170 shifted to 1215 cm-l after evacuation of an
"amorphous" ferric hydroxide equilibrated with sulfate.
3.4 Materials and Methods
The infrared spectroscopy study was divided into two parts including: a) a FTIR characterization of the adsorbents, and b) a FTIR study of sulfate adsorption.
Most of this study was conducted in the Department of Soil
Science at the University of Florida using a Bohmen Model
DA 310 FTIR. This instrument was connected to a Vax station 2 Microvax computer for collection and manipulation of the data. Some characterization work was also performed in the Chemistry Department of The Ohio
State University using a Perkin Elmer 1710 FTIR. 114
3.4.1 FTIR Characterization of the Adsorbents
The adsorbents were characterized using both KBr
pellets and diffuse reflectance spectroscopy. For the
preparation of KBr pellets 1 mg of oxide or clay was added
to 500 mg of KBr and the mixture was finely ground in an
agate mortar. Three hundred mg of this mixture was
subsequently pressed in a 13-mra die under a load of 10
tons for 3 min. Twenty scans (interferograras) of 4 cm-1
resolution were collected using a Beckman 1710 FTIR and
averaged to obtained a primary spectrum in the range of
4000 to 400 cm-1.
Characterization by diffuse reflectance was performed
with the Bohraen Model DA 310 FTIR. The optical diagram of
the diffuse reflectance cell is shown in Fig. 3.3. Two
plane mirrors (Ai,2) direct the infrared beam to an
ellipsoidal mirror (Bi) located above the sample, which in
turn redirects the beam to the sample. The radiation
passes into the bulk of the sample where it undergoes
reflection, refraction, scattering and absorption to
varying degrees before re-emerging at the sample surface.
The light reflected in this manner is collected by another ellipsoidal mirror (B2 ) and is then redirected by another OUTPUT ELLIPSOID INPUT ELLIPSOID
SAM PLE C U P To From Detector Source
A 2
Fig. 3 .3 . Schematic diagram of the optical system of the diffuse reflectance accessary. 116 set of plane mirrors (A3,4) towards the detector. The whole diffuse reflectance accessory fits in the sample compartment of the FTIR.
This method requires very little sample preparation.
A 4% goethite dilution in KCl was finely ground in the agate capsule of a small vibratory mill. The powdered sample was then placed in the sample holder and gently pressed with a glass slide to create a smooth surface.
The sample holder was subsequently placed in the diffuse reflectance cell within the sample compartment of the
FTIR. All samples were evacuated to 0.2 torr before scanning. Five hundred scans (interferograms) collected at a resolution of 4 cm-1 were coadded to yield a primary spectrum in the range of 4000 to 400 cm-1. A reference spectrum of KCl was subtracted to yield the sample spectrum.
3.4.2 FTIR Study of Sulfate Adsorption on Goethite
The sulfate adsorption study was performed using a cylindrical internal reflectance cell (CIRCLE) (Fig. 3.4).
The optics system of this cell consist of a rod-shaped crystal of ZnSe assembled into an open boat structure.
Sample suspensions are placed in this open boat structure OMNMU'CIIL-
0 CONE roaoio QtRCLE OPTICS Fig. 3.4. Schematic representation of the cylindrical internal reflectancecell CCIRCLE3. 117 118 to contact and coat the ZnSe crystal. A set of mirrors at the entrance of the cell brings the IE beam into focus at an angle of 45* inside the crystal. The IR radiation undergoes multiple reflections inside the crystal which results in multiple contacts with the sample. The multiple reflections increase the energy input into the system. The IR radiation leaving the crystal is redirected to the detector by another set of mirrors placed at the exit of the cell.
For this study only two samples of each series were used. Oxides S-0(No Al) and S-150B(7.98 mole % Al) were selected from the Al-substituted series, and S-OC and
S-0C(200 *C) from the hydrothermal series. These samples represent the extremes of each series. In each case the cell was coated with a A % suspension of goethite and allowed to dry. The cell was then placed in the FTIR and evacuated to 0.2 torr. The recorded spectrum (4000 to 400 cm-1) was the result of 500 coadded scans, all ratioed against the spectrum of the empty cell. Eight ml of
0.0024 M Na2S04 solution (pH 4) were then added to the oxide coated cell and allowed to react for 30 min. The excess of solution was poured out, the cell was allowed to dry, and the spectra collected under the same conditions stated previously. 119
3.5 Results and Discussion
3.5.1 Characterization of the Adsorbents
The FTIR spectra of a synthetic goethite using KBr pellets and diffuse reflectance (DR) are shown in Figs.
3.5 and 3.6, respectively. The spectra of all samples
(Al-substituted and hydrothermally-treated) were identical in terms of the number of peaks observed except for example S-200 (9.33 mole % Al), which showed an additional peak at 565 cm*i due to the hematite present in the sample. Both types of spectra show OH stretching (vOH),
OH bending (5OH, tfOH) and Fe-0 stretching bands around
3140, 900, 800 and 640 cm-1, respectively. The small, broad bands between 1600 to 1700 cm-1 have been attributed to bending modes of adsorbed water (Farmer, 1974).
The DR spectra show two additional peaks that are not present in the KBr spectra. Diffuse reflectance clearly resolved the Fe-0 stretching frequency at 491 cm*1 and also enabled the detection of two surface hydroxyl stretching frequencies around 3660 cm*1. The peak around
3490 cm*1 corresponds to A-type hydroxyls and the one around 3661 cm*1 to B and C-type hydroxyls (Fig. 3.7). 1650
3140 643
798 895 o 4000 3600 3200 2800 2400 2000 1600 1600 1400 1200 1000 800 600 400 em-1 Fig. 3 .5 . FTIR spectra of synthetic goethite using KBr peI Ib ts.
to Kubelka Munk units 0.40 0.30 0.00 0.20 0.10 i . .. TR pcr o goethite of FTIR spectra 3.6. Fig. 400 ?S CM O) c*> O) oo D ifs Reflectance). CD iffuse ae u e (crrr1) ber num Wave 1600 CM a> 2800 r ( m c CM CO CO in oo m CO 4000 Kubelka Munk units 0.60 0.80 0.40 0.20 .0 I 0.001 . s d n a b showing andsurfacebulkhydroxyl Fig. 80 00 20 40 60 3800 3600 3400 3200 3000 2800 3.7. ______I - i I— ^ i i i - i I I FTIR spectraFTIRofgoethite v number cm-1) m (c r e b m u n ave W CM “ T CO o> CO in n (O
122 123
The large number of scans coadded and the more concentrated sample used with DR may have contributed to the enhanced resolution of these surface hydroxyl groups.
Fuller and Griffiths (1978, 1980) also observed that the band intensities in DR spectra of samples prepared by mixing an organic compound with KBr or KCl were larger than the band intensities for the same sample pressed into a halide pellet. Diffuse reflectance has the added advantage that pure samples can be analyzed. Such analyses are restricted to high frequency regions of the spectrum because resolution decreases sharply at lower frequencies.
Aluminum substitution resulted in a shift of the bulk
OH bending bands to higher frequencies (Table 3.2, Fig.
3.8) suggesting the formation of a stronger H-bond (Fey and Dixon, 1981). The OH bending frequencies in the a-b plane (6OH) are more sensitive to Al substitution than the bending frequencies out of the a-b plane C&OH). An increase in Al substitution from 0 to 9.33 mole % resulted in an upward shift in the 6OH band of 24 and 18 cm-1 using
KBr pellets and diffuse reflectance, respectively. The
OH frequency showed a shift of 19 cm-i with KBr and 10 cm- i with DR. Linear regression plots of these two 124
Table 3.2. IR absorption frequencies of Al-goethites.
Fe-■0 OH 6 OH V OH Mole % Sample Al KBr* DR# KBr DR KBr DR KBr DR
S-0 0 640 651 798 796 895 895 3140 3143 S-10 1.07 643 642 798 796 895 895 3140 3134 S-30 1.65 645 642 798 795 897 895 3150 3133 S-80 4.75 660 649 810 796 908 899 3190 3137 S-120 6.54 661 655 812 797 911 902 3200 3135 S-150A 7.92 663 653 814 799 915 904 3190 3136 S-150B 7.98 665 652 815 798 917 902 3150 3146 S-2D0 9.33 660 653 817 806 924 913 3190 3143
* KBr Pellets # Diffuse Reflectance Kubelka Munk units 0.20 05 .0 0 0.10 0.15 0.00 Fig. CaD S-0 (No fll); (b) S-150R (7.92 Mole ZRl). frequencies of goethite with fil substitution. 0 0 8 0 5 7 3 . 8 . in Shift the s-OH and 6-OH bending v number I -11 cm r e b m u n ave W 0 5 8 0 0 9 0 5 9
125 126 bending frequencies (Figs. 3.9, 3.10) indicate that 6OH is better correlated with Al substitution than tiOH and that the degree of shift varies consistently with the IR technique employed. Specifically, the shift detected from
KBr pellets is consistently larger than observed with DR.
It has been suggested that ion exchange and changes in crystal structure can occur while KBr pellets are being pressed. Both effects are minimized with the DR technique
(Fuller and Griffiths, 1980). Moisture may have also influenced the shift of these bending frequencies. Fey and Dixon (1981) found that removal of water by evacuation or heating of samples at 110 °C caused an upward shift in
8OH of about 4 cnr*. In our experiment the samples used to prepare the KBr pellets were dried at 110 ®C for 24 hr and evacuated during pressing. These two pretreatments may have been more effective in removing adsorbed water than the evacuation to 0.2 torr applied to the DR samples.
As noted previously (Sec. 3.1), a stronger H-bond in the goethite structure should also cause a downward shift of the OH stretching frequency (vOH); however, the position of the OH stretching band shows a general shift to higher wavenumbers with increasing Al substitution cm- 825 815 820 805 810 ygg ygg 795 800 . 10 . 30 . 50 . 70 . 90 10.0 9.0 8.0 7.0 6.0 5.0 4.0 3.0 2.0 1.0 0.0 Fig. fgeht ihH substitution.HIwithgoethite of ------Kr e lets KBr Pel o Dfue Reflectance Diffuse ■ 3.9. 1 ------hf in tf-OH the Shift vibrationsbending 1 ------Mo Ie
1 ------X flIum inum Subs t i itut on 1 ------1 ------1 ------1 ------1 ------
1 ------127 930 o KBr pel lets C r=0.987 D 925 ■ Diffuse Reflectance 0=0.899!) o 920
915
910
905
900
895<
890 0. 3 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0 10 Mo Ie X RI urn i num Subs t i tut i on Fig 3.10. Shift in the 6-0H bending vibrations goe hite with fll substitution. 129
(Table 3.2). The relationship is not particularly good
with r values of 0.708 and 0.773 for KBr and DR,
respectively. Similar results were reported by Schulze
and Schwertmann (1984) who also noticed that the position
of this band varied considerably even for unsubstituted
samples. These authors suggested that some structural
defects in the samples caused an overall weaker H-bond.
