PHOSPHATE ADSORPTION BY MIXED AND REDUCED IRON PHASES IN

STATIC AND DYNAMIC SYSTEMS

A THESIS

SUBMITTED TO THE DEPARTMENT OF GEOLOGY

OF STANFORD UNIVERSITY

IN PARTIAL FULFILLMENT OF THE REQUIREMENTS

FOR THE DEGREE OF

MASTER OF SCIENCE

Theresa M. Barber

August 2002

ii

I certify that I have read this dissertation and that, in my opinion, it is fully adequate in scope and quality as a thesis for the degree of Master of Science.

Scott Fendorf, Principal Advisor

I certify that I have read this dissertation and that, in my opinion, it is fully adequate in scope and quality as a thesis for the degree of Master of Science.

Adina Paytan

iii

ABSTRACT

Phosphorus (P) availability is an important biogeochemical element and is often

the limiting nutrient in freshwater and estuarine ecosystems. It is a macronutrient and is a

critical component in DNA, RNA, ATP, and phospholipids. The prominent correlation

between phosphorus retention and iron has led to numerous examinations of P adsorption

capacities and mechanisms onto iron oxides and hydroxides in aerobic environments,

most commonly hematite and goethite.

Many studies have discounted iron oxides as playing a significant role in P

retention under anoxic conditions, suggesting P will be released to solution upon the

reductive dissolution of the iron (III) oxides or phosphates; however, there are conflicting

sentiments regarding P retention in anaerobic environments. The goal of this study was

to elucidate the binding capacity of phases common to anaerobic environments for P

retention. The specific objectives of this study were (1) to determine P sorption

capacities of siderite, magnetite, and sulfate green rust and compare retention to that of

lepidocrocite, goethite-coated sand, ferrihydrite-coated sand, and ferrihydrite slurry, (2)

to determine the influence of pH and ionic strength on retention capacity, and (3) to

examine phosphate retention under flow through heterogeneous and homogenous

mixtures of iron phases.

The results of the static experiments indicate ferrous minerals have phosphate

adsorption capacities that are much greater than those of the ferric phases examined.

However, the ferrous minerals all had much lower binding coefficients and were more

sensitive to changes in pH. The high capacity for phosphate of the ferrous minerals,

combined with the relatively low affinity and sensitivity to pH suggests that the ferrous

minerals investigated may provide appreciable, although transitory phosphorus sinks in

some reductomorphic environments.

The hydrodynamic experiments indicated that the pH of a system will determine

the reaction time necessary for a particular sorbent to become saturated with respect to

PO4, but if the reaction time is adequate pH will not diminish the capacity of the sorbent.

Additionally, the influent PO4 concentration will control the adsorption capacities

exhibited by an iron oxide sorbent. The presence of elements such as Ca and Fe(II) may

promote the precipitation of phosphate minerals, which increases the apparent adsorption

iv

capacity of a sorbent. Reductive dissolution of ferric oxides leading the precipitation

secondary ferrous minerals create complicated mixtures of sorbents, each of which will

exhibit different sorption capacities and responses to changes in pH. Sorption of ferrous

iron may passivate reactive sites on the primary and secondary phases that could

otherwise sorb phosphate. This work has demonstrated that phosphate retention in

dynamic systems is dependant upon many interrelated factors, each of which must be

considered in order to fully understand and predict P retention, even in these relatively

simple geochemical systems.

v

ACKNOWLEDGMENTS

I would like to thank my advisor, Dr. Scott Fendorf, for his guidance and patience

throughout the last two years. His excitement for science is contagious and kept me

motivated even during the times I felt like giving up. Colleen Hansel, Dr. Shawn Benner,

and Dr. Bruce Wielinga were all instrumental in the completion of this work, providing

guidance in the lab, as well as thought-provoking discussions about experimental results.

I would also like to thank Dr. Adina Paytan for her time and input during the final stages

of this project. Dr. Guangchao Li assisted with many of the chemical analyses and

offered valuable advice regarding experimental design. I would like to thank several

other graduate students, Colin Doyle, Chris Oze, T.J. Kiczenski, Megan Young, Kristen

Averyt, and Chris Van de Ven, for providing counsel, insight, and encouragement

throughout the past two years. I also wish to thank Tim Himmer for his patience and

support throughout the past year. Finally, I would like to thank the Department of

Geological and Environmental Sciences and the McGee Foundation for their financial

support.

vi

TABLE OF CONTENTS

Abstract.....................................................................................................................iii

Acknowledgements ..................................................................................................v

Table of Contents .....................................................................................................vi

List of Tables ............................................................................................................viii

List of Figures...........................................................................................................x

Chapter 1: Literature Review

Importance of phosphate in the natural environment..................................... 1

Global cycling and phosphate chemistry ....................................................... 4

Iron: mineral occurrence, chemistry, and importance toward phosphate...... 9

Adsorption and adsorption models: a general overview................................12

Phosphate sorption onto iron oxides ..............................................................15

Need and research objectives .........................................................................22

References......................................................................................................26

Chapter 2: Phosphate Adsorption by Mixed and Reduced Iron Minerals Under

Static Conditions

Introduction

Importance and cycling of phosphate.................................................33

Phosphorus retention processes .........................................................34

Need and research objectives .............................................................36

Materials and Methods

Mineral preparation............................................................................39

Mineral characterization ....................................................................40

Phosphate adsorption experiments.....................................................41

Release of sulfate from green rust......................................................42

Results

Adsorption capacities & isotherms ....................................................43

Influence of pH and ionic strength.....................................................45

Release of sulfate from green rust......................................................45

Discussion ......................................................................................................53

vii

Conclusions....................................................................................................59

Appendix: Isotherm data collected during sorption experiments ..................60

References......................................................................................................71

Chapter 3: Phosphate Adsorption by Heterogeneous Mixtures of Ferrous and Ferric

Oxides Under Hydrodynamic Conditions

Introduction

Importance and cycling of phosphate in the natural environment .....75

Need and research objectives .............................................................78

Materials and Methods

Ferrihydrite-coated sand preparation .................................................79

General column set-up and sampling ................................................79

Aerobic columns (ferrihydrite-coated sand only) ..............................80

Anaerobic columns (multiple iron phases coating sand) ...................82

Static adsorption experiments (on sand) ............................................83

Surface Area Analysis........................................................................83

Results

Aerobic columns ................................................................................85

Anaerobic columns ............................................................................90

Discussion ......................................................................................................103

Conclusions....................................................................................................107

References......................................................................................................109

viii

LIST OF TABLES Chapter 1: Literature Review

Table 1.1. Summary of previously determined phosphate adsorption capacities........................................................................................................17

Chapter 2: Phosphate Adsorption by Mixed and Reduced Iron Minerals Under Static Conditions

Table 2.1. Maximum observed phosphate adsorption on a mass basis of

the mineral adsorbent (µmols PO4/g mineral)................................................44

Table 2.2. Surface area (SA) and SA normalized adsorption maxima for phosphate on various iron oxides...................................................................44

Table 2.3. Summary of adsorption capacities (b-values) and KL-values for minerals conforming to the Langmuir isotherm.............................................44

Table 2.4. Average adsorption (µmols PO4/g) at high and low ionic strength...........................................................................................................46

Table 2.5. Concentration data and saturation indices for phosphate solutions after 2 d reaction with geologic magnetite, geologic siderite, and sulfate green rust. ....................................................................................55 Table 2A.1. Summary of data used to establish adsorption isotherm of PO4 by naturally occurring (geologic origin) magnetite. ...........................61 Table 2A.2. Summary of data used to establish adsorption isotherm of PO4 by sulfate green rust ...........................................................................62 Table 2A.3. Summary of data used to establish adsorption isotherm of PO4 by naturally occurring (geologic) siderite ..........................................63 Table 2A.4. Summary of data used to establish adsorption isotherm of PO4 by ferrihydrite (as a slurry). ................................................................64 Table 2A.5. Summary of data used to establish adsorption isotherm of PO4 by ferrihydrite-coated sand.................................................................65 Table 2A.6. Summary of data used to establish adsorption isotherm of PO4 by goethite-coated sand. .....................................................................66 Table 2A.7. Summary of data used to establish adsorption isotherm of PO4 by lepidocrocite. .................................................................................67

ix

Table 2A.8. Summary of data used to establish adsorption isotherm of PO4 by synthetic siderite............................................................................68 Table 2A.9. Summary of data used to establish adsorption isotherm of PO4 by naturally sulfate green rust. ...........................................................69

Chapter 3: Phosphate Adsorption by Heterogeneous Mixtures of Ferrous and Ferric

Oxides Under Hydrodynamic Conditions Table 3.1. Average feed solution compositions used during biotic column experiments. ...................................................................................................84 Table 3.2: Distribution of iron (mole percent) in abiotic, anaerobic column upon experiment termination. ...........................................................91 Table 3.3. Mineralogical composition (in terms of mole percent iron) with distance downfield in GW and BC columns..........................................98

x

LIST OF FIGURES Chapter 1: Literature Review Figure 1.1. Depiction of the phosphorus cycle.............................................6 Figure 1.2. Speciation of phosphate as a function of pH. ............................8 Figure 1.3. Diagram of surface complexes. ..................................................13 Figure 1.4. Example of an experimental adsorption envelope for goethite......................................................................................21 Chapter 2: Phosphate Adsorption by Mixed and Reduced Iron Minerals Under

Static Conditions

Figure 2.1. Phosphate adsorption curves for minerals fit to a Freundlich isotherm (a) geologic magnetite, (b) sulfate green rust and (c) geologic siderite .................................................................................47 Figure 2.2. Adsorption isotherms of ferric (hydr)oxides all fit to the Langmuir isotherm (a) ferrihydrite, (b) ferrihydrite-coated sand, (c) goethite-coated sand, and (d) lepidocrocite. .............................................48 Figure 2.3. Phosphate adsorption curve fit to a Langmuir isotherm for (a) synthetic siderite and (b) synthetic magnetite. ..................................49 Figure 2.4. Influence of pH on phosphate adsorption to various iron oxides......................................................................................................50 Figure 2.5. Sulfate released from green rust with increasing ionic strength (as NaCl). ................................................................................51 Figure 2.6. Phosphate adsorption and sulfate release with increasing phosphate concentration. Also shown is phosphate sorbed and sulfate released into 100 µM phosphate solution when no NaCl is present ..............52

Chapter 3: Phosphate Adsorption by Heterogeneous Mixtures of Ferrous and Ferric Oxides Under Hydrodynamic Conditions

Figure 3.1. Schematic of column apparatus used to evaluate PO4

3- transport through packed mineral beds. .........................................................81 Figure 3.2. Average influent and effluent P concentration and pH for aerobic columns PB1 and PB2 (a) and aerobic well-buffered column (b). .....................................................................................................86

xi

Figure 3.3. Side-port solution P concentration and pH through time for aerobic columns PB1 (a-b) and PB2 (c-d)................................................87 Figure 3.4. Elemental concentration (mM) of P and Fe on the solid-phase after 50 d (a-b) and 21 d (c) of reaction for PB1 and PB2 (a–b) and the well-buffered column (c)..................................................88 Figure 3.5. Aqueous P concentrations along the flow-path of the well-buffered, aerobic column. ......................................................................89 Figure 3.6. Schematics of ripened anoxic, abiotic column before (a) and after (b) introduction of phosphate solution...........................................92 Figure 3.7. Influent and effluent P concentration and pH data (a) for abiotic, anoxic experiment and side-port solution P concentration and pH through time.. ....................................................................................93 Figure 3.8. Side-port solution Fe concentration through time for the anaerobic, abiotic experiment. .................................................................94 Figure 3.9. Elemental concentration of the solid-phase after 11 d of reaction for the anoxic, abiotic column .....................................................95 Figure 3.10. Mineralogical composition as determined by EXAFS analysis upon termination of anoxic, abiotic column. ...................................96 Figure 3.11. Influent and effluent P and pH data for the GW (a) and BC (b) biotic columns .............................................................................99 Figure 3.12. Total (6 M HCl) and partial (0.5 M HCl) extraction data from GW (a) and BC (b) anaerobic, biotic columns......................................100 Figure 3.13. Fe and PO4 concentrations (mM) in pore water in GW (a) and BC (b) biotic columns at experiment termination...................................101 Figure 3.14. Mineralogical composition (in terms of mole % iron) of packed mineral beds within GW (a) and BC (b) columns upon experiment termination. .................................................................................102

1

CHAPTER 1

Literature Review I. Importance of phosphate in the natural environment Phosphate is an essential component of ATP (adenosine tri-phosphate), DNA

(deoxyribonucleic acid), RNA (ribonucleic acid), and phospholipids as a result of its

unique chemical properties. Phosphoric acid has three ionizable groups; within DNA

molecules two groups link the nucleotides while the third inhibits hydrolysis and

maintains the structural integrity of the molecule. Phosphate also serves as an

intermediary in metabolism and during energy transfer involving ATP. Available pools

of phosphate and nitrogen control the expanse and activity of the biosphere on a global

level. The critical role occupied by phosphate in combination with relatively low

availability leads to very rapid biological cycling and a high level of phosphate

conservation in undisturbed ecosystems (Schlesinger, 1991). In many forests, phosphate

input through atmospheric deposition is greater than loss to runoff (Berner and Berner,

1996). Berner and Berner (1996) also refer to phosphate at being the most ‘biologically

active’ of all the macronutrients, meaning the ratio of phosphate stored in the biomass to

that lost to runoff is very high. However, in areas cleared for agricultural production,

phosphate loss may be significant due to increased soil erosion and removal during the

harvesting process (Brady and Weil, 1999). The high amount of phosphate lost from

agricultural ecosystems is also due in part to the large quantities of soil amendments used

in those areas (Berner and Berner, 1996).