A weaker H-bond and structural defects in the Al-
substituted samples seem unlikely. All previous
characterization data (XRD, DSC, TEM, and surface area
measurements) indicate an improvement in crystallinity
with Al substitution. Also, the position of the very
intense and sharp OH bending bands show an upward shift
suggesting a stronger H-bond. Considering the rather poor
relationship between Al substitution and the position of
the vOH band, plus the fact that we are measuring the
position of a very broad absorption band, the apparent
shift to higher wavenurabers may be within the range of
measurement error.
The position of the Fe-0 stretching band also shifts
to higher wavenumbers with Al substitution (Table 3.2).
This observation is in agreement with a stronger Al-0 bond
(507.5 ± 10.5 KJmoi’i) compared to Fe-0 bond (408.8 ± 13
KJmoi-i) (Kerr, 1984). 130
The hydrothermal treatment of goethite improved the crystallinity of the oxide and resulted in an enhancement of the H-bond strength as indicated by a shift of the 5OH band to higher wavenumbers (Table 3.3, Fig. 3.11). Very little change in the position of the #0H band was observed. The position of the 6OH band is well correlated with temperature as indicated by r values of 0.929 and
0.946 for the KBr and DR data, respectively (Fig. 3.12).
The full magnitude of the shift in 6OH is comparable to that observed with Al substitution. In contrast to the
Al-goethites, KBr shifts are consistently less than obtained by DR. The reason for this difference is not apparent. Perhaps, the hydrothermally-treated goethites retain less adsorbed water due to their lower surface area and this water is effectively removed by evacuation (see p. 126).
The formation of a stronger H-bond as crystallinity improves indicates that a shortening of these bonds has taken place. It is possible that during hydrothermal treatment some tilting between pairs of octahedra takes place reducing the angle that the H-bond forms with the
O i-Oii vector (see section 2.1). A smaller angle results in a shorter and stronger H-bond. The formation of a stronger H-bond is also indicated by a shift of the vOH 131
Table 3.3. IR absorption frequencies of hydrothermally- treated goethites.
Fe--0 8 OH 6 OH V OH
Sample KBr* DR# KBr DR KBr DR KBr DR
S-OC 642 645 796 796 892 894 3142 3135 S-0C(50 ’C) 642 642 795 792 892 692 ,, 3139 3126 S-0C(100 •C) 641 641 796 795 892 893 3139 3136 S-0C(150 *C) 637 646 797 797 894 897 3138 3128 S-0C(200 *C) 634 635 797 796 897 898 3132 3127
* KBr Pellets # Diffuse Reflectance Fig. ihhdohra tetet (a)S-OC;with hydrothermal treatment.
Cb) S-OC (200 °C3. Kubelka Munk units 0.30 0.45 5 50 5 70 5 950 850 750 650 550 450 3 . 11 . inShift the IR frequencies ofgoethite Wave number number Wave vO n m in cm-1! m [c CO ' O CO ' O 00 “ b
132 i . .2 Sit in Shift the 6-OH Fig.bending 3.12.vibrations of goethitewith hydrothermal treatment. cm- 896 898 900 894 890 50 o o ■Diffuse Reflectance KBr Pel lets 75 yrteml ramet CD C ent Treatm Hydrothermal 100 125 150 175 200
134 band to lower wavenumbers (Table 3.3, Fig. 3.13), an effect that was not observed with A1 substitution.
The position of the Fe-0 stretching band shifted to lower wavenumbers with hydrothermal treatment suggesting an increase of this bond length with improved crystallinity as the pairs ofe octahedra become more tilted and the H-bond becomes stronger. Cambier (1986b) observed a similar decrease in the frequency of this band and suggested that the frequency may not correspond to a pure stretching mode along this single bond.
The IR spectrum of the Nipe clay (Fig. 3.14) indicates the presence of kaolinite, gibbsite, goethite and hematite in agreement with the XRD analysis. Although kaolinite is not the major mineral component of the clay, it dominates the IR spectrum with very intense absorption peaks in the range of 3600 to 3700 cm-1 and 900 to 1100 cm-1. The dominance of kaolinite bands in the latter region would clearly hinder or prohibit the detection of adsorbed sulfate on goethite.
3.5.2 FTIR Study of Sulfate Adsorption on Goethite
The sulfate adsorption study was performed using a cylindrical internal reflectance cell (CIRCLE) which is a Kubelka Munk units 0.30 0.20 0.00 0.10 0 0 7 3 0 0 5 3 0 0 3 3 0 0 1 3 0 0 9 2 0 0 7 2 Fig. r amn . (a) S-OC; (b) S-OCtreatment. frequenciesof goethite with hydrothermal 3.13. hf in Shift the vOH (stretching) v number (cm*1) r e b m u n ave W CM (200°C).
135 Fig. 3.14. FTIR spectra of Nipe clay (Diffuse Reflectance) Reflectance) (Diffuse clay Nipe of spectra FTIR 3.14.Fig. K=kaolinite, Gb=gibbsite, Gt^gaethite and Ht^hematite. and Gt^gaethite Gb=gibbsite, K=kaolinite, Kubelka Munk units 0.35 0.30 0.25 0.20 0.05 0.10 450 Z I Gb Gl fcm-1) r e b m u n e v a W 1650 40502850 137 relatively new variation of internal reflectance spectroscopy. The internal reflectance technique involves passing the infrared radiation through an infrared transmitting crystal of high refractive index, and allowing the radiation to reflect off the crystal walls many times. The reflected radiation penetrates the surrounding medium producing a spectrum of the sample in contact with the crystal.
The CIRCLE cell is especially designed for use with
FTIR spectrophotometers. This cell presents several advantages over previous internal reflectance cell designs. In addition to the higher energy throughput of the system, the configuration of the sample chamber (open boat structure) allows for free flow of the sample so that it fills without trapping bubbles and empties completely, leaving no pockets of residual sample. Both are problems often encountered with other internal reflectance cell designs.
One of the major advantages of the CIRCLE cell is that it allows for analysis of suspensions practically without any sample preparation. Suspensions can also be dried on the surface of the crystal for analysis. Since this IR technique does not require dilution of the sample with foreign substances like KBr or KC1, it can be 138 considered a more "natural" way to analyze properties of adsorbents and adsorption reactions.
The spectrum of a Na2S04 solution (0.1 M) collected with the CIRCLE cell is shown in Fig. 3.15. This spectrum was obtained after subtracting the spectrum of the CIRCLE cell + water. A very intense sulfate peak is observed at
1099.5 cra-i which corresponds to the V3 fundamental vibrational mode of the anion (Nakamoto, 1986). Any changes in the symmetry of the sulfate molecule upon adsorption on goethite will result in a splitting of this peak, which in turn supplies information about the type of complex formed (see Table 3.1).
The FTIR spectrum of sulfate adsorbed on goethite
(Sample S-0) shows a splitting of the V3 stretching band into three components at 1111.5, 1135.7, and 1169.2 cm-1
(Fig. 3.16). This indicates that the symmetry of the anion has been reduced from Td to C 2v upon bonding to Fe through two oxygens (Parfitt and Smart, 1977; Nakamoto,
1986).
Nakamoto (1986) suggests that it is possible to distinguish between bidentate and bidentate bridging sulfate complexes (Fig. 3.17), both with C2v symmetry, on the basis that the highest SO frequency of the bidentate groups lies above 1200 cm-i while the bridging complexes 139 0.1128
0.0771
& 0.0414
0.0057
- 0.0300 700 1800 2900 4000 Wave number (cm”1) Fig. 3.15. FTIR spectra of 0.1M sodium sulfate (CIRCLE cell). 140
S -0 0% Al
o u C D n
o «/) <
930 1030 1130 1230 1330 W ave number (cm-1)
Fig. 3.16. FTIR spectra of sulfate adsorbed on synthetic goethite (S-0). (a) oxide? (b) oxide + S04? (c) b - a. (CIRCLE CELL) a b Fig. 3.17. Schematic representation of S04-oxide complexes, (a) bidentate. (b) bidentate bridging. 142 show frequencies below 1200 cm-1. Using this criterion our results indicate the formation of a bidentate bridging
S04-goethite complex with sulfate replacing two OH groups from adjacent Fe atoms; however, such a distinction between the two types of complexes should probably be made with caution. For example, Harrison and Berkheiser (1982) found that the SO frequency of the bidentate bridging S04-
Fe oxide complex shifted from 1170 to 1215 cm-1 just by submitting air dry samples to evacuation.
There is the possibility of having some monodentate oxide-sulfate complexes if protonated sulfate (HS04-) is adsorbed instead of SO42-. The syraetry of this monodentate complex is C 2v like the bidentate complex.
However, the contribution of the former species to the observed peaks should be minimum under the pH range
(4 - 6.5) used in the experiment.
Surface OH groups absorb radiation in the range between 3500 and 3600 cm-i. Although the existence of three different types of hydroxyl groups in goethite has been proposed (Russell et al., 1974), only two OH stretching frequencies are clearly observed around 3660 and 3490 cnr* (see Fig. 3.7). Parfitt et al. (1976) found that F“ completely replaced A-type OH groups in goethite, 143 but little or no replacement of B or C-type OH occurred even when the F" concentration was increased to 1000 iimol g-1. This led the authors to conclude that A-type OHs are the only surface groups significantly involved in anion adsorption. Similar results were observed after phosphate adsorption.
Any anion undergoing ligand exchange reactions with a particular type of surface hydroxyl should decrease the intensity of the corresponding surface OH adsorption band.
Infrared spectra of the surface OH region of goethite before and after SO 42- adsorption are shown in Fig. 3.18.
The intense, broad band around 3135 cra-i (a,b) corresponds to stretching vibrations of bulk hydroxyls. A band of very low intensity corresponding to surface OH groups can be observed at 3638 cm~i. After sulfate adsorption (b), this small band disappears indicating that sulfate has replaced these surface hydroxyls. The spectrum shown in c is a differential or subtraction spectrum. The intense absorption band of the bulk OH groups, a feature common to both spectra a and b, is no longer present and only a small negative band indicative of the replacement of surface OH by SO 42- is visible around 3638 cm~i. 144
S-0 0% Al
0) v C o
k. 3638 \ f l - O «/> <
3638
2925 3125 3325 3525 3725 W ave number (crrr1)
Fig. 3.18. FTIR spectra of surface and bulk hydroxyl region of synthetic goethite. (a) oxide; (b) oxide + S04. Cc) b - a. (CIRCLE cel I)> 145
The differential or subtraction spectrum is an
excellent tool for resolving small treatment effects. To
observe in more detail the negative peak around 3638 cm-i,
the 3300 to 3900 cm-i region of the subtraction spectrum was expanded (Fig. 3.19). A sharp negative peak is
observed around 3638 cm-i which corresponds, supposedly,
to B and C-type OH groups. In addition, a small negative peak is now evident at 3481 cm-*, which agrees well with the vibrational frequency of the proposed A-type OH.
These results suggest that B and C-type OH groups also play a significant role in anion adsorption, contrary to the results reported by Parfitt et al. (1976).
The spectrum of surface adsorbed S042~on sample S-
150B (7.98 mole % Al) is shown in Fig. 3.20. Although of lower intensity than the 3042- bands in sample S-0, the same splitting of the v3 band into three components is observed, again suggesting the formation of a bidentate bridging complex. Sample S-150B has lower surface area than S-0 (32.0 vs 42.0 m2/g). The lower surface area may have resulted in a lower total S042- adsorption; however, that conclusion can not be stated with certainty from the
IR spectra alone since the amount of oxide coating the
ZnSe crystal may have varied between samples. A b sorbance -0.0230 -0.0205 -0.0130 -0.0105 -0.0155 0.0180 30 50 70 3900 3700 3500 3300 after S04 adsorption (CIRCLE ceil). (CIRCLE adsorption S04 after goethite of spectra subtraction FTIR 3.19.Fig. S0 soigngtv ufc yrxl bands hydroxyl surface negative showing (S-0)
v nmbr (cm*1) ber num ave W
3481 146 147
J S-150B S 7.98% Al_
a U C o jD
930 1030 1130 1230 1330 Wave number (cm-1)
Fig. 3.20. FTIR spectra of sulfate adsorbed on synthetic goethite CS-I5QB). (a) oxide; Cb) oxide + S04; Cc) b - a. (CIRCLE CELL) 148
The subtraction spectrum of the surface OH region shows a single negative peak around 3640 cm~i corresponding to B and C-type OH's, but no negative peak corresponding to A-type OH is observed (Fig. 3.21). The decrease in surface area and multidomainic character of the sample with Al substitution may have reduced the proportion of A-type over B and C-type hydroxyls or has made the former less reactive. The traditional view of goethite morphology (Fig. 2.3) includes prominent (100),
(001) and (010) faces with the former dominating the total surface area and supposedly displaying a combination of A,
B, and C-type hydroxyls (Cornell et al., 1974; Russell et al., 1974). Recently, thin sections of goethite crystals cut perpendicular to the z axis have suggested that the surface of both multidomainic and monodomainic goethites is almost exclusively formed by (110) and (001) faces
(Schwertmann, 1984; Mann et al., 1985) (Fig. 3.22). Both the (110) and (001) faces reportedly contain only A-type
OH groups (Schwertmann, 1985b; Cornell et al., 1974).
Domain formation decreases with increasing Al substitution and ultimately single domain crystals are formed with well developed but fewer (110) faces. Consequently, the A-type hydroxyls may decrease in total abundance due to reduced surface area but should increase in proportion to B and A bsorbance 0.0070 0.0080 .0040 0 .0050 0 0.0060 0.0030 Fig. 3.21. FTIR subtraction spectra of goethite goethite of spectra subtraction FTIR 3.21.Fig. after S04 adsorption (CIRCLE cell!. (CIRCLE adsorption S04 after SI0) hwn eaiesraehdoy bands hydroxyl surface negative showing CS-I50B) 20 40 60 3840 3640 3440 3240 Wave Wave number (cm-1) 7.98%AI- S-150B
149 001
Aluminum Hydrothermal Substitution T reatment 100
001 010
Fig. 3.22. Effect of hydrothermal treatment ana fll substitution on crystal morphology. 150 151
C-type hydroxyls. Even though Parffit et al. (1976) argue
for preferred displacement of A-type hydroxyls by flouride and phosphate, our data clearly suggest that sulfate adsorption at B and C type sites is preferred and also
"clouds" previously accepted assignments of various OH groups to specific types of faces. Diffuse reflectance spectra like that shown in Fig. 3.7 for the entire series of Al-goethites would help define relative changes in the intensities of the two surface hydroxyl bands as a function of Al substitution, but these data are not currently available.
The spectra of adsorbed sulfate on samples S-OC and
S-OC(200 °C) (untreated and hydrothermally-treated) also indicate the formation of a bidentate bridging complex
(Figs. 3.23 and 3.25). The subtraction spectrum from the surface OH region of sample S-OC shows two well defined negative peaks at 3480 and 3640 cm-i corresponding, once again, to A-type OH, and B and C-type OHs, respectively
(Fig. 3.24). These results are similar to the results observed for sample S-0, indicating that sulfate has undergone ligand exchange reactions with all three types of surface hydroxyls. The corresponding spectra of sample
S-QC(200 °C) (Fig. 3.26) are very similar to the spectra of the Al-substituted sample (S-150B). Only one negative 152
0 u C 0 & fc. o (/> S i <
930 1030 1130 1230 1330 Wave number (cm*1)
Fig. 3.23. FTIR spectra of sulfate adsorbed on synthetic goethite (S-OC). (a) oxide; (b) oxide + S04; Co) b - a. (CIRCLE CELL) Absorbance 0. 5 3 0 .0 -0 -0.0015 -0.0055 0025 2 0 .0 0 0.0005 Fig. 3.24. FTIR subtraction spectra of goethite goethite of spectra subtraction FTIR 3.24. Fig. fe S4 dopin CRL cel I). (CIRCLE adsorption S04 after CS-QC) showing negative surface hydroxyl bands bands hydroxyl surface negative showing CS-QC) 0 0 3 3 Wavenumber (cm-1) 3485 37003 0 0 5 S-OC
0 0 9 3 153 154
S-OC 2 0 0 °C
o u c o JD w O (/> S i <
930 1030 1130 1230 1330 Wave number (cm -1)
Fig. 3.25. FTIR spectra of sulfate adsorbed on synthetic goethite CS-OC (200 C)). (a) oxide; Cb) oxide + S04; (c) b - a. (CIRCLE CELL) Absorbance 0080 8 0 .0 0 0103 0 1 .0 0 6 2 1 .0 0 3 7 1 0 0 0.0150 Fig. 3.26. FTIR subtraction spectra of gosthlts gosthlts of spectra subtraction FTIR 3.26. Fig. ad atr Q asrto (ICE cell). (CIRCLE adsorption SQ4 aftsr bands CS-0CC200*C>) showing negative surface hydroxyl hydroxyl surface negative showing CS-0CC200*C>) 0 0 3 3 ae number Wave 0 0 5 3 0 0 7 3 (cmr (cm-1) *C 0 0 2 - C S-O 0 0 9 3 156
peak is observed at 3642 cm-1 in the subtraction spectrum.
Hydrothermal treatment like Al substitution reduced the
surface area and the multidomainic character of goethite.
Consequently, the proportion of A-type to B and C-type OH
groups may have also been reduced in this sample by
reducing the overall exposed (110) face.
3.5.3 Conclusions from the IR Study
Aluminum substitution and hydrothermal treatment caused a systematic increase in OH bending frequencies of goethite. Higher OH bending frequencies suggest stronger
H- bonding within the goethite structure and, therefore,
improved crystallinity. This is consistent with the
results obtained by XRD and DSC.
Sulfate seems to be adsorbed as a bidentate bridging complex under the conditions employed in this experiment.
Sulfate replaced not only A-type OH groups as suggested by
Parfitt and Smart (1978), but also B and C-type OHs as well. Aluminum substitution and hydrothermal treatment seem to decrease the total abundance of A-type hydroxyls due to reduced surface area. However, the proportion of
A-type to B and C-type hydroxyls should simultaneously increase due to a better development of (110) and (001) faces. CHAPTER IV
ADSORPTION ISOTHERMS
4.1 Surface Charge of Oxides as Related to Sulfate Adsorption
Iron and A1 oxides occur in soils as discrete particles and as coatings on other minerals. In either case these oxides may develop considerable positive surface charge under acid soil conditions. As a result, most anionic species in the soil solution are strongly attracted to oxide surfaces. Surface charge on the Fe and
A1 oxides develops due to unequal adsorption of potential determining ions (H30+ and OH-). Generally, oxygen atoms do not persist at the surface of an oxide once the oxide is exposed to gaseous or liquid water. The oxygens react rapidly with H+ and OH- from dissociated water resulting in a mineral with hydroxylated surfaces. Depending on the pH, a positive, neutral or negative charge will develop.
157 158
Parks and de Bruyn (1962) represented this reaction as
follows:
+3 -3 OH, OH F/ ' p/ p/ l \ 3HjO* l \ 30H ‘ I \ 3H ,0 + O OH, O OH -----► O O +3H,0 1 / 1 / 1/ Fe Fe Fe \ \ \ OH, OH O
At some intermediate pH, the surface will experience a net
charge of zero. This pH is commonly called the zero point
of charge (ZPC). For Fe oxides this pH falls in the range
from 7-10. The higher the ZPC of an oxide the greater is
its potential to develop a net positive surface charge
under normal conditions.
Although the sign of the surface charge plays a major
role in the adsorption of anions, it is not the only
factor influencing the process. Bowden et al. (1977)
suggested that the free energy of adsorption (AGads) at
oxide surfaces is the result of three components, namely
AGcoul, AGint and AGchem. The coulombic component
(AGcoui) represents the pure electrostatic interaction
between a point charge and an electric field. If 159
AGcoul were the sole factor affecting the free energy of
adsorption, an ion would adsorb only on oppositely charged
surfaces, and ions of similar charge would adsorb in
strict proportion to their activities in solution.