From the agricultural standpoint, phosphate is an important factor in crop vitality.

Phosphate availability is often a problem due to three primary factors. First, the

phosphate content of soil is relatively low, usually falling within the range of 0.001

mg/kg (infertile) to 1 mg/kg (heavily fertilized). Secondly, phosphate compounds and

complexes within soils are often insoluble and not readily available for biological uptake.

Finally, the complexes previously mentioned fix (irreversibly adsorb) phosphorus added

as a soil amendment. Phosphate deficiency can impede numerous developmental

processes including photosynthesis, flowering, fruiting, maturation, nitrogen fixation, and

root growth. The combination of strong retention mechanisms within soils and the

2

importance of phosphate to plant health has led to the over-application of phosphate in

many agricultural areas (Brady and Weil, 1999).

Freshwater and estuarine ecosystems are often nutrient limited and are

particularly sensitive to increases in phosphate (Reddy et al., 1999; Manahan, 1994).

Phosphate input to lakes is normally limited to rainfall and most of the phosphate in a

given lake is sequestered in the biomass (plankton). Cycling of phosphate is very rapid in

these systems, but a small amount is lost from the biomass to the sediments. Reductive

dissolution of ferric iron phases within the sediments may release phosphate back into the

water column but this is typically a very small input (Schlesinger, 1991). Soil and

sediment erosion can contribute appreciable quantities of particulate phosphate to aquatic

systems, especially since erosion will preferentially remove clay-sized particles and

organic matter, both of which are strong phosphate sorbents. Dissolved phosphate may

also be introduced to aqueous systems by surface runoff from drainage areas with

unincorporated fertilizers on the land surface. While phosphate is not directly toxic to

aquatic organisms, eutrophication and toxic algal blooms, such as pfisteria, do cause fish

kills. Eutrophication occurs when a limiting nutrient is added to a water body. Decay of

resulting algal blooms consumes oxygen and prevents light infiltration. As the algae

begin to die and decompose the biological oxygen demand spikes, resulting in anoxia

(Brady and Weil, 1999). Phosphate is not often a problem in groundwater due to strong

sorption by Fe and Al oxides common in aquifer materials (Sparks, 1995).

Recent studies in the Everglades have documented that small additions of

phosphate, due to increased runoff from agriculture, appear to be altering the ecology of

the area. Prior to urbanization and intensified agriculture, rainfall was the predominant

source of phosphate to the system and the Everglades were oligotrophic (DeBusk et al.,

1994). Minor increases of phosphate, from ~0.07 µM/L to ~16 µM/L, shift the

environment to a more eutrophic state resulting in the replacement of cyanobacterial-

diatom populations by a filamentous green algae, the advance of cattails into areas that

were once strictly sawgrass, and endemic periphyton replaced by algal species normally

found in more eutrophic areas (DeBusk et al., 2001; Noe et al., 2001; DeBusk et al.,

1994). DeBusk et al. (1994) also suggested that increased phosphate loading might result

in even larger scale changes in tropic state, including possible increased biomass

3

production, changes in species composition, and a decrease of species diversity. In

addition to changes in vegetation, Wright and Reddy (2001) noted that phosphate was a

factor in peat accumulation rates and changes in enzyme activity within the Everglades.

Finally, the high binding strength of phosphate to ferric oxides has implications

on the fate of other contaminants in the natural environment. Phosphate competes with

and displaces numerous other anions including AsO43-, SO3

2-, MoO42-, CrO4

2-, SO42-,

SeO42-, Br-, I-, and Cl- (Geelhoed et al., 1997; McBride, 1994). Ryden et al. (1977a)

examined phosphate sorption and competition by other inorganic anions and found that

SO42-, SeO4

2-, Cl-, and NO2- had no effect on phosphate sorption; however, H2AsO4,

HSeO3-, and MoO4

2- all depressed phosphate adsorption. These authors also found that

H2AsO4- was more competitive for sites when added at the same time as phosphate.

Arsenic migration, in particular, has received a great deal of attention in recent

years due to the implications on human health. Application of phosphate fertilizers has

been correlated with increased As mobility (Darland and Inskeep, 1997; Peryea and

Kammereck, 1997) in soils, and in some cases phosphate and hydroxide have been used

specifically to remove As from soils (Jackson and Miller, 2000). Su and Puls (2001)

studied As removal by zero-valent Fe and observed that phosphate caused a substantial

decrease in As retention, with As(III) retention inhibited more at lower pH and As(V)

retention hindered more at higher pH. In soil columns, phosphate addition increased As

mobility but did not yield total As desorption even when the amount of phosphate added

was greater than the sorption capacity of the soil (Darland and Inskeep, 1997). Reynolds

et al. (1999) found that phosphate enhanced As transport in well-aerated soils but had

very little effect on As mobility in soils that had undergone cycles of flooding and

subsequent aeration. This phenomenon was attributed to the formation of more

amorphous iron oxides in the flooded soils, leading to increased sorption capacity for

both arsenic and phosphate.

Several researchers have examined the competition of arsenic and phosphate for

sorption sites in more controlled, single sorbent systems. O’Reilly et al. (2001)

demonstrated arsenic release from goethite upon addition of phosphate to the surrounding

solution. Hongshao and Stanforth (2001) also used goethite as a sorbent and found that

the order in which arsenic and phosphate were added to solution made a significant

4

difference in the adsorption and desorption characteristics. They observed roughly equal

adsorption when arsenic and phosphate were added concurrently, resulting in an overall

surface coverage greater than that observed during single anion addition. When the ions

were added successively the amount of exchange was found to be dependant upon how

long the first anion had been allowed to react. A positive correlation between amount of

exchange and equilibration time was observed. Based on these observations, a two-stage

reaction was proposed, beginning with very rapid surface complexation followed by

slower precipitation, with the second step being the most influential on exchange. Jain

and Loeppert (2000) examined the influence of phosphate on arsenate and arsenite

adsorption by ferrihydrite. The presence of phosphate reduced the amount of As sorbed

by ferrihydrite, but this occurrence was conditional and dependent upon the pH and

phosphate concentration. Phosphate had a greater effect on arsenate sorption at higher

pH and on arsenite at lower pH. They also suggested some surface sites had a greater

affinity for arsenite than phosphate as well as a “moderate preference for arsenate”.

II. Global cycling and phosphate chemistry

Weathering of geologic materials is the ultimate source of phosphate to the

biosphere; however, the amount of phosphate released is a small fraction of that required

for the current levels of biological production. This deficiency results in phosphate being

highly conserved and very rapidly cycled within the biosphere (Filippelli and Souch,

1999; Berner and Berner, 1996). Vegetation, periphyton, plankton, and microbes all act

as biotic phosphorus sinks. However, biotic cycling is not without P loss to abiotic

mechanisms, including adsorption by soils and sediments, mineral precipitation, and

sedimentation. Phosphate lost from the biomass is quickly complexed by iron and

aluminum oxides or sorbed by clay minerals (Berner and Berner, 1996). Dissolution of

phosphate minerals and complexes may be expedited by plant root exudates or

mycorrhizae, but the major environmental transformations of phosphate are not

biologically or chemically driven. The largest flux of phosphate in terms of geochemical

cycling is fluvial transport to the oceans. Phosphate cycling within the water column is

quite rapid, but constant losses of small amounts from the biomass to the sediments and

5

soils have resulted in oceanic sediments being the largest phosphate reservoir on earth

(Schlesinger, 1991). Figure 1.1 is a flowchart of the global phosphorus cycle.

Phosphorus almost always occurs as the oxyanion PO43-, the behavior of which

may be used as a proxy for arsenate (O’Reilly et al., 2001; Reynolds et al., 1999). The

acid-dissociation constants for phosphate are listed below (Stumm and Morgan, 1981).

[H3PO4] ↔ [H+] + [H2PO4-] Ka = 10-2.2 [1.1]

[H2PO4-] ↔ [H+] + [HPO4

2-] Ka = 10-7.2 [1.2]

[HPO42-] ↔ [H+] + [PO4

3-] Ka = 10-12.3 [1.3]

One should note that the pKa values for the arsenate anion are 3.60, 7.25, and 12.52

(Benjamin, 2002). At circumneutral pH, both HPO42- and H2PO4

- are important anions,

with HPO42- being more prevalent in slightly alkaline conditions and H2PO4

- dominating

in slightly acidic environments (Figure 1.2).

Phosphorus adsorption by soils and lake sediments has often been correlated to

amorphous iron, aluminum, and manganese oxides and hydroxides, as well as calcium

and magnesium (Brady and Weil, 1999; Reddy et al., 1999; Torrent, 1997; Golterman

1995; Olila and Reddy, 1995; Blanchar and Frazier, 1991; Williams et al., 1971). Once

sequestered within soils or sediments, phosphate falls into one of four categories:

phosphate minerals, labile or non-labile sorbed phases, or organic phosphate. Non-labile

and labile soil phosphate pools fall into calcium bound phosphate and iron or aluminum

bound phosphate. Non-labile phosphorus is usually incorporated within the matrices of

Fe and Al minerals and are usually quite insoluble. Labile, or more bioavailable,

phosphate is sorbed to the surface of silica, calcium carbonate, or other soil minerals. In

acidic soils, iron and aluminum phosphate minerals and complexes are dominant, while

calcium phosphates are the main controlling phases in alkaline environments. Several

common soil phosphate minerals are apatite [Ca5(PO4)3(OH,F,Cl)], hydroxyapatite

[Ca5(PO4)3OH], fluorapatite [Ca5(PO4)3F], octocalcium phosphate [Ca8H2(PO4)2·5H2O],

strengite [FePO4·2H2O], vivianite [Fe3(PO4)2·8H2O], variscite [AlPO4·2H2O], and

wavellite [Al3(PO4)2(OH)3·5H2O] (Brady and Weil, 1999; Reddy et al., 1999).

6

Xenobiotic Organisms

Soluble inorganic P (HPO42-,

H2PO4-), polyphosphates

Assimilation by organisms Precipitation

Biological P: ADP, ATP, DNA, nucleic acids, phospholipids

Biodegradation

P in sediments: biological, organic, inorganic

Insoluble P (inorganic): Ca & Fe phosphates

Fertilizer runoff, wastewater, etc

Dissolution

Figure 1.1. Depiction of the phosphorus cycle (taken from Manahan, 1994).

7

Organic phosphate is typically categorized into three groups. Inositol phosphates

make up 10 to 50 percent of the total organic phosphate pool in soils. These compounds

are phosphate esters of sugar-like molecules and are usually very stable in alkaline

conditions and are likely to react with higher molecular weight compounds. The second

group consists of phosphate contained within nucleic acids, which may be sorbed by

silicate clays and humic acids. Phospholipids, molecules critical to the cytoplasmic

membranes of cells, comprise the final group. Most organic phosphate is associated with

the fulvic acid fraction of soil organic matter (Brady and Weil, 1999).

Organic phosphate must be converted back to an inorganic form via

mineralization in order to be usable by higher plants. Humus and organic matter will

decompose to release HPO42- or H2PO4

-, depending on pH. Soil conditions and the

presence or absence of carbon will dictate whether the released phosphate is rapidly

reincorporated into the biomass, precipitates as a mineral, or retained as a sorbed phase

(Brady and Weil, 1999; Manahan, 1994). The cycling of phosphate within soils can be

simply expressed (Brady and Weil, 1999):

1 2 Organic P H2PO4 Fe, Al, Ca phosphates 1. Microbial catalyzed reaction in both directions 2. Reaction requires presence of Al, Fe, Ca, or organic matter. This

phosphate is fixed and not readily available.

Seasonally flooded soils pose unique conditions on phosphate dynamics due to

the changing oxidation-reduction conditions. Periodic flooding may occur naturally or be

the result of rice production or be a means to control weeds, insects and soil subsidence

(Diaz et al., 1993). Episodic flooding generally leads to amplified P release due to

changes in iron mineralogy (i.e., reductive transformation) (Pant and Reddy, 2001a;

Quang and Dufey, 1997; Diaz et al., 1993). Pant and Reddy (2001a) in fact proposed that

water table fluctuations within wetlands destabilize sorbents and lead to phosphate

release.

8

pH

0 2 4 6 8 10 12 14

Frac

tion

of S

peci

es

0.0

0.2

0.4

0.6

0.8

1.0H2PO4

-HPO42- PO43-H3PO4

Figure 1.2: Speciation of phosphate as a function of pH.

9

Riparian zones, areas adjacent to rivers where the vegetation is predominantly

influenced by the presence of water, can take many different forms, from narrow bands of

land along headwater streams to well-developed floodplains along larger rivers (Carlyle

and Hill, 2001). Phosphate retention within riparian zones is particularly important

because these systems provide a crucial buffer between surface run-off and phosphate-

limited freshwater systems. Riparian zones have long been recognized as natural filters,

removing excess nutrients, heavy metals, and other contaminants from non-point sources

(Bridgham et al., 2001; Carlyle and Hill, 2001; Fisher and Reddy, 2001; Wright and

Reddy, 2001). Riparian zones, do however, exhibit differing abilities to sequester

phosphate, a property of which is thought to be related to the amount of dissolved oxygen

present in the system; higher dissolved oxygen levels have been associated with higher

levels of retained phosphate (Carlyle and Hill, 2001). Within wetlands three biological

pools predominate: plants, microbes and organic matter (peat) are partially responsible

for phosphate accumulation. While retention by plants and microbes is seasonal, peat

accumulation in wetlands can store large amounts of phosphate for long periods of time

(Bridgham et al., 2001; Carlyle and Hill, 2001; Fisher and Reddy, 2001). Bridgham et al.