The interaction component (AGint) arises from the
specificity in bonding of different ions due to their
charge, size and polarizability. When this component is
significant, adsorption is no longer in strict proportion
to the activity of an ion in solution. The chemical
component (AGchem) arises from the electronic nature of
ions comprising the adsorbate and adsorbent; it is
composed of coordination, van der Waals, and polarization
forces. Thus, an ion can be adsorbed by an uncharged or
even a similarly charged surface provided Gchem is large
enough to overcome the electrical forces of repulsion.
When AGchem is small in a system, ions are adsorbed predominantly on oppositely charged surfaces. This type of adsorption is referred to as non-specific. Nitrate and chloride are examples of ions that are largely retained by
simple electrostatic attractions. In contrast, for
AGchem > 0, some anions are adsorbed regardless of the surface charge or their activity in solution. This type of adsorption involves ligand exchange reactions and is termed "specific." 160
Some anions that are widely recognized as forming inner sphere complexes (specific adsorption) at oxide surfaces include phosphate, selenate, silicate and flouride (Hingston et al., 1967; Parfitt and Smart, 1977;
Hayes et a l . , 1987; Martin and Smart, 1987). A specifically adsorbed anion should affect the surface potential and lower the ZPC of the oxide surface to more acid values. Such a shift in the ZPC has been observed after selenite adsorption on goethite (Hingston et al.,
1972; Sigg and Stumra, 1981) and for phosphate adsorption on an Oxisol (Wann and Uehara, 1978).
There is considerable debate as to the raechanism(s) responsible for sulfate adsorption, with some workers
(Hingston et al., 1967; Parfitt, 1978; Rajan, 1979;
Harrison and Berkheiser, 1982) suggesting ligand exchange, others suggesting the occurrence of both ligand exchange and outer sphere complexes (Sigg and Sturam, 1981; Marsh et al., 1987), and others postulating the formation of outer sphere complexes only (Yates and Healey, 1975).
Determinations of ZPC by titration curves have yielded contradictory results with respect to the shift of the crossover point of the titration. Pyman et al. (1979) found that when sulfate was present in small, constant 161 amounts in varying concentrations of indifferent electrolyte the ZPC was shifted to lower pHs. On the other hand, Sigg and Stumm (1981) observed a shift in the crossover point of the curve to higher pHs in the presence of increasing concentrations of sulfate with the background electrolyte (0.1 M NaClCU) kept constant.
Uehara and Gillman (1981) differentiate between two types of specific adsorption namely "high affinity" and
"low affinity" with the former referring to chemical coordination to the surface metal ion and the latter to adsorption in the Stern layer. They considered phosphate and sulfate as examples of high affinity and low affinity specific adsorption, respectively. Sulfate adsorption in the Stern layer should induce the development of additional positive charge at the surface (Fig. 4.1) and thereby shift the ZPC to higher pHs. Phosphate, on the other hand, becomes part of the surface and transfers its charge to the solid. The additional negative charge will require a lower pH to achieve neutralization of the surface charge (i.e., lower ZPC). Fig. 4.1. Induction of positive charge onto the surface brought about by adsorption of sulfate. 163
4.2 Adsorption Isotherms
Adsorption isotherms have been used extensively to
study sulfate adsorption on soils and oxide surfaces (Chao
et al., 1964; Bornemisaa and Llanos, 1966; Aylmore et al.,
1967; Hasan et al., 1970; Parfitt and Smart, 1978; Rajan,
1978; Fox, 1982; Harrison and Berkheiser, 1982). Chao et al. (1964) and Bornemisza and Llanos (1966) reported
sulfate adsorption isotherms on temperate and tropical soils, respectively, which they modeled with Freundlich- type equations. No indications of adsorption maxima were observed in either study.
Rao and Sridharan (1984) found that sulfate adsorption on a kaolinite from India also followed a typical Freundlich-type behavior within a solution concentration range of 0.25 mmol/L to 50 mmol/L. Based on the amount of OH released with adsorption and corresponding changes in surface charge, they concluded that at low concentrations (0.35 to 0.45 mmol/L) sulfate was adsorbed at positive sites forming a binuclear bridging complex, whereas at higher solution concentrations the sulfate ion formed both monodentate and bidentate complexes. 164
Sulfate adsorption isotherms obeying a Langmuir-type
equation were reported by Aylmore et al. (1967) for
pseudoboehmite, hematite and kaolinite. The Al oxide
showed an adsorption maximum of 420 mmol/Kg, whereas the
adsorption maximum of hematite was considerably lower (67 mmol/Kg). Two distinct regions of sulfate adsorption were
observed for the kaolinites studied, which led the authors to suggest the existence of reaction sites of different binding energy. Muljadi et al. (1966) drew similar conclusions for phosphate adsorption on kaolinite, and
Hasan et al. (1970) also suggested the existence of energetically different adsorption sites on some Hawaiian soils.
Langmuir adsorption isotherms have also been used to describe sulfate adsorption on hydrous alumina (Rajan,
1978) and other Fe oxides (Parfitt and Smart, 1977;
Harrison and Berkheiser, 1982). Sulfate adsorption maxima for these oxides were attained at very low solution concentrations (0.2 - 0.5 ramol/L). The observed adsorption maxima ranged from 85 to 2000 mmol/Kg with
"amorphous" Fe oxides showing the highest adsorption capacities (Table 4.1). 165
Table 4.1 Sulfate adsorption capacity of various oxides.
SCU2- adsorbed Oxide mmol/Kg pH Reference
Hematite 85 3.5 Parfitt and Smart (1978) •i Goethite 125 3.4 M Akaganeite 410 3.2 ii Lepidocrocite 330 3.5 ii Fe(OH)3 600 4.5 Fe(OH)3 2000 5.8 Harrison and Berkheiser (1982) Hydrous Alumina 350 5.0 Rajan (1978) 166
Holford et al. (1974) used a two-term Langmuir equation to describe phosphate adsorption on soils. The two-term equation described phosphate adsorption almost perfectly on 41 soils of varying physical and mineralogical properties; however, the validity of this equation to predict the existence of different adsorption sites should be taken cautiously. Posner and Bowden
(1980) demonstrated that phosphate adsorption isotherms generated by a coordination-chemistry model that assumes a perfectly uniform adsorbing surface could be represented by a sum of two or more traditional Langmuir equations.
The fact that these authors have obtained similar results for isotherms generated with a model which considers only one type of adsorption site throws doubt on the validity of multiple site interpretations.
The use of Freundlich and Langmuir equations to describe ion adsorption has received much negative criticism in recent years (Veith and Sposito, 1977; Posner and Bowden, 1980; Sposito, 1982) because many researchers have attempted to explain adsorption mechanisms based solely on the interpretation of curve shapes and adsorption maxima. Veith and Sposito (1977) demonstrated that the Langmuir equation cannot differentiate between 167 adsorption and secondary precipitation at the adsorbate surface, and concluded that the Langmuir parameters will have no particular chemical meaning unless it can be independently demonstrated that only adsorption occurs.
The general failure of these equations to explain adsorption mechanisms has led to increasing popularity and use in recent years of chemical models of adsorption
(Stumm et al., 1970; Stumm et al., 1976; Stumm et al.,
1980; Sigg and Stumm, 1980; Bowden et al., 1977; Bowden et al., 1980; Goldberg and Sposito, 1984; Goldberg, 1985).
These models differ from the Langmuir and Freundlich adsorption equations in that they take into consideration the reaction of individual ions with the surface and account for electrostatic effects. For example, since ion adsorption decreases the electrostatic potential in the plane of adsorption, each increment of adsorption should make the surface less favorable for adsorption. This overcomes one of the weakest points of the Langmuir equation which assumes adsorption sites of uniform energy.
A description of some of these models has been provided by
Sposito (1983) and Barrow (1985). The models differ primarily in the number of planes that are specified, the specific allocation of ions in those planes and in the type of surface reactions considered. 168
The present study deals with the use of adsorption isotherms to study sulfate adsorption on a series of goethites varying in chemical composition and crystallinity. The adsorption curves should provide useful information to compare the adsorption capacities of the adsorbents under stated conditions. Also, the adsorption data will supplement the characterization and infrared spectroscopy data and help explain differences in sulfate adsorption that may be the result of variations Al substitution or crystallinity.
4.3 Materials and Methods
4.3.1 Preparation of Sulfate Adsorption Isotherms
Eight SO42- solutions in the concentration range of 0 to 2.53 mmol/L were prepared with NaCl (0.01 M) as the background electrolyte. This low sulfate concentration range was selected to minimize the potential for precipitation at the adsorbent surface. Possible Fe and
Al-sulfate precipitation products in the oxides and Nipe 169
clay are jarosite (KFe3(0H)6(S04)2), coquinbite
(Fe2(S04)3.5H20), jurbanite (A10HS04.5H 2O), alunite
(KAl2(0H)e(S04)2) and basalumnite (Al4(0H)10SO4.5H20)
(Charlet, 1986). The pH of the solutions was adjusted to
4.00 ± 0.1 using HC1. Preliminary tests indicated that after sulfate adsorption the pH of the suspension
increased above 7.0. This preliminary adjustment of the pH ensured that after sulfate adsorption the pH would
still be below 7.0 (=6.5) which is more representative of soils with mineralogy dominated by Fe and Al oxides.
Duplicate, 50 - mg samples of oxide or clay were placed in
50 ml nalgene plastic centrifuge tubes and 30 ml of S042- solution were added. The tubes were tightly capped and placed longitudinally in an Eberbach shaker bath with the temperature set at 298 *K. The tubes were then continuously shaken for 12 hr at 120 oscillations per min.
This equilibration period was selected because previous sulfate adsorption studies have indicated that the reaction of sulfate with soils and minerals is rapid, achieving equilibration in a matter of hours (Chao et al.,
1962; Bornemisza and Llanos, 1967; Rajan, 1978). 170
After the shaking period, the oxide or clay suspensions were centrifuged for 25 rain at 10000 rpra
(13776 G) using a high speed Beckman centrifuge
(Model J2-21). The supernatant was then filtered through a 0.2 um filter membrane to remove any suspended mineral particles. The pH values of the solutions were measured using an Orion model 701 A/digital analyzer pH meter with an Orion 8102 combination electrode.
The sulfate that remained in solution after equilibration was determined using a Dionex 2000i Ion
Chromatograph, equipped with a fast separator anion-1 column. A standard curve was prepared in the range from 0 to 0.15 mmol/L. This concentration range was found to yield an excellent linear relationship (r = 0.9999) when plotted against peak area. The highest sample dilution required to fall into this range was 1/25. All reported results were the average of two injections per sample; thus, each point on an isotherm represents four measurements (duplicate samples with duplicate sulfate determinations). Sulfate adsorption (mmol/Kg and mmol/m2) was plotted against sulfate equilibrium solution concentration (mmol/L). 171
The amount of OH released into solution as a result
of sulfate adsorption was calculated from the difference
between the initial suspension pH and the pH at the end of the equilibration period. The values reported
represent the OH activity in solution. Hydroxyl release
(mmol/ra2) was plotted against sulfate equilibrium concentration, and the OH released/SCU2- adsorbed ratio was plotted against mole % Al substitution and temperature of the hydrothermal treatment.
4.4 Results and Discussion
4.4.1 Al-substituted Goethites
Sulfate adsorption isotherms for the Al - substituted goethites on both a mass and surface area basis are presented in Figs. 4.2 and 4.3. The curves show a distinct s-shape which become less pronounced as Al substitution increases. For all goethites, the sulfate adsorption capacity, expressed on a mass basis was very similar at low equilibrium concentrations (0.20 - 1.20 mmol SO4-2/L) in solution; however, at high concentrations
(2.00 -2.4 mmol SO4-2) the adsorption capacity decreased SQ4 adsorbed Cmmol/Kg oxide) 200 150 100 50 .0 .50 10 15 2.00 2 3.00 0 5 2. 0 0 . 2 1.50 1.00 0 5 0. 0.00 Fig. substitutedgoethites on a 0 o S-0 No fil No S-0 o A S-10 S-10 A o o * * 5 7 . 4 0 8 - S 0 3 - S 4.2. 1.85 1.85 Mole Sulfateadsorption isotherms on fll- EquiIibrium concentration Cmmol/L) concentration EquiIibrium 7 0 . 1 Mole Mole X X X fil fil fil o S-150 fl S-150 o + S-120 + S-120 ■ S-150 B B ■ S-150 mass basis. 4 5 . 6 8 9 . 7 2 9 . 7 Mole Mole Mole Mole Mole Mole X fil X X fil fil
173
fll fll Rl X X X 2.50 3.00 2.00 + S-120 6 .5 4 Mole ■ ■ S-150B 7.98 Mole o S-150R 7.92 Mole 1.50 Rl Rl Rl X X X onconoontrat Cnvnol/L.3 I 1.00
I Ibrlum I Equl 0.50 □ □ S-30 * S-80 1.85Mole 4.75Mole a S-10 1.07Mole o S-0 No fll - - substituted goethites on a surface area basis. Fig. 4.3. Sulfate adsorption isotherms on fll- 0.00 0 0 . 0.00 2 3.00 - 4.00 5.00
C-01 X C2 W/|ouiuj) P0q-IO8PD frQS 174 with Al substitution (Fig. 4.2, Table 4.2). Sample S-0
(no Al) showed an adsorption capacity of 161 mmol/kg, whereas S-150B (7.98 mole % Al) showed an adsorption capacity of 121 mmol/kg which represents 75 % of the sulfate adsorbed by the unsubstituted sample at the highest equilibrium concentration. The decrease in adsorption was not directly proportional to Al substitution. Sample S-120 (6.54 mole % Al) showed the lowest sulfate adsorption capacity, but contains less Al than samples S-150A (7.92 mole % Al) and S-150B (7.98 mole
% Al). The lower sulfate adsorption capacity of sample
S-120 may be an effect of surface area. This sample has the lowest surface area (28 m2/g) among all samples in this series (see Table 2.4, page 63).
On a surface area basis, the differences in adsorption at high sulfate solution concentrations are negligible (Fig. 4.3, Table 4.2). This suggests that the higher adsorption capacity observed on a mass basis for the unsubstituted sample (S-0) and the low Al-substituted samples (S-10, S-30, and S-80) at high equilibrium sulfate concentrations is mostly a surface area effect. Aluminum substitution decreased the total surface area by reducing the porosity and raultidomainic character of the goethites. Table 4.2. Sulfate adsorbed by Al - substituted goethites at the highest S04“2 equilibrium concentration.
Sample mmol/Kg mmol/m2 x 10“3*
S-0 161 3.83 S-10 153 3.73 S-30 141 3.52 S-80 152 4.23 S-120 99 3.54 S-150A 129 4.29 S-150B 121 3.79
* Average value corresponding to the last two points of the curve. 176
The formation of a more uniform surface with Al
substitution should render the oxides less reactive; however, sulfate adsorption increases with Al substitution at intermediate sulfate equilibrium concentrations
(0.50 - 1.50 mmol S04-2/L). This effect of Al substitution can be seen clearly in Fig. 4.4 where only three curves are plotted. A point in the middle of this region (1.2 mmol S04_2/L) was selected to observe the variation in sulfate adsorption with Al substitution.
This quantity is designated as SO4-2 and is presented in
Table 4.3. The SO4-2 of adsorption increases from 0.51 to 1.04 mmol/m2 x 10"3 as Al substitution increases from
1.07 mole % (S-10) to 7.98 mole % (S-150B). The same general effect was observed when the areas comprised by the curve of the unsubstituted sample (S-0) and the various Al-substituted samples were measured using a digitizing tablet. The area values presented in Table
4.3 are the average of four measurements per sample. The area of the SO4-2 region increased from 2.792 x 10-* to
1.221 x IO-3 ra2 as Al substitution increased from 1.07 to
7.98 mole %. 177 3 . 0 0 2.50 2.00 1.50 Rl Z Rl Rl oonoentratlon Cmmol/LD Z 1.00 Mol© Mole 7 . 9 8 EqulIIbrlum 0.50 Sulfate adsorption isotherms on fll-substItuted fll-substItuted on isotherms adsorption Sulfate s-jsoB m o S-8Q 4.75 o S-0 No Rl No S-0 o ------4.4. 1.00 2.00 3. 0 0 4 . 0 0 5 . 0 0 1 intermediate sulfate equilibrium concentration. equilibrium sulfate intermediate 0 —01 X C2UJ/IOUJUO peqJ08pD frOS at adsorption sulfate Fig. in increase the showing goethites 178
Table 4.3 Variation in S042- adsorbed with aluminum substitution.
AS04 2 - adsorbed Sample Hole % Al mmol/m2 x 10-3* Area (m2)**
S-10 1.07 0.51 2.792 x 10-4 S-30 1.65 0.19 1.981 x 10-4 S-80 4.75 0.89 9.084 x 10-4 S-120 6.54 0.97 1.162 X 10-3 S-150A 7.92 1.15 1.358 X 10-3 S-150B 7.98 1.04 1.221 X 10-3
* Difference in adsorption between S-0 (no Al) and the indicated sample at 1.20 mmol/L SO4-2 equilibrium concentration.
** Area comprised between S-0 curve and the indicated Al - substituted sample curve. 179
The increase in sulfate adsorption in the mid region of the isotherms as Al substitution increases suggests a greater affinity for sulfate in these samples. Apparently,
Al substitution has created a limited number of "high" affinity adsorption sites. These sites become "saturated" with increasing sulfate concentration so that at higher sulfate levels no significant differences are observed between samples. Evidence supporting the higher affinity for sulfate displayed by the Al-substituted samples comes from the heats of formation (AHf) of Fe2(S04)3 and
Al2(SO4)3. These compounds have a AHj of -2576.93 and
-3440.84 KJ/mol, respectively (Robie et al., 1978). The more negative the AHf the greater the tendency of the compound to form. Given the more negative value for Al, it seems reasonable that reaction of sulfate with the oxide surface would be enhanced by the substitution of Al for Fe at adsorption sites.
Aluminum substitution may have also enhanced the formation of bidentate bridging sulfate complexes by reducing the separation between surface hydroxyls. When only Fe atoms are present in goethite, the separation of these hydroxyls is approximately 0.276 nm (Schwertmann,
1985b). The smaller size of the Al3+ ion will result in a 180
shorter distance between hydroxyls which will be closer
to the 0.24 nm 0-0 separation in the sulfate molecule
(Parfitt and Russell, 1977), thus facilitating the
formation of bidentate bridging complexes.
4.4.2 Hydrothermally-treated Goethites and the Nipe Clay
Hydrothermal treatment of goethites resulted in a
sharp and systematic decrease in sulfate adsorption on a mass basis (Fig. 4.5). Sulfate adsorption decreased from
107 to 42 mmol/Kg as the temperature of the hydrothermal treatment was increased from 50 *C to 200 °C (Table 4.4).
Like Al substitution, hydrothermal treatment decreased the porosity and multidomainic character of the crystals resulting in a more uniform surface. Unlike the Al- substituted samples, which yielded isotherms with a fairly linear shape, the hydrothermal isotherms maintain a distinct s-shape. The s-shaped isotherms suggest that the hydrothermally treated goethites have less affinity for sulfate at low solution concentrations than do the Al- substituted ones (Sposito, 1984). 200 o S-OC a S-OC C50 °CD
CD □ S-OC C100 °CD 150 o S-OC Cl50 °CD ■ S-OC C200 °C)
8 100
•Q
0.50 1.00 1.50 2.00 2.50 3.00 Equilibrium conc©ntration Cmmol/L) Fig. 4.5. Sulfate adsorption isotherms on hydrothermal Iy- treated goethites on a mass basis. Table 4.4. Sulfate adsorbed by hydrothermally - treated goethites at the highest S04-2 equilibrium concentration.
Sample mmol/Kg mmol/m2 x 10-3*
S-OC 96 2.52 S-OC (50 1*C) 107 3.03 S-OC (100 *C) 89 2.53 S-OC (150 •C) 68 2.96 S-OC (200 *C) 42 2.59
* Average value corresponding to the last two points of the curve. 183
The sharp decrease in sulfate adsorption with increasing temperature of hydrothermal treatment is not seen when the data are expressed on a surface area basis
(Fig. 4.6), which indicates that the observed differences are mostly due to a surface area effect (see Table 2.7).
The increase in sulfate adsorption observed for intermediate sulfate equilibrium concentrations in the Al- substituted goethites is also not observed in this series of goethites. Sample S-0C(200 ®C) shows a small increase in this region but it is not comparable to the increase observed for the A1 substituted samples.