(2001) and Carlyle and Hill (2001) both cite the main phosphate removal processes as

sedimentation of particulates and adsorption, with the greatest phosphate removal taking

place within the first few meters of the wetland.

III. Iron: mineral occurrence, chemistry, and importance toward phosphate

Iron oxides account for the majority of metallic oxides in soils and are found in

most climatic regions. These oxides are particularly important because they are often

contributing, if not controlling, factors in the fate and transport of nutrients and

contaminants (Huang and Wang, 1997). The various iron oxides usually have high

surface areas and chemically reactive functional groups. Surface complexation not only

results in the retention of contaminants and nutrients but also alters the redox properties

of the iron phase (in particular, reductive dissolution) (Stumm and Sulzberger, 1992).

The primary source of iron to the aqueous environment is the weathering of

igneous and metamorphic minerals. Ferrous (II) iron is present in these minerals and

oxidation to ferric (III) iron can weaken the silicate structure and enhance the weathering

10

process. Once ferric iron is released into solution, it rapidly hydrolyzes and subsequently

precipitates as an oxide (Huang and Wang, 1997; Langmuir, 1997). The iron cycle is

closely tied to phosphorus, sulfur, oxygen, carbon, and heavy metals, as well as light and

biological organisms. Due to the highly reactive surfaces and large surface areas of most

iron oxides and hydroxides, surface reactions are a controlling factor of iron flux (Stumm

and Sulzberger, 1992).

The mobility of iron in aqueous systems is dependant on the redox potential and

pH of the environment. Under reducing conditions, ferrous iron predominates and is

generally relatively soluble up to near pH 8. However, if sulfide is present, ferrous iron

will be removed as FeS. In addition to low redox conditions, low pH (pH<4) will

enhance dissolution of most ferric iron phases. Complexation of ferric iron with organic

ligands (humic and fulvic acids) may also lead to mobile complexes. Ferric iron forms

strong, inner-sphere complexes with ligands such as OH-, PO43-, and F- and is likely to be

present as complexes in most aquatic environments (Langmuir, 1997).

The common ferric oxides and hydroxides are goethite, hematite, lepidocrocite,

and ferrihydrite (or hydrous ferric oxide). The arrangement of the anions is the main

control on the crystal structure and interconversion of the oxides. An octahedron,

Fe[O(H)]6, is the basic structural unit in these oxides, which are linked by sharing edge

surfaces, faces, or corners (Cornell and Schwertmann, 1996). The surface areas and

reactivities of iron oxides are highly variable and dependent upon crystallinity and

porosity. The porosity of iron oxides may be caused by structural defects, aggregation, or

dehydroxylation and contributes to the internal surface area. Common ranges of surface

areas are: goethite 8-80 m2/g, lepidocrocite 15-100 m2/g, ferrihydrite 100-400 m2/g,

hematite 2-90 m2/g, and magnetite 4-100 m2/g (Cornell and Schwertmann, 1996; Deng

and Stumm, 1994).

Goethite (α-FeOOH) is the most widespread Fe oxide in soils and sediments in

temperate climates. Hematite (α-Fe2O3) is a common weathering and oxidation product

of ferrous silicates and may form as a result of ferrihydrite conversion. Like goethite,

hematite is quite stable and is often an end member of transformations and is more

common than goethite in older sediments and desert soils, since higher temperatures and

arid conditions favor hematite formation. Organic matter will also dictate whether

11

hematite or goethite forms, particularly in surface soils, due to its ability to complex Fe.

Complexation of organic matter with Fe will suppress hematite formation since it

prevents the solubility product of ferrihydrite from being exceeded, but not that of

goethite (Langmuir, 1997; Cornell and Schwertmann, 1996).

Lepidocrocite (γ-FeOOH) is a small component of noncalcareous soils in humid

and temperate regions. This oxyhydroxide is particularly apt to form in soils having a

fluctuating water table (reductomorphic), since formation is favored by rapid oxidation at

low pH, low total iron, low temperature and the absence of aqueous ferric iron (Huang

and Wang, 1997). Ferrihydrite or hydrous ferric oxide (Fe(OH)3·nH2O) is very prevalent

in surface conditions and precipitates from the rapid hydrolysis of Fe(III) or oxidation of

Fe(II). Ferrihydrite is poorly ordered and is likely to convert to hematite in warm areas

and goethite in cooler areas. The degree of ordering of both synthetic and naturally

occurring ferrihydrite is variable, producing differing XRD patterns (2-line and 6-line).

The 2-line and 6-line ferrihydrites form under different conditions, and 2-line ferrihydrite

is not thought to convert to the 6-line form (Langmuir, 1997; Cornell and Schwertmann,

1996).

Mixed and low valent iron minerals, such as magnetite, siderite and green rust, are

also common in natural environments. Magnetite (Fe3O4) contains both ferrous and

ferric iron and has been found in soils as lithogenic magnetite in the mineral fraction and

as authigenic magnetite. Formation of authigenic magnetite occurs abiotically but is

facilitated by bacteria (Cornell and Schwertmann, 1996; Zachara et al., 2002). Siderite

(FeCO3) is the third most common carbonate mineral in ocean sediments (Schwertmann

and Taylor, 1977) is a by-product of microbial respiration of iron (Zachara et al., 2002);

authigenic siderite has also been observed in anoxic sediments (Boughriet et al., 1997;

Moore and Reddy, 1994).

Green rust is a blue-green Fe2+-Fe3+ hydroxide with a pyroaurite structure and

may contain Cl-, SO42-, or CO3

2- in the interlayer. The structure of the mineral consists of

FeII(OH)6 octahedra in which a fraction of the FeII is replaced with FeIII. The FeII/FeIII

ratio can vary from 0.8 to 3.6, creating a positive layer charged which is balanced by the

interlayer anions (Cornell and Schwertmann, 1996). Green rusts were first noticed as

products of steel corrosion and were later observed in lake sediments and soils (Boughriet

12

et al., 1997). They can form as byproducts of dissimilatory iron reduction by bacteria and

as intermediates in the formation of iron oxides and hydroxides. The most common

interlayer anions observed in nature are SO42- and CO3

2-. The sulfate form of green rust is

examined in this study and has the general formula Fey2+Fex

3+(OH)3x+2y-2z(SO4)z (Randall

et al., 2001).

IV. Adsorption and adsorption models: a general overview

Stumm (1992) describes adsorption as “the accumulation of matter at the solid-

water interface”. A major ramification of this process is an alteration in the distribution

of molecules between the solid and solution, subsequently influencing their transport.

Adsorption can also change the electrostatic properties of particles and affect flocculation

and adhesion. Finally, adsorption will affect the reactivity of surfaces, altering rates of

precipitation, dissolution, and subsequent ion retention. Surface precipitation and

polymerization are not considered as forms of adsorption, but collectively the three

processes are referred to as ‘sorption’ (Sparks, 1995; Stumm, 1992).

Adsorption mechanisms are commonly divided into inner-sphere complexation,

outer-sphere complexation, and retention in the diffuse swarm. Inner-sphere complexes

are a result of a chemical reaction, are highly specific, and are quite strong. Electrostatic

attraction is responsible for outer-sphere complexes and retention in the diffuse swarm.

Outer-sphere complexes result from localized charge imbalances and are not specific

(Sparks, 1995). Figure 1.3 illustrates several inner- and outer-sphere complexes.

13

Figure 1.3: Diagram of various surface complexes. The top three (a-c) are examples of outer sphere complexes, while d (monodentate) and e (bidentate) are inner sphere complexes (from McBride, 1994).

-

+1-

2+

Oxygen Hydrogen Metal/divalent cation

+

+

1-

14

Anions are commonly observed to sorb to oxide and silicate minerals within soils

and in some cases bind to organic matter, with iron or aluminum typically acting as a

bridging ion. Chemisorption of anions occurs on ionizable functional groups on mineral

surfaces. Adsorption occurs through a ligand exchange process in which the anion

displaces H2O or OH-. Since H2O is a more mobile ligand than OH-, sorption is often

favored at lower pH. Four key characteristics of ligand exchange are the release of

hydroxyls, high specificity toward binding sites, apparent hysteresis, and the surface

charge becoming more negative after adsorption (McBride, 1994).

Adsorption isotherms are used to describe the adsorbed mass (normalized to the

mass of the substrate) as a function of the equilibrium concentration of the adsorbing

molecule at constant temperature. Isotherms have been classified into four main types: S,

L, H, and C. The L-type isotherm exhibits a decreasing slope with increasing

concentration, with the slope eventually becoming zero. This shape suggests a high

affinity between the adsorbed molecule and the substrate and often represents

chemisorption. The H-type isotherm is an exaggerated form of the L-type, suggesting

very strong interactions such as the formation of inner-sphere complexes. The C-type

(constant partitioning) isotherms insinuate partitioning between the solution and substrate

without any type of specific bonding. S-type isotherms usually have a slope that

increases with adsorptive concentration until sites are filled, when the slope decreases to

zero (Sparks, 1995; McBride 1994).

Several equilibrium-based models are used to describe adsorption on mineral

surfaces. The Freundlich equation, Langmuir equation, the diffuse double-layer model,

and the Stern model are commonly used. However, the Langmuir and Freundlich

equations are two of the most widely used and are incorporated in the following work.

Both models allow one to summarize and compare adsorption data, and possibly even to

predict adsorption behavior outside of experimental conditions. However, these

isotherms do not specifically indicate the type of sorption mechanism responsible for

adsorption, nor do they provide the speciation of surface complexes (Cornell and

Schwertmann, 1996).

The Langmuir equation was developed to describe gas adsorption on metal

surfaces but has been adapted for solid-solution interactions given the following

15

assumptions. First, a surface has a fixed number of identical sites which can each retain

one molecule, resulting in an adsorption maxima or monolayer coverage. Second,

adsorption is also reversible, and third, the molecules do not move laterally along the

surface of the substrate. Finally, adsorbate molecules do not react with each other and

the adsorption energy for all sites is identical, regardless of surface coverage (Sparks

1995).

The Freundlich equation is an empirical model that was first used to describe gas

and solute adsorption; it has also been used to describe ion retention by adsorption and

precipitation processes. This model does not predict an adsorption maximum and implies

that the adsorption energy of a surface varies with surface coverage (Sparks, 1995).

V. Phosphate sorption onto iron oxides

The prominent correlation between phosphorus and iron minerals in estuaries

(Pant and Reddy, 2001b; Roden and Edmunds, 1997), lake and ocean sediments (Young

and Ross, 2001; Olila and Reddy, 1997; Slomp et al., 1996; Olila and Reddy, 1995;

Moore and Reddy, 1994), and soils (Golterman, 1999; DeMello et al., 1998; Szilas et al.,

1998; Willett and Higgins, 1978) has led to numerous examinations of P adsorption

capacities and mechanisms by iron oxides and hydroxides. The most commonly studied

ferric minerals have been hematite and goethite since they are the most stable and most

abundant, although not necessarily the most reactive, iron minerals found in soils. They

are also easily synthesized in the lab and have well defined crystal structures (Cornell and

Schwertmann, 1996; Sparks, 1995; Torrent et al., 1994).

The mechanism of phosphate adsorption onto ferric oxides is generally dominated

by ligand exchange in which two singly coordinated hydroxyl groups or water molecules

are replaced by a single phosphate anion resulting in the formation of a bidentate,

binuclear complex (Reddy et al., 1999; Torrent, 1997; Colombo et al., 1994; Goldberg

and Sposito, 1985; Parfitt and Russell, 1977; Parfitt et al., 1975). The phosphate surface

complexes are very stable and result in slow exchange rates and an apparent

irreversibility (hysterisis) of phosphorus adsorption. The strength of these complexes

leads to long-term phosphorus storage in soils, sediments, and wetlands (Reddy et al.,

1999; Goldberg and Sposito, 1985; Parfitt et al., 1975).

16

Spectroscopic methods have been used to determine the type of surface complex

formed by phosphate on various oxides. Parfitt et al. (1975) and Parfitt and Russell

(1977) identified the formation of binuclear complexes on lepidocrocite, goethite,

hematite, ferric hydroxide gels, and β-ferric hydroxide using infrared (IR) spectroscopy.

Persson et al. (1996) examined phosphate adsorption by goethite and hematite also using

IR spectroscopy and suggested a monodentate inner-sphere complex was the primary

sorption mechanism. Three distinct surface complexes on goethite were observed, each

having a different degree of protonation as a function of pH. Unexpectedly, differences

between the hematite and goethite samples were noted. The authors proposed the

observed variations were due to surface precipitation on the hematite surface that resulted

from mineral synthesis. Arai and Sparks (2001) used attenuated total reflectance Fourier

transform infrared (ATR-FTIR) spectroscopy to show that the phosphate adsorption

mechanism on ferrihydrite was a non-protonated, binuclear complex at pH greater than or

equal to 7.5 and was a protonated, inner-sphere complex at sorption levels between 0.38

and 2.69 µM P/m2 from pH 4 to 6. Goethite, lepidocrocite, and hematite were all found

to form bidentate, binuclear complexes with phosphate, though the possibility of a

monodentate, mononuclear complex on goethite surfaces was not ruled out.