The sulfate adsorption isotherms of Nipe clay on both a mass and surface area basis are presented in Figs. 4.7 and 4.8 and are compared to some of the synthetic goethites. This clay, which contains considerable amounts of goethite, also shows an s-shaped isotherm. The adsorption capacity at the highest equilibrium sulfate level is larger on a mass basis than that of S-0C(200 °C) but smaller than that of the Al-substituted sample
(S-150B). In contrast, when the data are presented on a surface area basis, the Nipe clay curve lies below all the others. This lower adsorption capacity may be the result 184
2 . 5 0 3 . 0 0 ■ S-OC C200 °CD C200 S-OC ■ o S-OC Cl50 °CD Cl50 S-OC o 2.00 1.50 1.00
EqulIibrlum concGntration Cmmol/LD concGntration EqulIibrlum 0 . 5 0 S-OC C50 C50 °CD S-OC a o S-OC o - □ S-OC Cl00 °CD Cl00 S-OC □ - 1.00 2.00 4 . 0 0 3 . 0 0 5.00 treated goethites on a surface area basis. C-Ol X C2W/IOUW) peqjospD frQS Fig. 4.6. Sulfate adsorption isotherms on hydrothermal Iy- S04 adsorbed (mmol/Kg) 200 100 150 .0 .0 .0 .0 .0 .0 3.00 2.50 2.00 1.50 1.00 0.50 0.00 £ Q Fig. otie adNp ca o am9 basis. ma98 a on clay Nipe and goethites ------o o S-OCC200 °C) S0 o Rl No S-0 o A S-150B 7 .9 8 Mole Mole 8 .9 7 S-150B A Np Clay Nipe ■ 4.7. 4.7. 1 lbrum cnetain Cmmol/L) concentration m riu ilib u q E ------Sulfate adsorption isotherms on synthetic synthetic on isotherms adsorption Sulfate o-
v 1 ______X fil
I ______I ______
I ______185 S04 adsorbed Cmmol/m2) x 10-3 5.00 3.00 2.00 4.00 0.00 1.00 .0 .0 .0 .0 .0 2.50 2.00 1.50 1.00 0.50 0.00 Fig. goethites and Nipe clay on a surface area basis. area surface a on clay Nipe and goethites o o S-0 No fii o o S-OC °CD(200 a a S-150B 7.98 Mole ■ Nipe Clay . 8 . 4 Sulfate adsorption isotherms on synthetic synthetic on isotherms adsorption Sulfate EquiIIbrium concentration Cmmol/LD concentration EquiIIbrium
X fii
3.00 186 187
of surface site contamination by other anions including
sulfate, or to the presence of less reactive mineral
surfaces like kaolinite.
4.4.3 Sulfate Adsorption and OH Release
The suspension pH after equilibration increased with
increasing concentration of sulfate in solution and thus
with sulfate adsorption. Plots of sulfate equilibrium
concentration vs. amount of OH- released are presented in
Pigs. 4.9 and 4.10 for all adsorbents. Figure 4.9 shows
that the maximum amount of OH" released decreases from
0.41 x 10"6 to 0.23 x 10"6. mmol/m2 as A1 substitution
increases from 1.07 to 7.98 mole % ’, however, sample S-0
(No Al) deviates from the observed trend. Due to the similar properties of sample S-0 and S-10, the curve of the former should lie above or close to the curve of the latter. The behavior of this sample cannot be explained.
Hydrothermal treatment also resulted in a decrease of OH" release but the effect was more marked than for the Al- substituted samples (Fig. 4.10). The maximum amount of
OH- released decreased from 0.58 x 10-6 (S-OC) to 0.03 x
10-6 (S-0C(200 *C)) mmol/m2 . □H released Cmmol/m2) x 10-0 0 . 00 0.00 ^ Fig. ocnrto (lsbtttd goethites}. (Rl-substItuted concentration * S-10 Cl.07 Cl.07 2 HID Mole S-10 * ni) CNo S-0 o * * □ 5 7 . 4 C 0 8 - S 0 3 - S . 9 . 4 EqulIIbrlum concentration Cmmol concentration EqulIIbrlum .0 .0 .0 .0 2.50 2.00 1.50 1.00 0.50 Cl.65 Cl.65 Mole OH released vs. sulfate equilibrium equilibrium sulfate vs. released OH Mole Z Z FII J HI 5 ■ ■ OS-I50FI OS-I50FI S-120 + 8 9 . 7 C B 0 5 1 - S 4 5 . 6 C 2 0 . 7 C D L / 4 0 S Mole Mole Mole 2 fllJ2 Mole Mole Mole Z Z RIJ RIJ flli 3.00 OH released Cmmol/m2) x 10-0 0.70 0.80 0.30 0.60 0.10 0.20 0.40 Fig. 4.10. OH released vs. sulfate equilibrium equilibrium goethites sulfate Iy-treated vs. (Hydrothermal released OH concentration 4.10. Fig. n Np clay). Nipe and o o S-OC a S-OC C50 °C) o o S-OCC150 S-OCC50°C)°C) EquiIibrium concentrationCmmol 0.50 1.00 □ S-OC C100°C) 1.50
2.00 * Nipe Clay ■ S-OCC200 °C) S04/L) .03.00 2.50
189 190
The curve for Nipe clay lies above all other samples, but most of the hydroxyls driven into solution do so only in the presence of NaCl so that very little increase in OH release can be attributed to sulfate adsorption. The small amount of OH released with added sulfate is in agreement with the fact that the Nipe clay had the lowest sulfate adsorption capacity among all adsorbents on a surface area basis. The different behavior of the clay can be attributed to the heterogeneity of the sample, which includes a mixture of goethite, hematite, gibbsite and kaolinite.
Another way to evaluate the relationship between displaced surface hydroxyls and adsorbed sulfate is to examine the ratio of OH released/sulfate adsorbed (Fig.
4.11 - 4.12). Aluminum substitution in goethite results in a decrease of the OH released/S04 adsorbed ratio
(Fig. 4.11) in agreement with the previously suggested decrease in total number of surface hydroxyls (inferred from FTIR data) and the observed decrease in surface area induced by A1 substitution. The low Al-substituted samples show more variability in OH/S04 ratio with increasing sulfate equilibrium concentration than do the high Al-substituted ones. The 0H/S04 ratios tend to 0H/S04 x 10-4 4.00 3.00 2.00 0.00 1.00 Fig. 4.11. OH released/SQ4 adsorbed ratio as a function function a as ratio adsorbed released/SQ4 OH 4.11. Fig. f fliof substitution. . 10 . 30 . 50 . 70 . 9.0 8.0 7.0 6.0 5.0 4.0 3.0 2.0 1.0 0.0 CO.15} o ) 0 3 . 2 C ■ Sulfate equilibrium concentration Cmmol/L) Cmmol/L) concentration equilibrium Sulfate Mole a + Cl.52) + 2
Rl substitution ) 5 3 . O C C1.92) o □ □
) 3 7 . O C 191 0H/S04 x 10-4 4.00 3.00 2.00 0.00 Fig. 4.12. OH released/S04 adsorbed ratio as a function function a as ratio adsorbed treatment. released/S04 OH hydrothermal of 4.12. Fig. 50 - o CO.16} CO.16} o l2) 1.563 C + ■ Cl.21) * ) 9 3 . 2 C Sulfate equilibrium concentration (mmol/L) (mmol/L) concentration equilibrium Sulfate 0 150 100 Temperature C °C) Temperature a
} 5 7 . O C □ ) 6 3 . O C o (2.02) o 200 193 converge for samples S-120 (6.54 mole % Al), S-150A (7.92 mole % Al) and S-150B (7.98 mole % Al) suggesting the presence of a more uniform surface with less variable adsorption sites in these samples.
Hydrothermal treatment resulted in a sharp decrease in the 0H/S04 ratio (Fig. 4.12). This again agrees with
FTIR results which suggest a decrease in the number of surface OHs with hydrothermal treatment. Like Al substitution, hydrothermal treatment reduced the surface area and the domainic character of the crystals. Such changes resulted in the formation of more uniform crystals where the nature of surface hydroxyls undergoing exchange reactions should be more similar. The fact that the
0H/S04 ratios at all sulfate equilibrium concentrations converge as the temperature of the hydrothermal treatment increases (samples S-OC (150 °C) and S-OC (200 °C)) supports this concept.
4.4.4 Conclusions
Variation in surface area was the dominant factor influencing sulfate adsorption capacity of both the hydrothermal and Al-substituted goethites. However, Al 194 substitution clearly enhanced sulfate adsorption in the
0.5 - 1.5 rnraol/L sulfate equilibrium range, perhaps due to the creation of a limited number of "high" affinity
* adsorption sites. Similar results were not observed with the hydrothermal goethites. The increase in pH as a result of sulfate adsorption suggests that ligand exhange reactions have occurred at the surface of the oxides; however, no conclusion about the type of bonding reaction can be drawn purely from the adsorption data. The decrease in the 0H-/S04 ratio with increasing Al substitution and hydrothermal treatment suggests a decrease in the overall amount of surface hydroxyls as previously indicated by FTIR. The lower surface area, more monodomainic character and improved crystallinity brought about by Al substitution and hydrothermal treatment resulted in the formation of more uniform crystals with more uniform adsorption sites and was reflected in decreased variability of the 0H-/S04 ratios at high Al substitution and high hydrothermal treatment.
The sulfate adsorption characteristics of the Nipe soil clay were comparable to those of the synthetic goethites. The lower total adsorption per unit surface 195 area may reflect contamination of adsorption sites by other anions, including sulfate, or the presence of less reactive mineral phases like kaolinite. CHAPTER V
SUMMARY
Goethite (a-FeOOH) is the most abundant iron oxide in
soils and normally occurs in solid solution with aluminum.
Aluminum may substitute for over 30 mole % of the
structural iron in natural specimens. Goethite is thought
to play a major role in anion adsorption and has been the
subject of many phosphate adsorption studies. In most
such studies, the effect of systematic variations in
adsorbent properties were not evaluated.
The present study was conducted to test the
hypothesis that Al substutition in goethite should alter
the physicochemical properties of the oxide and therefore
influence anion adsorption. Sulfate was selected as the adsorbate of interest because of controversy surrounding
its mode of bonding to oxide surfaces.
A series of Al-substituted goethites (0-9 mole % Al) were successfully prepared by coprecipitating Fe(N03)3-
A1(N03)3 solutions with KOH; however, Al substitution was found to systematically decrease the surface area of the oxides. To separate the effects of surface area and Al substitution on sulfate adsorption, a second series of goethites were produced by hydrothermally treating
196 197
subsaroples of unsubstituted goethites. Aluminum
substitution and hydrothermal treatment increased the
cyrstallinity and dehydroxylation temperature of the
synthetic goethites while reducing surface area and the
presence of multiple domains or intergrowths in the
goethite crystals.
Fourier transform infrared spectroscopy studies
indicated that sulfate was specifically adsorbed at the
oxide surface. The splitting of the V3 sulfate band into
three components suggested the formation of a bidentate
bridging complex with two Fe atoms. Sulfate replaced not
only A-type OH groups, but also B and C-type OHs as well.
Even though Parfitt et al. (1976) argue for preferred
displacement of A-type hydroxyls by fluoride and
phosphate, our data clearly suggested that sulfate
adsorption at B and C-type sites is preferred and also
"clouds” previously accepted assignments of various OH
groups to specific types of faces.
Aluminum substitution and hydrothermal treatment seem
to decrease the total abundance of A-type hdroxyls due to
reduced surface area. However, the proportion of A-type
to B and C-type hydroxyls should simultaneously increase due to a better development of the (110) and (001) faces where the former predominates. 198
Sulfate adsorption isotherms indicated that variation in surface area was the dominant factor influencing sulfate adsorption in both goethite series. However, Al substitution clearly enhanced sulfate adsorption in the
0.5 - 1.5 mmol/L sulfate equilibrium range, perhaps due to the creation of a limited number of "high" affinity adsorption sites. The lower surface area, more monodomainic character and improved crystallinity brought about by Al substitution and hydrothermal treatment presumably also resulted in the formation of crystals with more uniform adsorption sites. Evidence of this effect was obtained in the form of decreased variability of
0H/S04 ratios at high Al substitution and hydrothermal treatment.
The sulfate adsorption capacity of a clay extracted from an Oxisol (Nipe - Typic Acrorthox) was compared to the adsorption capacities of the synthetic goethites. The adsorption characteristics of the clay were comparable to those of the synthetic goethites; however, the total adsorption per unit surface area of the clay was lower.
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SULFATE ADSORPTION ISOTHERMS DATA
215 216
Table A.I. Sulfate adsorption isotherms data (mmol/Kg) for samples S-0, S-10, S-30 and S-80.
Sulfate adsorbed (romol/Kg)
Replication 1 Replication 2 Sulfate equilibrium Injection Injection
Sample nmol/L) 1 2 1 2 Averai
S-0 0.14 61.0 61.0 49.0 53.0 56.0 0.33 67.0 68.0 58.0 55.0 62.0 0.73 65.0 69.0 48.0 55.0 59.0 1.18 55.0 67.0 72.0 84.0 70.0 1.53 81.0 90.0 92.0 97.0 90.0 1.89 137.0 156.0 166.0 164.0 156.0 2.25 161.0 179.0 155.0 168.0 166.0
S-10 0.16 ____ 43.0 45.0 44.0 0.34 67.0 68.0 44.0 47.0 57.0 0.74 55.0 56.0 56.0 56.0 56.0 1.14 83.0 89.0 92.0 93.0 89.0 1.51 107.0 109.0 92.0 98.0 102.0 1.90 130.0 160.0 142.0 171.0 151.0 2.26 181.0 149.0 147.0 141.0 155.0
S-30 0.15 46.0 51.0 41.0 47.0 46.0 0.36 44.0 44.0 49.0 52.0 47.0 0.73 55.0 57.0 64.0 62.0 60.0 1.16 87.0 53.0 68.0 89.0 74.0 1.53 92.0 101.0 89.0 100.0 96.0 1.93 149.0 129.0 113.0 141.0 133.0 2.28 142.0 145.0 146.0 163.0 149.0
S-80 0.15 36.0 42.0 54.0 54.0 47.0 0.36 44.0 50.0 43.0 48.0 46.0 0.71 67.0 76.0 66.0 57.0 67.0 1.14 87.0 83.0 96.0 102.0 92.0 1.50 102.0 120.0 115.0 108.0 111.0 1.89 140.0 153.0 158.0 156.0 152.0 2.26 142.0 141.0 158.0 169.0 153.0 217
Table A.2. Sulfate adsorption isotherms data (mmol/Kg) for samples S-120, S-150A, and S-150B.
Sulfate adsorbed (mmol/Kg)
Replication 1 Replication 2 Sulfate ______equilibrium Injection Injection concentration ______Sample (mmol/L) 1 2 1 2 Average
S-120 0.16 38.0 44.0 44.0 45.0 43.0 0.36 41.0 40.0 46.0 51.0 45.0 0.71 55.0 63.0 68.0 76.0 66.0 1.17 73.0 81.0 57.0 84.0 74.0 1.55 69.0 75.0 83.0 70.0 74.0 1.99 70.0 115.0 102.0 106.0 98.0 2.36 87.0 123.0 75.0 116.0 100.0
S-150A 0.15 45.0 48.0 42.0 47.0 46.0 0.36 45.0 46.0 43.0 46.0 45.0 0.73 54.0 64.0 56.0 57.0 58.0 1.15 86.0 70.0 97.0 83.0 84.0 1.51 114.0 90.0 102.0 77.0 96.0 1.93 105.0 122.0 142.0 145.0 129.0 2.31 114.0 131.0 134.0 135.0 129.0
S-150B 0.16 44.0 43.0 47.0 48.0 46.0 0.36 43.0 42.0 50.0 48.0 46.0 0.74 43.0 57.0 43.0 49.0 48.0 1.14 80.0 86.0 86.0 94.0 87.0 1.50 97.0 112.0 95.0 109.0 103.0 1.94 111.0 141.0 129.0 149.0 133.0 2.34 112.0 100.0 119.0 109.0 110.0 218
Table A.3. Sulfate adsorption isotherms data (mmol/m^ x 10-6) for samples S-0, S-10, S-30, and S-80.
Sulfate adsorbed (mmol/m2 x 10~3)
Replication 1 Replication 2 Sulfate ______equilibrium Injection Injection concentration ______Sample (mmol/L) 1 2 1 2 Average
S-0 0.14 1.46 1.45 1.16 1.25 1.33 0.33 1.59 1.62 1.39 1.32 1.48 0.73 1.55 1.65 1.13 1.31 1.41 1.18 1.31 1.60 1.71 2.00 1.66 1.53 1.92 2.15 2.19 2.31 2.14 1.89 3.25 3.72 3.95 3.89 3.70 2.25 3.84 4.25 3.70 4.00 3.95
S-10 0.16 ------— _ —— 1.04 1.08 1.06 0.34 1.64 1.66 1.06 1.14 1.38 0.74 1.34 1.36 1.37 1.35 1.36 1.14 2.02 2.18 2.24 2.26 2.17 1.51 2.61 2.66 2.23 2.39 2.47 1.90 3.18 3.90 3.45 4.17 3.68 2.26 4.42 3.64 3.59 3.43 3.77
S-30 0.15 1.19 1.27 1.03 1.16 1.16 0.36 1.09 1.09 1.22 1.30 1.18 0.73 1.38 1.43 1.60 1.56 1.49 1.16 2.16 1.32 1.70 2.21 1.85 1.53 2.30 2.52 2.22 2.50 2.39 1.93 3.71 3.24 2.82 3.52 3.32 2.28 3.55 3.63 3.65 4.06 3.72
S-80 0.15 1.00 1.16 1.49 1.50 1.29 0.36 1.22 1.38 1.19 1.32 1.28 0.71 1.85 2.12 1.83 1.59 1.85 1.14 2.41 2.29 2.67 2.84 2.55 1.50 2.82 3.33 3.20 3.00 3.09 1.89 3.89 4.26 4.39 4.33 4.22 2.26 3.95 3.92 4.38 4.69 4.24 219
Table A.4. Sulfate adsorption Isotherms data (mmol/m2 x 10“®) for samples S-120, S-150A, and S-150B.
Sulfate adsorbed (mmol/m2 x 10“2)
Replication 1 Replication 2 Sulfate ______equilibrium Injection Injection concentration ______Sample (mmol/L) 1 2 1 2 Average
S-120 0.16 1.37 1.59 1.58 1.61 1.54 0.36 1.45 1.43 1.66 1.82 1.59 0.71 1.97 2.25 2.43 2.73 2.35 1.17 2.61 2.91 2.02 3.00 2.63 1.55 2.48 2.66 2.95 2.49 2.65 1.99 2.50 4.09 3.65 3.79 3.51 2.36 3.09 4.39 2.66 4.15 3.57
S-150A 0.15 1.49 1.61 1.40 1.57 1.52 0.36 1.51 1.55 1.42 1.53 1.50 0.73 1.78 2.12 1.88 1.90 1.92 1.15 2.87 2.33 3.24 2.77 2.81 1.51 3.80 3.01 3.38 2.58 3.19 1.93 3.51 4.07 4.72 4.84 4.29 2.31 3.79 4.37 4.45 4.51 4.28
S-150B 0.16 1.35 1.34 1.46 1.48 1.41 0.36 1.33 1.30 1.55 1.49 1.42 0.74 1.33 1.77 1.35 1.53 1.50 1.14 2.51 2.67 2.70 2.93 2.70 1.50 3.03 3.50 2.96 3.40 3.22 1.94 3.47 4.41 4.03 4.66 4.14 2.34 3.50 3.13 3.70 3.40 3.43 220
Table A.5. Sulfate adsorption isotherms data (mmol/Kg) for samples S-OC, S-0C(50 ’C), S-0C(100 #C), and S-0C(150 *C).
Sulfate adsorbed (mmol/Kg)
Replication 1 Replication 2 Sulfate ______equilibrium Injection Injection concentration ______Sample (mmol/L) 1 2 1 2 Average
S-OC 0.15 51.0 51.0 50.0 50.0 51.0 0.35 51.0 51.0 49.0 49.0 50.0 0.74 38.0 45.0 51.0 48.0 46.0 1.18 68.0 65.0 63.0 65.0 65.0 1.53 69.0 90.0 105.0 114.0 95.0 2.00 96.0 98.0 89.0 93.0 94.0 2.36 106.0 92.0 94.0 97.0 97.0
S-0C(50 #C) 0.15 48.0 49.0 49.0 50.0 49.0 0.35 52.0 52.0 49.0 50.0 51.0 0.74 60.0 49.0 41.0 47.0 49.0 1.21 61.0 67.0 42.0 60.0 58.0 1.51 98.0 101.0 68.0 81.0 87.0 1.96 91.0 112.0 121.0 126.0 113.0 2.35 98.0 101.0 94.0 106.0 100.0
S-OC(100 *C) 0.15 45.0 47.0 49.0 49.0 48.0 0.36 44.0 43.0 44.0 47.0 45.0 0.74 42.0 53.0 47.0 52.0 49.0 1.22 35.0 38.0 63.0 60.0 49.0 1.56 43.0 63.0 76.0 86.0 67.0 2.00 87.0 94.0 91.0 90.0 91.0 2.39 72.0 83.0 93.0 96.0 86.0
S-0C(150 *C) 0.16 42.0 43.0 39.0 41.0 41.0 0.37 33.0 36.0 41.0 43.0 38.0 0.75 47.0 47.0 43.0 41.0 45.0 1.21 44.0 46.0 43.0 43.0 44.0 1.59 60.0 74.0 35.0 46.0 54.0 2.04 68.0 68.0 56.0 80.0 68.0 2.41 73.0 46.0 72.0 81.0 68.0 221
Table A. 6. Sulfate adsorption isotherms data (mmol/Kg) for samples S-OC, S-OC(200 °C) and Nipe clay.