Goldberg and Sposito (1985) cited several lines of evidence for a ligand exchange

mechanism including hydroxyl release, adsorption and desorption kinetics, and

stereochemical calculations as well as spectroscopic evidence. A bidentate complex on

goethite was suggested on the basis of their IR data; however, it was noted that the

addition of water to the system may favor a monodentate complex. Phosphate was also

observed to behave very differently than other ions that are not specifically adsorbed,

particularly chloride. At a pH higher than the point of zero charge of a given iron oxide,

chloride sorption does not occur to any significant extent because the interaction is

entirely electrostatic. Phosphate can by sorbed by iron oxides at pH levels above the

PZC, however.

17

Table 1.1. Summary of previously determined phosphate adsorption capacities. Mineral P adsorption capacity (µM/m2) Reference Hematite 0.2 – 1.7 (1 day) Torrent, 1994 0.8 – 4 .1 (75 days) Torrent, 1994 2.6 (average) Torrent, 1994 2.1 – 2.4 Borggaard, 1983 HFO gel 643 – 1020 Ryden et al., 1977a Goethite ~2.5 Strauss et al., 1997 2.1 – 2.6 Borggaard, 1983 Maghemite 1.8 – 3.0 Borggaard, 1983 Akaganeite 6.2 Borggaard, 1983 Lepidocrocite 2.4 – 2.5 Borggaard, 1983 Feroxyhite 2.5 Borggaard, 1983 Amorphous Iron Oxides 3.1 Borggaard, 1983

Phosphate sorption by soils and pure mineral phases occurs in two distinct steps -

rapid adsorption taking place during the first few minutes of the reaction followed by a

slower stage (Reddy et al., 1999; Strauss et al., 1997; McLaughlin et al., 1977). Barrow

et al. (2000) proposed that lag times seen in experiments were due to variations in the soil

to solution ratio. Nooney et al. (1996) conducted sorption experiments of PO43- on

hematite under vacuum and observed very rapid adsorption within the first ten minutes,

followed by surface precipitation after 20 minutes. Crosby et al. (1984) found that

phosphate adsorption by iron hydroxides in estuaries obeyed first- and second order rate

expressions and that the adsorption rates were highly dependent upon the age of the

precipitate, the initial iron oxidation state, pH, and temperature. Torrent (1997)

summarized several other proposed theories explaining the slow adsorption step. These

include the initial formation of mononuclear complexes and rearrangement to form

binuclear complexes, replacement of silicate in the sorbent, competing anions on the

surface, and surface precipitation processes. Slow adsorption increases with specific

surface area, micro- and meso-porosity, and ferrihydrite impurities (Strauss et al., 1997;

Torrent 1997; Torrent et al., 1994), and has also been attributed to diffusion of phosphate

into the structural pores and defects of soil minerals (Ryden et al. 1977c). This slow

sorption has also been associated with the irreversibility of phosphate retention by soils

and pure iron oxides (Madrid and DeArambarri, 1985; Ryden and Syers, 1977).

The phosphate adsorption capacity of iron oxides generally averages 2.5 µM P/m2

and ranges from 1.5 to 3.5 µM/m2 (Reddy et al., 1999; Torrent 1997; Enyard, 1994;

18

Torrent et al., 1990; Borggaard, 1983). The findings of several previous studies are

summarized in Table 1. Torrent et al. (1994) found phosphate adsorption by natural

hematite specimens increased between 1 and 75 days. They also observed that hematite

had a lower affinity for phosphate and thus, a slower reaction rate than goethite.

McLaughlin et al. (1981) found that hematite, goethite, ferric hydroxide gel, and dried

ferric hydroxide gel all had very high affinities for phosphate with maxima of 6.9, 6.1,

6.9, and 4.8 µM/m2, respectively. The unusually high capacity of hematite in comparison

to goethite was attributed to variations during mineral synthesis, particularly the lack of a

sufficient crystallization period. In a study by Strauss et al. (1997), P adsorption on

crystalline goethite (surface area = 18 m2/g) had a sorption capacity near 2.5 µM P/m2

and was complete within a day. Several other less crystalline goethite samples reacted as

long as 3 weeks and exceeded the 2.5 µM P/m2 adsorption capacity. Borggaard (1983)

listed the ranges of phosphate adsorption by several iron oxides; overall, sorption ranged

from 1.8 to 6.2 µM P/m2.

Phosphate adsorption by synthetic and naturally occurring iron oxides has been

described by Langmuir and Freundlich isotherms, each of which with limitations

(Golterman, 1995). McLaughlin et al. (1977) and Ryden et al. (1977a) used three

separate Langmuir isotherms to describe different regions of an adsorption curve of

phosphate on ferric hydroxide gel. The first two regions (I and II) occurred at lower

equilibrium concentrations and were attributed to chemisorption, or ligand exchange;

region III was thought to be a more physical, ‘potential-determining’, sorption. It was

also noted that the equilibration time had significant influence on sorption in region I, the

steepest section of the adsorption curve, leading to increased sorption with time. Even

after 28 d sorption continued in region I, increasing the H-type character of the

adsorption curve, whereas regions II and III achieved equilibrium in less than a day.

These authors found that sorption by ferric hydroxide gel did not conform to first order

kinetics under the experimental conditions and that extractable phosphate decreased with

reaction time, suggesting that increased ligand exchange was the result of diffusion into

the gel. Ryden et al. (1977c) also observed that increases in ionic strength or decreases in

pH amplified the H-type character of the adsorption isotherm. Torrent et al. (1994) fit

19

sorption by hematite with a Freundlich isotherm modified with a time function. The

initial fast reaction was observed to take less than a day.

Surface area, crystallinity, and morphology of iron oxides are the best indicators

of phosphorus retention capacities and adsorption kinetics. Often a linear relationship is

observed between surface area and adsorption capacity (Enyard, 1993). Colombo et al.

(1994) found adsorption on hematite ranged from 0.31 to 2.27 µM P/m2 and this range

was attributed to variations in crystal size and morphology. More platy hematite crystals

diminish the proportion of binding sites (i.e., increased proportion of the chemically inert

[001] plane), decreasing the overall phosphate sorption capacity. Torrent et al. (1994)

also found sorption capacities of natural hematite, ranging from 0.8 to 4.1 µM P/m2, to be

influenced by crystal morphology. This study also noted that sorption by hematite

specimens exhibited a similar average to that of goethite, but it was much more

inconsistent due to the high degree of variability among hematite samples. Ainsworth et

al. (1985) found that aluminum substitution in goethite led to decreased crystal size,

increased surface area, and increased phosphate adsorption.

Solution pH has been found to influence which metal oxides will preferentially

sorb phosphate; it will also determine the extent, rate, and mechanism of adsorption by

iron oxides. In acidic soils, phosphate is most likely to be sorbed by amorphous iron or

aluminum oxides, while in more alkaline soils, phosphate is generally associated with

calcium and magnesium carbonates. Since the sorption capacity of CaCO3 is quite low,

this association has been credited to ferrihydrite impurities within the mineral

(McLaughlin et al., 1981) or, more probable, the formation of Ca-phosphate minerals. In

general, it appears that phosphorus sorption by ferric oxides decreases with increasing pH

(Figure 3) (Geelhoed, 1997; Willett and Cunningham, 1983; Ryden et al., 1977b;

McLaughlin et al., 1977; and Hingston et al., 1972). Strauss et al. (1997) found P

adsorption onto goethite is greatest at pH 2 and that pH was particularly important during

the initial fast reaction, with high pH slowing the reaction kinetics. Shang et al. (1992)

examined the kinetics and overall adsorption of phosphate by amorphous iron precipitates

and also found that as pH increased sorption decreased with the most prominent

difference in the fast adsorption step, defined as 0.5 h. The authors also found that total

phosphate sorption decreased with pH and attributed this occurrence to the presence of

20

fewer fully protonated (i.e, water) leaving groups, which are preferred over hydroxyl

groups. Torrent (1997) stated that phosphate formed mono- and bi-nuclear complexes on

goethite at low pH and binuclear complexes were only predominant near neutral or

higher pH. Arai and Sparks (2001) found phosphate sorption on ferrihydrite decreased

with pH from 3.5 to 9. Arai and Sparks (2001) also noted that at pH greater than 7.5,

phosphate formed non-protonated, bidentate, binuclear complexes on ferrihydrite, while

from pH 4 to 6 their evidence suggested the presence of protonated complexes.

Ionic strength and competition with other ions also influence phosphate retention.

Ryden et al. (1977c) and McLaughlin et al. (1977) found P adsorption onto ferric

hydroxide gel to increase with ionic strength. Ryden et al. (1977c) also noted that the

ionic strength effects on chemisorption were predominantly kinetic and that the effects on

physical adsorption were absolute. The authors attributed this behavior to the decreased

thickness of the diffuse layer with increased ionic strength, which would result in an

increased concentration of ions in the diffuse layer. Arai and Sparks (2001) found that

from pH 4 to 7.5, phosphate adsorption on ferrihydrite was unaffected by changes in

ionic strength; however, at pH greater than 7.5, adsorption increased with ionic strength.

Celi et al. (2000) observed increased phosphorus sorption by goethite with ionic strength

in agreement with Ryden et al. (1977c). However, Geelhoed et al. (1997) found that

increasing the ionic strength at low pH resulted in lower phosphate adsorption, but at

higher pH an increase in sorption was observed. Rietra et al. (2001) examined the

influence of calcium on phosphate adsorption by soils and goethite and found that at high

pH and at high concentrations of calcium, phosphate will be more strongly retained.

21

pH

2 4 6 8 10 12 140

20

40

60

80

100

Figure 1.4. Example of an experimental adsorption envelope for goethite (Hingston et al., 1972).

µM PO4 sorbed/ g goethite

µM P

O4 s

orbe

d/ g

goe

thite

22

Organic ligands are ubiquitous in soils and sediments and must not be ignored

when considering phosphate dynamics. Plants will often respond to phosphate deficiency

by releasing various organic ligands as root exudates, which make phosphate available by

promoting dissolution of phosphate minerals or desorption from iron oxides (Geelhoed et

al., 1999). The relationship between plant exudates and phosphate release has led to

studies of the competition of citrate and phosphate for sorption sites on goethite. Citrate

causes partial desorption of phosphate below pH 7, with the greatest effect at pH 4.5 to 5

(Geelhoed et al. 1999). Additionally, phosphate decreased citrate sorption over the entire

pH range (2 to 10) studied. In contrast, particulate organic matter did not hinder

phosphate adsorption to iron oxides in soils and would likely increase phosphate sorption

by inhibiting crystallization of iron and aluminum oxides (Borggaard et al., 1990).

Nevertheless, organic matter in solution could possibly inhibit sorption due to

competition for sorption sites. Gerke (1993) examined phosphate retention by iron oxide

mixtures in the presence of humic acids, which were found to reduce sorption initially.

However, sorption increased in long-term trials due to inhibition of iron oxide

crystallization by the formation of Fe-organic complexes. This study also observed that

sorption of phosphate by iron oxides decreased with aging when no humic substances

were present and that humic-Al or –Fe complexes exhibited higher phosphate sorption

capacities than the metal oxides alone.

VI. Need and Research Objectives

Many studies have discounted iron oxides as playing a significant role in

phosphate retention under anoxic conditions, suggesting phosphate will be released to

solution upon the reductive dissolution of the iron oxides or iron (ferric) phosphates.

However, Patrick and Khalid (1974) observed that anaerobic soils would release

phosphate to solutions having low initial phosphate concentrations, but the same soils

would also retain more phosphate from solutions with high initial concentrations than

their aerobic counterparts. The authors postulated that reductive dissolution of iron

oxides and the precipitation of ferrous minerals with high surface areas led to increased P

adsorption capacity, but that P would be less tightly held than in the aerobic environment.

Willett and Higgens (1978) examined reduced and re-oxidized rice soils and found

23

phosphate to be controlled by clays in the non-flooded soils, but by both ferric and

ferrous iron oxides in the seasonally flooded soils. These authors also observed increases

in phosphate adsorption capacity with flooding, which were correlated to oxalate

extractable iron. Upon drainage, the phosphate sorption capacity quickly decreased, but

still remained higher than the original level. The authors felt the ferrous hydroxides

produced under reducing conditions were likely to have higher surface areas than the

crystalline oxides present in the original soil. During re-oxidation, ferrous hydroxides

would convert to ferric phases, but with less crystallinity than the original materials.

Holford and Patrick (1979) examined the effect reduction and pH changes have

on phosphate sorption and mobility in acidic soils. They observed that anaerobic soils

could retain appreciable amounts of phosphate, owing to the precipitation of ferrous

phosphates and to sorption onto freshly precipitated ferrous hydroxides. Manahan (1994)

also observed that if phosphate concentrations are high in solution, iron, manganese, and

calcium minerals were all capable of sorbing phosphate even in anaerobic conditions.

Roden and Edmunds (1997) conducted several experiments in naturally anoxic lakes,

wetlands, and coastal marine sediments. Their findings indicated that large amounts of

phosphate released due to reductive dissolution of ferric phases were subsequently

retained by Fe2+ hydroxide phosphate complexes or ferrous phosphate minerals. The

authors also noted that if sulfate reduction was favored, Fe2+ would be sequestered in iron

sulfides and that phosphate would be released. Thus, P retention in anaerobic

environments is complex and site specific. Accordingly, the goal of this study is to

elucidate the binding capacity of phases common to anaerobic environments for

phosphate retention. Specifically, we investigate phosphate sorption on a number of

biogenic ferrous-bearing minerals that may be formed by microbial reduction of ferric

oxides and hydroxides in anaerobic environments (Fredrickson et al., 1998; Zachara et

al., 1998). Identified ferrous phases include siderite, vivianite, magnetite, and green rust.

Characterizing phosphate retention by these minerals is important since

microbially and chemically induced mineral transformations within soils, sediments, and

riparian zones can have a substantial influence on the mobility of phosphate, other

nutrients, and contaminants (Loyaux-Lawniczak et al., 2000; Erbs et al., 1999; Myneni et

al., 1997; and Hansen et al., 1996). Riparian zones are often important in determining the

24

behavior of contaminants in freshwater systems, since these areas can be pathways

between surface run-off and rivers. The changing oxidation – reduction conditions in

riparian zones will dictate the formation of specific iron minerals, partially through

biogenic processes. The results from this study will help determine how nutrients,

specifically phosphate, interact with iron phases in suboxic environments. This research

will also help to characterize interaction between iron and phosphate in riparian systems,

which will lead to a better understanding of the fate and transport of phosphate in the

freshwater environment. Knowledge of the extent of phosphate retention by mixed and

reduced iron minerals will also have implications in nutrient management from an

agricultural standpoint, particularly with respect to fields subjected to periodic flooding.

Alteration between oxic and suboxic conditions could lead to formation of the mixed and

reduced iron minerals formed and these phases will need to be considered when making

decisions about nutrient management and when evaluating phosphate dynamics in a

system.

The specific objectives of the first part of this study are (1) to determine

phosphate sorption capacities of siderite, magnetite, and sulfate green rust and compare

retention to that of lepidocrocite, goethite-coated sand, ferrihydrite-coated sand, and

ferrihydrite slurry, (2) to determine the influence of pH and ionic strength on retention

capacity, and (3) to compare phosphate sorption behavior on natural magnetite and

siderite to synthetic specimens. Sorption isotherms were determined using batch

experiments allowed to react for 48 h. Initial phosphate concentrations ranged from 10 to

1500 µM, with the pH adjusted to 7 and ionic strength to 0.1 M. Ionic strength effects

and pH effects were evaluated using 100 µM phosphate solutions.

The second component of this project examines phosphate retention under flow

through heterogeneous and homogenous mixtures of iron phases. A total of six

hydrodynamic column experiments were run. Three were anoxic and three were oxic.

The three oxic columns all contained only one iron phase, ferrihydrite-coated sand, and

two of the columns were replicates that were not buffered while the third column was

buffered. The three anoxic columns all consisted of heterogeneous mixtures of iron

oxides and hydroxides. In two of the anaerobic columns the iron reducing bacteria

Shewanella putrefaciens (CN 32) was used to facilitate the conversion of ferrihydrite to

25

magnetite, siderite, green rust, lepidocrocite, and goethite. Ferrous chloride was used to

drive magnetite precipitation in the third anoxic column.

The flow experiments are necessary in order to evaluate phosphate retention

dynamics under more realistic conditions. The specific objectives of the second section

of this project were to compare the phosphate adsorption of minerals in static and flow

conditions. The retention time and capacity of the columns were also examined; in

addition to assessing which iron phases were responsible for phosphate retention within

the columns. These experiments were also used to determine what other factors are

likely to influence phosphate retention in hydrodynamic systems.

26

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33

CHAPTER 2

Phosphate Adsorption by Mixed and Reduced Iron Minerals Under Static Conditions

INTRODUCTION I. Importance and cycling of phosphate

Phosphorus (P) availability is an important factor in biogeochemical cycling and

is often the limiting nutrient in freshwater and estuarine ecosystems (Reddy et al., 1999).

It is a macronutrient required by plants and animals and is a critical component in DNA,

RNA, ATP, and phospholipids (Schlesinger, 1991). Too little available P in agricultural

systems can result in poor productivity and crop failure. Conversely, increased P input to

fresh or saltwater systems via erosion can lead to large algal blooms and subsequent

eutrophication of water bodies (Brady and Weil, 1999). Noe et al. (2001) have recently

documented that even small P increases due to increased agricultural runoff in the

Everglades appears to be altering species distribution and, subsequently, the ecology of

the system.

The ultimate source of P to the biosphere is weathering of rocks, and the amount

required for biological production is much greater than the amount released by

weathering. This deficit results in P being highly conserved and rapidly cycled within the

biosphere (Filippelli and Souch, 1999; Berner and Berner, 1996). Vegetation, periphyton,

plankton, and microbes all act as biotic phosphorus sinks. However, biotic cycling is not

without P loss to abiotic mechanisms, including adsorption by soils and sediments,

mineral precipitation, and sedimentation. Minerals controlling P in soils and sediments

include apatite, hydroxyapatite, fluorapatite, octocalcium phosphate, strengite, vivianite,

variscite, and wavellite (Reddy et al., 1999). Phosphorus adsorption by soils and lake

sediments has often been correlated to amorphous iron, aluminum, and manganese oxides

and hydroxides, as well as calcium and magnesium (Brady and Weil, 1999; Reddy et al.,

1999; Torrent, 1997; Golterman 1995; Olila and Reddy, 1995; Blanchar and Frazier,

1991; Williams et al., 1971). The relationship between P and ferric oxides is usually

more apparent in acidic to neutral soils, while in alkaline soils P is often associated with

calcium or magnesium (Reddy et al., 1999).

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II. Phosphorus retention processes

The prominent correlation between phosphorus retention and iron has led to

numerous examinations of P adsorption capacities and mechanisms onto iron oxides and

hydroxides in aerobic environments. The most commonly studied ferric minerals have

been hematite and goethite since they are the most abundant, although not necessarily the

most reactive, iron minerals found in soils (Torrent et al., 1994; Sparks, 1995).

The mechanism of phosphate adsorption onto ferric oxides is dominated by a

ligand exchange reaction in which two singly coordinated hydroxyls or water molecules

are replaced by a single phosphate anion resulting in the formation of a bidentate,

binuclear complex (Reddy et al., 1999; Torrent, 1997; Colombo et al., 1994; Parfitt and

Russell, 1977; Parfitt et al., 1975). Parfitt et al. (1975) used infrared (IR) spectroscopy to

verify the formation of binuclear complexes on lepidocrocite, goethite, hematite, ferric

hydroxide gels, and β-ferric hydroxide. Arai and Sparks (2001) used attenuated total

reflectance Fourier transform infrared (ATR-FTIR) spectroscopy to show that the P

adsorption mechanism on ferrihydrite was a non-protonated, binuclear complex at pH

greater than 7.5. Binuclear complexes are quite strong and result in slow exchange rates

and an apparent irreversibility of phosphorus adsorption. The strength of these

complexes lead to long-term phosphorus storage in soils, sediments, and wetlands (Reddy

et al., 1999; Parfitt et al., 1975). However, formation of mononuclear complexes cannot

be overlooked. Torrent (1997) stated that P formed mono- and binuclear complexes on

goethite at low pH and binuclear complexes were only predominant near neutral or

higher pH. Nonetheless, it is clear that phosphorus forms very stable surface complexes

on iron (hydr)oxide solids.

Two stages of phosphate sorption transpire. Rapid adsorption takes place during

the first few minutes of the reaction and is followed by a slower stage, which may be

measurable for days to weeks, attributed to diffusion into the solid phase (Reddy et al.,

1999; Strauss et al., 1997; McLaughlin et al., 1977). Torrent (1997) summarized several

other proposed theories explaining the slow adsorption step. These include the initial

formation of mononuclear complexes and rearrangement to form binuclear complexes,

replacement of silicate in the sorbent, competing anions on the surface, and surface

precipitation processes. Slow adsorption increases with specific surface area, micro- and

35

meso-porosity, and ferrihydrite impurities (Strauss et al., 1997; Torrent 1997; Torrent et

al., 1994). This slow sorption has also been associated with irreversible phosphate

retention (Madrid and DeArambarri, 1985).

Phosphate adsorption capacity generally averages 2.5 µmols P/m2 for most iron

oxides, ranges from 1.5 to 3.5 µmols/m2 (Reddy et al., 1999; Torrent 1997; Enyard, 1994;

Torrent et al., 1990; Borggaard, 1983). Surface area and crystallinity are the best

indicators of phosphorus retention capacities (Strauss et al., 1997; Colombo et al., 1994;

Enyard, 1993; Borggaard, 1983; McLaughlin et al., 1981). Phosphate adsorption on

synthetic iron oxides has been described by the Langmuir and Freundlich isotherms

(Miltenburg and Golterman, 1998; Torrent, 1997; Golterman, 1995; Gerke, 1993;

McLaughlin et al., 1977).

Colombo et al. (1994) found adsorption on hematite ranged from 0.31 to 2.27

µmols P/m2 and was described by the Freundlich isotherm. This range was attributed to

variations in crystal morphology and size, and it is supported by the work of Torrent et al.

(1994), where adsorption capacities of natural hematite were reported between 0.8 and

4.1 µmols P/m2. Sorption onto hematite was more variable than onto goethite. The rate

and adsorption capacity of phosphate on goethite is highly dependant on crystallinity

(Strauss et al., 1997). Crystalline goethite with a surface area of 18 m2/g had an

adsorption capacity near 2.5 µmols P/m2, and adsorption was complete within a day.

Several other less crystalline goethite samples reacted as long as 3 weeks and exceeded

the 2.5 µmols P/m2 adsorption capacity.

In general, it appears that phosphorus sorption by ferric hydroxide gel and

goethite decreases with increasing pH (Geelhoed, 1997; Willett and Cunningham, 1983;

Ryden et al., 1977a; McLaughlin et al., 1977; Hingston et al., 1972). Strauss et al. (1997)

found P adsorption is greatest at pH near 2. Shang et al. (1992) suggested pH has an

appreciable influence on reaction rate in the beginning of sorption reactions and noted an

overall decrease in phosphorus adsorption capacity with pH from pH 4 to 6. Strauss et al.

(1997) also observed P sorption onto goethite occurred most rapidly at pH 2 and that pH

was particularly important during the initial fast reaction, with high pH slowing the

reaction kinetics. Shang et al. (1992) examined the kinetics and overall adsorption of

phosphate by amorphous iron precipitates and also found that as pH increased sorption

36

decreased with the most prominent difference in the fast adsorption step, defined as 0.5 h.

The authors also found that total phosphate sorption decreased with pH and attributed this

occurrence to the presence of fewer fully protonated (i.e, water) leaving groups, which

are preferred over hydroxyl groups. Torrent (1997) stated that phosphate formed mono-

and binuclear complexes on goethite at low pH and binuclear complexes were only

predominant near neutral or higher pH. Arai and Sparks (2001) found phosphate sorption

on ferrihydrite decreased with pH from 3.5 to 9. Arai and Sparks (2001) also noted that

at pH greater than 7.5, phosphate formed non-protonated, bidentate, binuclear complexes

on ferrihydrite, while from pH 4 to 6 their evidence suggested the presence of protonated

complexes. Understanding how the sorption characteristics of a given oxide change with

pH is important when considering phosphate cycling. Oxides with well defined, sharp

adsorption envelopes are likely to be more sensitive to pH fluctuations and are less likely

to strongly retain phosphate under changing pH conditions.

Ionic strength may also have an important influence on phosphate retention.

Ryden et al. (1977b) and McLaughlin et al. (1977) found P adsorption onto ferric

hydroxide gel to increase with ionic strength. Arai and Sparks (2001) found that from pH

4 to 7.5, P adsorption on ferrihydrite was unaffected by changes in ionic strength;

however, at pH greater than 7.5, adsorption increased with ionic strength. Work by Celi

et al. (2000) on goethite has shown increased ionic strength results in increased

phosphorus sorption as well. However, Geelhoed et al. (1997) found that increasing the

ionic strength at low pH resulted in lower P adsorption but at higher pH an increase in

sorption was observed.

III. Needs and Research Objectives

Many studies have discounted iron oxides as playing a significant role in P

retention under anoxic conditions, suggesting P will be released to solution upon the

reductive dissolution of the iron (III) oxides or phosphates. However, Patrick and Khalid

(1974) found the adsorption capacity for iron-rich sediments increases under reducing

conditions. They postulated that reductive dissolution of iron oxides and hydroxides and

subsequent precipitation of ferrous minerals with high surface areas increased P

adsorption capacity in an aerobic environment, but that the P would be less tightly held

37

than in the aerobic environment. Thus, there are conflicting sentiments regarding P

retention in anaerobic environments. Accordingly, the goal of this study is to elucidate

the binding capacity of phases common to anaerobic environments for P retention.

Specifically, we investigate P sorption on a number of biogenic ferrous-bearing minerals

that may be formed by microbial reduction of ferric oxides and hydroxides in anaerobic

environments. Identified ferrous phases include siderite, vivianite, magnetite, and green

rust (Fredrickson et al., 1998; Zachara et al., 1998).