Sulfate adsorbed (mmol/Kg)
Replication 1 Replication 2 Sulfate equilibrium Injection Injection concentration Sample (mmol/L) 2 Average
S-0C(200* C) 0.18 28.0 29.0 28.0 30.0 29.0 0.39 24.0 25.0 37.0 31.0 29.0 0.78 26.0 31.0 19.0 29.0 26.0 1.23 32.0 38.0 34.0 40.0 36.0 1.60 45.0 49.0 40.0 46.0 45.0 2.08 34.0 57.0 33.0 53.0 44.0 2.45 28.0 53.0 34.0 40.0 39.0
Nipe clay 0.18 32.0 29.0 32.0 31.0 31.0 0.39 30.0 28.0 25.0 27.0 28.0 0.78 29.0 26.0 29.0 29.0 28.0 1.21 49.0 51.0 65.0 49.0 54.0 1.55 75.0 72.0 81.0 82.0 78.0 2.01 80.0 105.0 72.0 92.0 87.0 2.41 79.0 75.0 64.0 49.0 67.0 222
Table A.7. Sulfate adsorption isotherms data (mmol/m* x 10-3) for samples S-OC, S-0C(50 *C), S-0C(100 *C) and S-0C(150 °C).
Sulfate jadsorbed (mmol/m^ x 10- »)
Replication 1 Replication 2 Sulfate equilibrium Injection Injection concentration Sample (mmol/L) 1 2 1 2 Average
S-OC 0.15 1.35 1.35 1.31 1.31 1.33 0.35 1.34 1.34 1.29 1.29 1.31 0.74 0.99 1.19 1.34 1.22 1.20 1.18 1.79 1.71 1.65 1.70 1.71 1.53 1.81 2.38 2.76 3.00 2.49 2.00 2.52 2.59 2.35 2.46 2.48 2.36 2.79 2.43 2.48 2.54 2.56
S-0C(50 *C) 0.15 1.38 1.40 1.41 1.43 1.41 0.35 1.49 1.50 1.40 1.42 1.45 0.74 1.71 1.40 1.17 1.34 1.41 1.21 1.73 1.90 1.19 1.72 1.64 1.51 2.81 2.87 1.94 2.31 2.49 1.96 2.59 3.21 3.46 3.60 3.21 2.35 2.81 2.88 2.69 3.03 2.85
S-0C(100 *C) 0.15 1.29 1.35 1.40 1.41 1.36 0.36 1.26 1.23 1.24 1.33 1.27 0.74 1.20 1.52 1.34 1.48 1.39 1.22 0.99 1.10 1.80 1.77 1.40 1.56 1.23 1.80 2.16 2.45 1.91 2.00 2.50 2.69 2.60 2.57 2.59 2.39 2.06 2.38 2.66 2.75 2.46
S-0C(150 °C) 0.16 1.82 1.89 1.71 1.79 1.80 0.37 1.43 1.57 1.78 1.89 1.67 0.75 2.06 2.06 1.86 1.78 1.94 1.21 1.92 1.99 1.87 1.88 1.91 1.59 2.62 3.22 1.52 1.98 2.34 2.04 2.96 2.96 2.43 3.47 2.96 2.41 3.16 1.99 3.13 3.52 2.95 223
Table A.8. Sulfate adsorption isotherms data (mmol/m2 x 10-®) for samples S-0C(200 *C) and Nipe clay.
Sulfate adsorbed (mmol/m2 x 10~®)
Replication 1 Replication 2 Sulfate equilibrium Injection Injection concentration Sample (mmol/L) Average
S-0C(200 *C) 0.18 1.72 1.83 1.74 1.84 1.78 0.39 1.51 1.58 2.31 1.92 1.83 0.78 1.60 1.90 1.16 1.83 1.62 1.23 1.99 2.39 2.10 2.50 2.25 1.60 2.83 3.09 2.52 2.89 2.83 2.08 2.10 3.55 2.09 3.29 2.76 2.45 1.73 3.33 2.14 2.51 2.42
Nipe clay 0.18 0.58 0.53 0.57 0.57 0.56 0.39 0.55 0.51 0.46 0.50 0.51 0.78 0.52 0.47 0.52 0.52 0.51 1.21 0.89 0.92 1.18 0.89 0.97 1.55 1.36 1.31 1.48 1.49 1.41 2.01 1.45 1.91 1.31 1.66 1.59 2.41 1.44 1.37 1.17 0.88 1.21 224
Table A. 9. OH released (mmol/m2 x 1 0 -6 ) for samples S-0, S-l 0, S-30 and S-80.
OH released mmol/m2 x 10-6
Replications Sulfate equilibrium ______Sample concentration mrool/L 1 2 Average
S-0 0 0.002 0.002 0.002 0.14 0.032 0.030 0.031 0.33 0.035 0.047 0.041 0.73 0.044 0.071 0.058 1.18 0.053 0.079 0.066 1.53 0.066 0.075 0.071 1.89 0.085 0.076 0.081 2.25 0.072 0.056 0.064
S-10 0 0.045 0.033 0.039 0.16 0.225 0.232 0.228 0.34 0.371 0.304 0.337 0.74 0.270 0.351 0.310 1.14 0.459 0.375 0.417 1.51 0.487 0.334 0.411 1.90 0.591 0.369 0.480 2.26 0.372 0.455 0.414
S-30 0 0.019 0.020 0.020 0.15 0.170 0.217 0.194 0.36 0.238 0.150 0.194 0.73 0.294 0.176 0.235 1.16 0.347 0.318 0.333 1.53 0.331 0.383 0.357 1.93 0.336 0.412 0.374 2.28 0.346 0.356 0.351
S-80 0 0.023 0.017 0.020 0.15 0.172 0.175 0.174 0.36 0.252 0.261 0.257 0.71 0.221 0.222 0.222 1.14 0.278 0.210 0.244 1.50 0.343 0.269 0.306 1.89 0.286 0.288 0.287 2.26 0.326 0.335 0.331 225
Table A .10. OH released (raraol/ra2 x 10"6) for samples S-120, 5-150A and S-150B.
OH released romol/m2 x 10“6
Replications Sulfate equilibrium Sample concentration mmol/L Average
S-120 0 0.016 0.015 0.160 0.16 0.158 0.222 0.190 0.36 0.223 0.253 0.238 0.71 0.349 0.257 0.303 1.17 0.371 0.344 0.358 1.56 0.453 0.362 0.408 1.99 0.460 0.380 0.420 2.36 0.346 0.333 0.340
S-150A 0 0.013 0.012 0.013 0.15 0.113 0.110 0.112 0.36 0.146 0.156 0.151 0.73 0.202 0.224 0.213 1.15 0.218 0.225 0.222 1.51 0.226 0.205 0.216 1.93 0.312 0.303 0.308 2.31 0.252 0.269 0.261
S-150B 0 0.000 0.000 0.000 0.16 0.001 0.001 0.001 0.36 0.001 0.001 0.001 0.74 0.236 0.002 0.119 1.14 0.305 0.246 0.276 1.50 0.363 0.003 0.183 1.94 0.441 0.003 0.222 2.34 0.003 0.004 0.004 226
Table A. 11. OH released (ramol/m2 x 10-6) for samples S-OC, S- 0C(50 *C), S-0C(100 °C) and S-0C(150 °C).
OH released mmol/m2 x 10“6
Replications Sulfate equilibrium Sample concentration mmol/L Average
S-OC 0.15 0.037 0.029 0.033 0.35 0.253 0.414 0.334 0.74 0.360 0.412 0.386 1.18 0.420 0.362 0.391 1.53 0.374 0.392 0.386 2.00 0.491 0.558 0.525 2.36 0.554 0.599 0.577
S-0C(50 0 0.019 0.019 0.019 0.15 0.068 0.139 0.104 0.35 0.389 0.145 0.267 0.74 0.438 0.311 0.375 1.21 0.349 0.357 0.353 1.51 0.286 0.435 0.361 1.96 0.413 0.470 0.442 2.35 0.416 0.393 0.405
S-0C(100 0 0.015 0.014 0.015 0.15 0.135 0.178 0.157 0.36 0.193 0.196 0.195 0.74 0.257 0.268 0.263 1.22 0.333 0.450 0.392 1.56 0.405 0.356 0.381 2.00 0.357 0.426 0.392 2.39 0.362 0.487 0.425
S-0C(150 0 0.008 0.006 0.007 0.16 0.050 0.058 0.054 0.37 0.043 0.094 0.069 0.75 0.089 0.075 0.082 1.21 0.136 0.091 0.114 1.59 0.127 0.139 0.133 2.04 0.126 0.117 0.122 2.41 0.161 0.126 0.144 227
Table A.12. OH released (mmol/m2 x 10-8 ) for samples S-0C(200 *C) and Nipe clay.
OH released mmol/m2 x 10-6
Replications
Sample concentration mmol/L 1 2 Average
S-0C(200 *C) 0 0.003 0.004 0.004 0.18 0.009 0.014 0.012 0.39 0.017 0.012 0.015 0.78 0.019 0.021 0.020 1.23 0.023 0.021 0.022 1.60 0.024 0.024 0.024 2.08 0.026 0.025 0.026
Nipe clay 0 0.458 0.536 0.497 0.18 0.540 0.543 0.542 0.39 0.549 0.581 0.565 0.78 0.549 0.581 0.565 1.21 0.562 0.651 0.607 1.55 0.712 0.703 0.708 2.01 0.844 0.497 0.671 2.41 0.404 0.674 0.539