Characterizing P retention by these minerals is important since microbially and

chemically induced mineral transformations within soils, sediments, and riparian zones

can have a substantial influence on the mobility of P along with other nutrients, and

contaminants (Loyaux-Lawniczak et al., 2000; Erbs et al., 1999; Myneni et al., 1997; and

Hansen et al., 1996). Riparian zones are often important in determining the behavior of

contaminants in freshwater systems, since these areas can be pathways between surface

run-off and rivers. Changing oxidation – reduction conditions in riparian zones will

dictate the formation of specific iron minerals, partially through biogenic processes.

Results from this study will help determine how phosphate interacts with iron phases in

suboxic environments. More importantly, this study will help to elucidate phosphorus

dynamics within riparian zones, thus improving our understanding of the fate and

transport of phosphate in the freshwater environment. Knowledge of the extent of P

retention by mixed and reduced iron minerals will also have implications in nutrient

management from an agricultural standpoint, particularly with respect to fields subjected

to periodic flooding. Alteration between aerobic and anaerobic conditions could lead to

formation of the mixed and reduced iron minerals and these phases will need to be

considered when making decisions about nutrient management and when evaluating P

dynamics in a system.

The specific objectives of this study are (1) to determine P sorption capacities of

siderite, magnetite, and sulfate green rust and compare retention to that of lepidocrocite,

goethite-coated sand, ferrihydrite-coated sand, and ferrihydrite slurry, (2) to determine

the influence of pH and ionic strength on retention capacity, and (3) to estimate P

sorption within anaerobic and aerobic soils. Sorption isotherms were determined using

batch experiments allowed to react for 48 hours. Initial P concentrations from 10 to 1500

38

µM were evaluated. Ionic strength and pH effects were investigated using 100 µM P

solutions with ionic strength of 0.05 M to 0.1 M and pH from 4.5 to 9.

39

MATERIALS AND METHODS

Batch studies were conducted using several mixed and reduced iron minerals as

sorbents. Natural and synthetic samples of magnetite and siderite were studied in

addition to green rust, ferrihydrite, ferrihydrite coated sand, goethite coated sand, and

lepidocrocite.

I. Mineral preparation

Naturally occurring siderite was purchased from Ward’s Scientific (46E7351,

Antigonish County, Nova Scotia) and naturally occurring magnetite was obtained from

the Stanford University Research Mineral Collection. Synthetic magnetite was purchased

from Aldrich Chemical (1317-61-9). The natural (of geologic origin) magnetite and

siderite samples were crushed and passed through a 250 µm sieve. Both magnetite

samples were treated with 1 M HCl for 1 h to remove oxidized surface constituents,

rinsed three times with anoxic, deionized water and dried using vacuum filtration. The

siderite was treated using 0.05 M HCl for 1 h, rinsed, and dried. All mixed or reduced

iron minerals were stored in an anaerobic glovebox (Coy Products, Ann Arbor, MI)

having O2(g) levels less than 5 ppm.

Synthetic siderite was made inside an anaerobic glovebox using the method

described by Hallbeck et al. (1993). Anoxic water was used to prepare 0.1 M

Fe(NH4)2(SO4)2·6H2O and 0.1 M Na2CO3 solutions. Water was made anoxic by boiling

and bubbling with oxygen-free N2 for 1 h prior to making solutions. A cloudy white

precipitate formed immediately upon the addition of 500 mL of the Na2CO3 solution to

500 mL of the Fe(NH4)2(SO4)2. The solid was allowed to settle, rinsed three times with

anoxic deionized water, and dried using vacuum filtration followed by desiccation.

Ferrihydrite (2-line) was made using the method described by Cornell and

Schwertmann (1996). A ferric chloride solution (2.3 L of 0.062 M FeCl36H2O, pH near

2) was titrated with 0.2 M NaOH to pH 7. The solution was then stirred and allowed to

stabilize for several hours with pH maintained by base addition. Once the solution pH

was stable, the precipitate was allowed to settle and rinsed three times with deionized

water. Ferrihydrite coated sand was made by drying the slurry onto 660 g of Iota 6

40

Quartz Sand (Umimin Corporation). The sand was dried at room temperature, rinsed

several times with deionized water, and dried again.

Goethite was synthesized using a method described by Cornell and Schwertmann

(1996). A 1 M NaHCO3 (110 mL) was added to 1L of 0.5 M FeCl2 4H2O (made with

deoxygenated water) while the latter was stirred and bubbled with N2. Upon addition of

the NaHCO3 the flow of N2 was replaced with air at a rate of 35-40 mL/minute. The

solution was stirred and bubbled for 48 h. When oxidation was complete, solids were

allowed to settle and the supernatant decanted. The remaining slurry was then washed by

suspension in a 0.1 mM NaCl solution followed by centrifugation. This rinse procedure

was repeated a total of five times and after the last rinse the slurry was sonicated for 1 h.

Goethite-coated sand was made by drying the slurry onto 278 g of Iota 6 Quartz Sand

(Umimin Corporation). The sand was dried at room temperature, rinsed several times

with deionized water, and dried again.

Lepidocrocite was synthesized following a method described by Cornell and

Schwertmann (1996). A 0.6 M FeCl2 solution was titrated to pH 6.5 using 1 M NaOH.

Air was then introduced to the solution by bubbling at a rate of ~250 mL/min. The pH

was maintained between 6.3 and 6.5 using 1 M NaOH and the process terminated when

no additional base was necessary to maintain pH.

Sulfate green rust was precipitated using a modification of the method described

in Schwertmann and Fetcher (1994). A 0.5 M ferrous sulfate solution (pH near 3.7) was

prepared using anoxic deionized water. The ferrous sulfate solution was kept in an

anoxic reaction vessel and the pH was brought to 7.1 using 4 N NaOH. Air was then

introduced at 60 mL/min while additions of 4 N NaOH or 1 N HCl were added to

maintain the pH. The reaction was complete when the Eh stabilized.

II. Mineral characterization

X-ray diffraction was carried out using a Rigaku Geigerflex diffractometer

(Model CN2029) with Cu Kα radiation (35 kV, 15 mA). Samples were prepared for

analysis by pressing the powdered mineral material into a 0.5 mm depression on a Rigaku

mono-crystalline silica XRD slide. The synthetic siderite was easily oxidized and thus

the slide needed to be prepared in the glovebox. In order to prevent oxidation during

41

analysis, a small amount of the solid was mixed with glycerol to form a thick paste,

which was then packed into the depressed slide. The sample was analyzed immediately

upon removal from the glovebox.

Specific surface area was determined by N2 adsorption using the BET equation

(Brunauer et al., 1938). The adsorption curves were obtained using a Coulter SA 3100

surface area analyzer. All minerals, except synthetic siderite and sulfate green rust, were

degassed by heating overnight at 85˚C. Contact with the atmosphere was avoided by

maintaining positive pressure (helium) on the sample containers. Excess water was

removed from the synthetic siderite and green rust by dessication under a vacuum for 1

week.

III. Phosphate adsorption experiments

Solutions used to define the adsorption isotherms ranged from 10 µM to 1500 µM

PO4 and were made using Na2HPO4. Ionic strength was adjusted to 0.1 M using NaCl.

The pH was adjusted to 7 using anoxic 0.05 M HCl or 0.05 M NaOH. The pH effects

were evaluated using 100 µM PO43- solutions adjusted to pH 4, 5.5, 7, 8, and 9. Ionic

strength influences were investigated using 100 µM PO43- solutions at pH 7 with the ionic

strength adjusted between 0.1 and 0.005 M. All solutions used with mixed or reduced

iron minerals were made under anaerobic conditions.

All experiments using mixed or reduced iron minerals were prepared in an

anaerobic atmosphere. Acid-cleaned glassware was allowed to equilibrate in the

glovebox at least 12 h before use. Approximately 0.1 g mineral was weighed into 125

mL serum vials, followed by the addition of 100 mL of solution. The supply of green

rust was limited, so about 0.02 g mineral was weighed into a 25 mL serum vial and 15

mL of phosphate solution added. The vials were sealed, removed from the glovebox, and

shaken using an orbital shaker (120 rpm) for 48 h. All samples were run in triplicate,

with the exception of the sulfate green rust, which were run in duplicate. Upon

completion, 8 mL of solution were passed through a 0.45µm filter and acidified using two

drops of HNO3. Starting and ending solutions were analyzed using inductively coupled

plasma-optical emission spectroscopy (ICP-OES). The amount of phosphate adsorbed

was determined from the difference between the initial and final concentrations.

42

IV. Release of sulfate from green rust

Sulfate release from green rust with increasing ionic strength was evaluated by

allowing the mineral to react for 48 h with deionized water and anoxic NaCl solutions

ranging in ionic strength from 0.001 to 0.1 M. Approximately 0.02 g green rust was

weighed into a 25 mL serum vial, 15 mL solution added, and the sample placed on an

orbital shaker (120 rpm). A solution containing only 100 µM NaH2PO4 was also

assessed using the same procedure.

43

RESULTS

I. Adsorption capacities and isotherms

Phosphate adsorption capacities for several minerals were higher than expected.

Adsorption capacities were determined by measuring the amount of phosphate in solution

before and after reaction with a known quantity of mineral. The amount of phosphate

lost from solution was assumed to be adsorbed by the mineral phase. Experimental

adsorption capacities ranged from 4.73 to 10.37.55 µmols PO4/g mineral (Table 2.1).

Goethite coated sand exhibited the lowest adsorption capacity and ferrihydrite the

greatest capacity on a mass basis.

Mineral surface areas ranged from 0.13 to approximately 200 m2/g. On a surface

area basis, sulfate green rust exhibits the greatest capacity for phosphate (343.23 µmols

PO4/m2) and lepidocrocite the lowest (1.65 µmols PO4/m2). Surface area values and

normalized sorption maxima are summarized in Table 2.2.

Adsorption curves for geologic magnetite, green rust, and geologic siderite

(Figures 1a-c) were best described by the Freundlich equation (R2 = 0.88, 0.99, and 0.97,

respectively). The linear form of the Freundlich equation is

log (x/m) = (1/n) log C + log K [2.1]

where K is the distribution coefficient, C is the equilibrium concentration, x/m is mass

phosphate adsorbed/mass mineral, and n is a correction factor. The calculated

distribution coefficients (K) were 0.0031 (magnetite), 0.535 (green rust), and 0.017

(siderite). Correction factors (n) for the three minerals were 1.38, 1.35, and 1.47

respectively.

The Langmuir equation best describes the adsorption curves for ferrihydrite,

ferrihydrite-coated sand, goethite-coated sand, and lepidocrocite (Figure 2.2), as well as

synthetic siderite and synthetic magnetite (Figure 2.3). The calculated monolayer

capacities and binding coefficient values for these equations are listed in Table 2.3.

44

Table 2.1. Maximum observed phosphate adsorption on a mass basis of the mineral adsorbent (µmols PO4/g mineral). Mineral

Maximum observed adsorption (µmols/g)

SD*

Goethite coated sand 4.73 0.36 Natural magnetite 13.61 4.20 Ferrihydrite coated sand 14.39 0.38 Natural siderite 26.97 2.15 Synthetic magnetite 34.56 3.98 Lepidocrocite 123.82 0.52 Sulfate green rust 878.67 131.84 Synthetic siderite 957.53 21.46 Ferrihydrite 1037.55 1.07 *SD = standard deviation.

Table 2.2. Surface area (SA) and SA normalized adsorption maxima for phosphate on various iron oxides.

Table 2.3. Summary of adsorption capacities (b-values) and KL-values for minerals conforming to the Langmuir isotherm.

Mineral Monolayer capacity (µmols /g) Binding coefficient (KL) Ferrihydrite slurry 1022.32 0.145 Synthetic siderite 974.94 0.040 Lepidocrocite 122.99 0.163 Synthetic magnetite 34.11 0.0407 Ferrihydrite-coated sand 14.11 0.042 Goethite sand 4.63 0.145

Mineral

Surface Area (m2/g)

Maximum adsorption normalized to SA (µmols PO4/m2)

Sulfate green rust 2.56 343.23 Natural magnetite 0.13 104.69 Natural siderite 0.27 99.89 Synthetic magnetite 1.74 19.86 Synthetic siderite 63.05 15.19 Ferrihydrite-coated sand 2.28 6.31 Ferrihydrite slurry 200 (estimated) 5.19 Goethite sand 2.39 1.98 Lepidocrocite 75.21 1.65

45

II. Influence of pH and ionic strength

Generally P sorption decreased with increasing pH, at least beyond a critical value

(Figure 2.4). Goethite- and ferrihydrite-coated sand both exhibited linear decreases in

adsorption. Natural siderite and natural and synthetic magnetite all exhibit a relatively

well-defined adsorption envelope, with minimal P retention at pH 9. Lepidocrocite

appears to behave similarly to both magnetite specimens and natural siderite; however,

sorption was not examined at pH 9 and the envelope thus cannot be constrained.

Ferrihydrite and sulfate green rust showed only slight decreases in adsorption with

increased pH. Synthetic siderite noticeably deviates from the general trend. Phosphate

adsorption reaches a maximum at pH 7 and there was little difference in sorption at pH 4

and 9.

Overall, changes in ionic strength had a negligible influence on P adsorption

within the pH range examined. The only appreciable change in sorption observed with

ionic strength was on sulfate green rust (Table 2.4).

III. Release of sulfate from green rust

Sulfate release from the interlayer of the green rust is likely to be an important

consideration when attempting to evaluate P adsorption. Hansen and Poulsen (1999)

found that approximately 50% of the interlayer sulfate was replaced by phosphate,

ultimately leading to the formation of vivianite. The release of sulfate from green rust

increased from an average of 245 µmols SO4/g in deionized water to 632 µmols SO4/g in

0.1 M NaCl solution (Figure 2.5). The amount of sulfate released into the 0.1 M ionic

strength phosphate solutions used to establish sorption isotherms ranged from 437 to 677

µmols SO4/g (Figure 2.6). The average sulfate released into 100µM NaH2PO4 solution

with no NaCl was 306 µmols SO4/g and release into 100µM NaH2PO4 with ionic strength

adjusted to 0.1 M using NaCl was 558 µmols SO4/g (Figure 2.6). It should also be noted

that phosphate adsorption was greater, 95.8 µmols PO4/g, when the solution contained no

NaCl. Adsorption averaged 66.3 µmols PO4/g when the ionic strength was adjusted to

0.1 M using NaCl.

46

Table 2.4. Average adsorption (µmols PO4/g) at high and low ionic strength. Comparisons were done using an initial 100 µM NaH2PO4 solution at pH 7. Mineral 0.05 M SD* 0.1 M SD*

Green rust 110.14 14.22 66.55 3.37 Natural magnetite 1.16 0.11 0.84 0.63 Ferrihydrite 112.14 0.32 108.56 0.84 Synthetic magnetite 26.75 0.84 28.22 1.16 Lepidocrocite 84.87 1.26 90.98 2.53 Goethite coated sand 4.00 0.11 4.42 0.02 Synthetic siderite 109.40 12.00 122.57 9.90 Ferrihydrite coated sand 6.42 0.11 7.27 0.21 Natural siderite 2.74 0.42 3.79 0.21

*SD = standard deviation.

47

0 200 400 600 800 1000 1200 1400

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

log C0 1 2 3

log

(x/m

)

-2.0

-1.5

-1.0

-0.5

0.0

0.5

0 200 400 600 8000

20

40

60

80

100

log C0 1 2 3

log

(x/m

)

-0.5

0.0

0.5

1.0

1.5

2.0

2.5

0 200 400 600 800 1000 1200 1400

0.0

0.5

1.0

1.5

2.0

2.5

3.0

log C0.5 1.0 1.5 2.0 2.5 3.0 3.5

log

(x/m

)

-1.2

-0.8

-0.4

0.0

0.4

Figure 2.1. Phosphate adsorption curves for minerals fit to a Freundlich isotherm (linear form is inset): (a) geologic magnetite (R2 = 0.89), (b) sulfate green rust (R2

= 0.99), and (c) geologic siderite (R2 = 0.97). In the case of sulfate green rust data points obscure the error bars. Observed adsorption maxima were: 13.61 µmols PO4/g magnetite, 878.67 µmols PO4/g GR, and 26.97 µmols PO4/g siderite.

C (equilibrium concentration µM PO4)

x/m

(mg

PO4

/ gra

m m

iner

al)

a

c

b

48

Figure 2.2. Adsorption isotherms of ferric (hydr)oxides all fit to the Langmuir isotherm (linear from in inset). (a) ferrihydrite (R2 = 0.99), (b) ferrihydrite-coated sand (R2 = 0.99), (c) goethite-coated sand (R2 = 0.99), (d) lepidocrocite (R2 = 0.99). Error bars in a are obscured by data points. The tabulated Kl and b-values are listed in Table 3. Observed maxima were: 1037.55 µmols PO4/g ferrihydrite, 14.39 µmols PO4/g ferrihydrite-coated sand, 4.73 µmols PO4/g goethite-coated sand, and 123.82 µmols PO4/g lepidocrocite.

0 100 200 300 400 5000

20

40

60

80

100

C0 100 200 300 400 500 600

C/(x

/m)

0

1

2

3

4

5

6

0 200 400 600 800 1000 1200 1400-0.1

0.0

0.1

0.2

0.3

0.4

0.5

0.6

C0 400 800 1200 1600

C/(x

/m)

0

1000

2000

3000

4000

0 200 400 600 800 1000 1200 1400

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

1.6

C0 400 800 1200 1600

C/(x

/m)

0

400

800

1200

0 200 400 600 800 1000 1200 14000

2

4

6

8

10

12

14

C0 400 800 1200 1600

C/(x

/m)

0

20

40

60

80

100

120

140

c d

ba x/

m (m

g PO

4 / g

ram

min

eral

)

C (equilibrium concentration µM PO43-)

49

0 100 200 300 400 500

0

20

40

60

80

100

C0 100 200 300 400 500 600

C/(x

/m)

0

1

2

3

4

5

6

7

0 200 400 600 800 1000 1200 1400

0

1

2

3

4

C0 400 800 1200 1600

C/(x

/m)

0

200

400

600

Figure 2.3. Phosphate adsorption curve fit to a Langmuir isotherm for (a) synthetic siderite (linear form in inset, R2 = 0.96) and (b) synthetic magnetite (linear form in inset, R2 = 0. 98). The tabulated Kl and b-values are listed in Table 3. Observed adsorption maxima were 957.53 µmols PO4/g siderite and 34.56 µmols PO4/g magnetite.

Ads

orbe

d P

Con

cent

ratio

n, x

/m (m

g PO

4/g so

lid)

Equilibrium PO4 concentration, C (µM)

a

b

50

0

2

4

6

8

10

12

14

pH3 4 5 6 7 8 9 10

0.0

0.2

0.4

0.6

0.8

1.0

Figure 2.4. Influence of pH on phosphate adsorption to various iron oxides. Comparisons were done using an initial 100 µM NaH2PO4 solution with I = 0.1 M.

Ferrihydrite sand

Goethite sand

Geologic Siderite

Geologic magnetite

Green rust

Synthetic siderite

Ferrihydrite slurry

Ads

orbe

d P

Con

cent

ratio

n, x

/m (m

g PO

4/g so

lid)

Synthetic magnetite

51

NaCl Concentration (M)

0.00 0.02 0.04 0.06 0.08 0.10

mg

SO4

Rel

ease

d (m

g SO

4/g G

R)

10

20

30

40

50

60

70

Figure 2.5: Sulfate released from green rust with increasing ionic strength (as NaCl).

No phosphate

100 µM PO4 (no NaCl)

52

0 200 400 600 800 1000

Con

cent

ratio

n (m

g/g

GR

)

0

20

40

60

80

100

Figure 2.6. Phosphate adsorption and sulfate release with increasing phosphate concentration. Also shown is phosphate sorbed and sulfate released into 100 µM phosphate solution when no NaCl is present.

PO4 sorbed SO4 released

PO4 sorbed at 100 µ PO4, no NaCl SO4 released at 100 µ PO4, no NaCl

Equilibrium Phosphate Concentration (µM)

53

DISCUSSION

Sorption capacities for the ferric iron minerals examined in this study ranged from

1.65 to 6.31 µmols PO4/m2. Ferrihydrite coated sand had the greatest capacity (6.31

µmols PO4/m2), followed by ferrihydrite (5.19 µmols PO4/m2), goethite-coated sand (1.98

µmols PO4/m2), and lepidocrocite (1.65 µmols PO4/m2). A review by Torrent (1997)

listed the ranges of adsorption capacities for ferrihydrite (2.42 – 3.05 µmols PO4/m2),

goethite (2.42 – 4.63 µmols PO4/m2), and lepidocrocite (1.47 – 2.63 µmols PO4/m2). The

experimental results from this work are comparable with regard to lepidocrocite;

however, the results for the remaining ferric phases are not as easily compared to the

previously determined values.

The adsorption capacity determined in this study for ferrihydrite is higher than

those cited by Torrent (1997). This is likely the result of underestimating the surface

area, which has been found to range from 100 to 700 m2/g (Cornell and Schwertmann,

1996). A value of 200 m2/g was assumed during the data analysis of this work. Based

on the sorption capacity per weight of ferrihydrite in this study, a surface area of ~400

m2/g, well within the documented range, would have resulted in a value near that

previously published.

A batch sorption experiment conducted on uncoated silica sand indicated that the

sand alone adsorbed no measurable phosphate; therefore, phosphate adsorption by the

goethite- and ferrihydrite-coated sands is due entirely to the respective iron phase. This

is also reflected in the results of the surface area analyses. Uncoated silica had a surface

area of 0.14 m2/g, while ferrihydrite- and goethite-coated sand yielded values of 2.28

m2/g and 2.39 m2/g, respectively. Iron content of the sand (weight percent) was 0.35%

for the ferrihydrite-coated sand and 0.42% for the goethite-coated sand. When

comparing the sorption capacities of the ferrihydrite and ferrihydrite-coated sand

normalized to iron (weight percent), sorption by the coated-sand is much greater (~2000

µmols PO4/g Fe as compared to ~4100 µmols PO4/g Fe). This may be the result of

aggregation of the ferrihydrite slurry (no sand), which would decrease the number of

sorption sites readily available.

Phosphate sorption by all four ferric iron minerals was fit to a Langmuir isotherm.

The initial slopes of the adsorption curves for lepidocrocite, ferrihydrite, and goethite-

54

coated sand were quite steep; however, the initial section of ferrihydrite-coated sand

adsorption curve is more gradual. These observations are reflected in the values of the

binding coefficients (KL), which are 0.145 for ferrihydrite and goethite coated sand and

0.163 for lepidocrocite. The KL for ferrihydrite-coated sand was 0.042. The steep

adsorption curves and high binding coefficients indicate a high affinity for phosphate by

lepidocrocite, ferrihydrite, and goethite-coated sand, while the ferrihydrite-coated sand

has a lower affinity for phosphate on a mass basis.

The reduced and mixed valent iron phases studied exhibited sorption capacities

higher than the oxidized minerals when normalized to surface area, ranging from 15.16 to

343.27 µmols PO4/m2. Synthetic magnetite and synthetic siderite reach adsorption

maxima of 15.16 and 19.90 µmols PO4/m2, respectively. The binding coefficients are

similar to that of ferrihydrite-coated sand with the KL for siderite calculated at 0.040 and

magnetite at 0.0407. The slopes of the initial section of the adsorption curves reflect

these lower binding affinities, which are similar to that of ferrihydrite-coated sand.

The natural magnetite, natural siderite, and sulfate green rust conformed to the

Freundlich isotherm and thus no adsorption maximum was observed. This is likely

explained by surface precipitation and/or physical adsorption of phosphate in addition to

chemisorption at the mineral surface. To evaluate whether surface precipitation occurred,

samples of natural siderite, natural magnetite, and sulfate green rust were reacted with

1500 µM NaH2PO4 (pH 7 and ionic strength adjusted 0.1 M using NaCl) for 48 h, dried,

and analyzed using X-ray diffraction. The ending solution was analyzed for phosphate

and dissolved iron; the chemical speciation program MINTEQ was used to determine

saturation indices using parameters from Stumm and Morgan (1981) and Benjamin

(2002) (Table 2.5). X-ray diffraction patterns did not show evidence of iron phosphate

mineral phases; however, if surface precipitates were present as a small percentage of the

sample they would not be detected. Vivianite was found to be oversaturated in solutions

reacted with magnetite and siderite, so it is possible for precipitation to occur.

The solution reacted with sulfate green rust was slightly undersaturated with

respect to vivianite (Table 2.5); however, based on the work by Hansen and Poulsen

(1999) conversion of green rust to vivianite is a likely phosphate sink in these systems.

They reacted sulfate green rust with 50 mM Na2HPO4 (pH 9.3) for 60 d.

55

Table 2.5. Concentration data and saturation indices for phosphate solutions after 2 d reaction with geologic magnetite, geologic siderite, and sulfate green rust.

*SI= saturation index (log IAP/Ksp, where IAP is the ion activity product for the respective phase) Vivianite was detected within 1 dl of phosphate addition and most of the vivianite

precipitation occurred during the first 10 d of the reaction; however, continued

precipitation was noted after 40 d. The first phase of precipitation was attributed to

vivianite formation resulting from the reaction of Fe(II) in solution with the added

phosphate. The later phase was believed to follow a period of phosphate reaction with

Fe(II) within the green rust resulting in vivianite precipitation in the interlayer. They also

found that the interlayer thickness decreased and that sulfate was released into solution as

a result of exchange with phosphate. The authors also stated that even though phosphate

does exchange with sulfate in the interlayer, the affinity for phosphate is not great enough

to lead to complete substitution. Nevertheless, the continued reaction observed here may

be attributed to phosphate reacting in the interlayer of green rust.

Sulfate release from green rust during the sorption experiments was between 416

and 625 µmols SO4/g green rust using initial phosphate concentrations between 10 and

1500 µM and was not directly correlated to PO4 (Figure 2.6). The NaCl used to adjust

the ionic strength of the phosphate solutions, however, had an appreciable influence on

sulfate release (Figure 2.5). Green rust allowed to react with anoxic distilled water for

two days released an average of 244.6 µmols mg SO4/g and that reacted with 0.1 M NaCl

released 631.9 µmols SO4/g. When a 100 µM NaH2PO4 solution (without the ionic

strength adjusted with NaCl) was used about 301.9 µmols SO4/g green rust was released.

Additionally, green rust reacted with 0.1 M NaCl and 1500 µM NaH2PO4 released the

same amount of sulfate per gram as green rust treated with only 0.1 M NaCl. The

Mineral Fe (mM) P (mM) SI* Fe(OH)2 SI* Vivianite SI* Wustite Magnetite A 0.0302 1.487 -3.996 2.803 -2.554 Magnetite B 0.0397 1.457 -3.874 3.150 -2.438 Green Rust A 0.0030 0.343 -4.857 -1.048 -3.369 Green Rust B 0.0030 0.296 -4.846 -1.226 -3.358 Siderite A 0.0140 1.467 -4.329 1.797 -2.868 Siderite B 0.0146 1.474 -4.311 1.854 -2.852

56

presence of chloride is not only the predominant factor in how much sulfate is released,

but also is a factor in how much phosphate will be sorbed. Green rust was reacted with

100 µM NaH2PO4 solutions with and without the ionic strength adjusted; the samples

without NaCl sorbed more phosphate than those reacted with a solution containing 0.1 M

NaCl (66.3 vs. 95.8 µmols PO4/g green rust, respectively). These observations indicate a

preference by interlayer sulfate to exchange with chloride over phosphate. This

preference may only be due to the concentration gradient betweens Cl- and PO43-. Lewis

(1997) found that chloride green rust would convert to sulfate green rust when reacted

with solutions containing less than 100 mM SO4; however, the study did not examine

whether the reverse reaction would occur.

Generally, a decrease in adsorption is seen with increasing pH; however, this

decrease is minimal for the ferric oxides within the pH range found in most soils (4 – 7)

(Brady and Weil, 1999). The trends displayed by the goethite- and ferrihydrite-coated

sand are in agreement with the findings of Geelhoed (1997), Strauss et al. (1997),

Goldberg (1983), Willett and Cunningham (1983), Ryden et al. (1977a), and McLaughlin

et al. (1977). The maximum sorption on goethite and ferrihydrite may not have been

observed and may actually occur at pH near 2.5 (Strauss et al., 1997). Both magnetite

samples and the natural siderite had well defined adsorption envelopes, with adsorption

rapidly diminishing between pH 8 and 9 (Figure 2.4). The ferrihydrite and ferrihydrite-

coated sand behave differently with increased pH. Adsorption onto the coated sand

decreases much more rapidly with pH than does sorption onto the ferrihydrite slurry. The

calculated binding coefficients were 0.145 for the ferrihydrite slurry and 0.042 for the

coated sand. Since the ferrihydrite slurry has a greater affinity for phosphate, one would

expect the changes in phosphate adsorption with pH to be less dramatic than those of the

ferrihydrite-coated sand.

The trends in decreasing adsorption with increasing pH will have implications on

the fate and transport of phosphate in natural systems. When phosphate is adsorbed to a

ferric hydroxide, two hydroxyl groups are released to solution, resulting in an initial

removal of phosphate from a system along with an increase in pH. This pH increase will

slow the reaction kinetics of sorption (Shang et al., 1992) and even completely inhibit

adsorption - as observed in the case of both magnetite samples and the geologic siderite.

57

Inhibition of phosphate sorption may occur at pH higher than what was examined in this

study for other iron oxides. Hingston et al. (1972) showed that phosphate sorption onto

goethite decreased to zero at pH 13. In terms of practical application, if a permeable

reactive barrier utilizing iron oxides were to be used for phosphate removal, one would

need to consider ways to control the pH. In a system that is not well buffered, the initial

phosphate adsorption will lead to increased pH and subsequently inhibited phosphate

removal. The retention would be slowed or completely halted until the pH decreased

again; in either case, the result would be ineffective phosphate removal. However, in

most soil systems the buffering capacity is likely to be great enough to mitigate large pH

changes upon P sorption.

The sorption characteristics of the siderite and magnetite samples varied

considerably between the geologic and synthetic specimens. In the case of the magnetite,

both exhibited a well-defined adsorption envelope, but the synthetic magnetite reached an

apparent adsorption maximum and was fit with a Langmuir isotherm; the geologic

sample did not exhibit an adsorption maximum and was fit to the Freundlich equation.

Adsorption onto synthetic siderite did not appear to change dramatically with pH;

however, sorption by the natural sample decreased to insignificant amounts at pH 9.

Similarly to the magnetite, no adsorption maximum on geologic siderite was observed

and was fit with the Freundlich equation, whereas phosphate adsorption onto synthetic

siderite conformed to the Langmuir equation. Additionally, dissolved iron was detected

in phosphate solutions following a two-day reaction period with both natural samples.

These variations may be the result of a variety of factors. One consideration is mineral

purity; the synthetic specimens are likely to have a much more consistent composition.

Geologic magnetite is commonly found to have small amounts of Ca2+, Mn2+, and Mg2+

replacing Fe2+, and Al3+ can replace Fe3+. Other, less common, replacement elements

include Cr3+ and V3+ (for Fe3+) and Ni2+, Co2+, and Zn2+ (for Fe2+). Manganese and

magnesium are both common substitutions for Fe2+ in siderite since there is complete

solid solution between siderite and rhodochrosite and siderite and magnesite (Deer et al.,

1992). Ainsworth et al. (1985) found that Al-substitution in goethite had an appreciable

influence on the extent and rate of isotopic phosphate exchange. The siderite and

magnetite samples were treated with dilute acid prior to the experiments; however,

58

residual oxidative coatings and inclusions may have influenced the sorption behavior.

Magnetite has commonly been observed to convert to maghemite and hematite (Cornell

and Schwertmann, 1994), and magnetite crystals are often observed to contain detectable

amounts of maghemite (Deer et al., 1992). Siderite usually will alter to goethite and

sometimes hematite (Deer et al., 1992).

In general, phosphate adsorption changed very little over the range of ionic

strengths examined. Sorption by natural magnetite and ferrihydrite (slurry) decreased

slightly with ionic strength and the remaining minerals, with the exception of green rust,

increased in sorption capacity with ionic strength. Green rust, discussed previously,

exhibited a substantial decrease in phosphate adsorption with increased chloride

concentration. Ionic strength alone is likely to have less influence on phosphate retention

than the pH and substrate; however, competition of other ions in solution for binding sites

is a factor that must be considered. Ryden et al. (1977b) found that increased ionic

strength resulted in an increase in phosphate sorption in soils and that the increase was

more pronounced with Ca present than with Na in solution. Additionally, pH may

influence how dramatically fluctuations in ionic strength affect phosphate sorption. Arai

and Sparks (2001) found that from pH 4 to 7.5, phosphate adsorption on ferrihydrite was

unaffected by changes in ionic strength; however, at pH greater than 7.5, adsorption

increased with ionic strength.

Patrick and Khalid (1974) postulated that the reductive dissolution of ferric iron

minerals led to ferrous compounds having increased activities and surface areas but that

the ferric oxides had a higher affinity for phosphate. These postulates are supported by

the findings of our work. On a surface area basis, all the ferrous iron minerals had much

higher capacities to sorb phosphate (from 15.16 to 343.27 µmols PO4/g mineral) than the

ferric minerals (1.68 to 6.32 µmols PO4/g mineral). This would result in the increased

phosphate sorption capacity of reduced sediments and soils in many areas provided high

surface area phases developed; however, the calculated binding and distribution

coefficients indicate that the ferric minerals have a higher affinity for phosphate than do

the ferrous minerals. The findings of this research also help to explain other previous

observations and have implications in how excess phosphate could be removed from a

system. Patrick and Khalid (1974) also noted that even though anaerobic soils were

59

capable of sorbing more phosphate they also released more into the soil solution, an

observation believed to result from lower affinity for phosphate by the ferrous minerals.

The distribution and binding coefficients calculated during this work support that theory

(Table 3). The adsorption characteristics exhibited by the ferric and ferrous iron minerals

also help to explain why seasonally flooded soils have been observed to have higher

levels of non-labile (non-exchangeable/iron bound) phosphate (Reynolds et al. 1999).

When flooded soils are subsequently aerated the relatively fine-grained ferrous minerals

would re-oxidize resulting in ferric phases with higher surface area and thus more

sorption sites than the ferric minerals present prior to flooding. The higher affinities of

the ferric iron minerals result in more strongly held (less labile) phosphate in the soil.

Changes in pH have been observed to affect phosphate retention by numerous ferric

oxides (Geelhoed, 1997; Strauss et al., 1997, Shang et al., 1992; Willett and Cunningham,

1983; Ryden et al., 1977b; McLaughlin et al., 1977). The adsorption envelopes observed

in this study indicate pH could play a much larger role in influencing phosphate retention

in sub-oxic to anoxic soils than in well-aerated systems.

CONCLUSIONS

The results of this study indicate ferrous minerals have phosphate adsorption

capacities that are much greater than those of the ferric phases examined. However, the

ferrous minerals all had much lower binding coefficients and were more sensitive to

changes in pH. The high capacity for phosphate of the ferrous minerals, combined with

the relatively low affinity and sensitivity to pH suggests that the ferrous minerals

investigated may provide appreciable, although transitory phosphorus sinks in some

reductomorphic environments.

60

CHAPTER 2

APPENDIX: ISOTHERM DATA COLLECTED DURING SORPTION

EXPERIMENTS

61

Table 2A.1. Summary of data used to establish adsorption isotherm of PO4 by naturally occurring (geologic origin) magnetite.

SamplePO4 solution

(L) g magnetite �M P (end) �M P (start)Change in �M P

�M P/g magnetite

mg PO4/g magnetite Avg mg/g SD mg/g SA m2/g

�M/m2

magnetite Avg �M/ m2 SD �M/m2

10A 0.099 0.097 10.144 10.330 0.186 0.189 0.018 0.025 0.006 0.132 1.430 1.956 0.47310B 0.099 0.109 10.025 10.330 0.305 0.276 0.026 0.132 2.09110C 0.098 0.117 9.960 10.330 0.370 0.310 0.029 0.132 2.34625A 0.099 0.094 23.255 23.928 0.673 0.707 0.067 0.061 0.026 0.132 5.356 4.891 2.04825B 0.100 0.099 23.055 23.928 0.873 0.880 0.084 0.132 6.66625C 0.099 0.098 23.581 23.928 0.347 0.350 0.033 0.132 2.65150A 0.100 0.096 47.685 48.363 0.678 0.703 0.067 0.045 0.022 0.132 5.328 3.555 1.72650B 0.098 0.090 47.944 48.363 0.420 0.457 0.043 0.132 3.45950C 0.100 0.091 48.137 48.363 0.226 0.248 0.024 0.132 1.879

100A 0.099 0.089 95.952 96.372 0.420 0.466 0.044 0.075 0.057 0.132 3.529 6.014 4.559100B 0.098 0.096 95.952 96.372 0.420 0.427 0.041 0.132 3.236100C 0.099 0.097 94.919 96.372 1.453 1.488 0.141 0.132 11.275200A 0.098 0.094 191.355 191.517 0.161 0.169 0.016 0.077 0.054 0.132 1.280 6.160 4.297200B 0.099 0.119 190.032 191.517 1.485 1.238 0.118 0.132 9.377200C 0.099 0.093 190.548 191.517 0.969 1.033 0.098 0.132 7.822500A 0.099 0.090 485.249 487.509 2.260 2.486 0.236 0.278 0.087 0.132 18.836 22.141 6.913500B 0.098 0.096 485.249 487.509 2.260 2.310 0.219 0.132 17.500500C 0.099 0.097 483.635 487.509 3.874 3.971 0.377 0.132 30.086

1000A 0.100 0.096 996.003 1001.492 5.489 5.727 0.544 0.877 0.310 0.132 43.383 69.930 24.7681000B 0.098 0.091 992.452 1001.492 9.040 9.767 0.928 0.132 73.9911000C 0.100 0.090 990.515 1001.492 10.977 12.199 1.159 0.132 92.4161500A 0.099 0.103 1499.655 1516.927 17.273 16.578 1.574 0.969 0.627 0.132 125.593 77.277 50.0331500B 0.099 0.098 1506.434 1516.927 10.493 10.633 1.010 0.132 80.5511500C 0.099 0.099 1513.537 1516.927 3.390 3.391 0.322 0.132 25.688

pH 4 A 0.096 0.035 0.132 7.650 2.801pH 4 B 0.099 0.100 94.822 95.581 0.759 0.748 0.071 0.132 5.669pH 4 C 0.098 0.086 94.467 95.581 1.114 1.271 0.121 0.132 9.631

pH 5.5 A 0.241 0.001 0.132 19.237 0.076pH 5.5 B 0.099 0.092 94.661 97.017 2.357 2.532 0.240 0.132 19.184pH 5.5 C 0.099 0.092 94.661 97.017 2.357 2.546 0.242 0.132 19.291pH 8 A 0.097 0.096 95.597 96.226 0.630 0.639 0.061 0.055 0.037 0.132 4.842 4.400 2.923pH 8 B 0.099 0.123 96.017 96.226 0.210 0.169 0.016 0.132 1.282pH 8 C 0.099 0.104 95.242 96.226 0.985 0.934 0.089 0.132 7.077pH 9 A 0.099 0.111 98.825 98.390 -0.436 -0.389 -0.037 -0.045 0.040 0.132 -2.945 -3.574 3.179pH 9 B 0.099 0.119 99.503 98.390 -1.114 -0.927 -0.088 0.132 -7.020pH 9 C 0.099 0.112 98.503 98.390 -0.113 -0.100 -0.009 0.132 -0.757

I/S 0.005 A 0.098 0.091 97.211 98.357 1.146 1.234 0.117 0.112 0.005 0.132 9.351 8.940 0.422I/S 0.005 B 0.098 0.114 97.050 98.357 1.308 1.123 0.107 0.132 8.507I/S 0.005 C 0.098 0.095 97.211 98.357 1.146 1.183 0.112 0.132 8.964

n/a

n/a

62

Table 2A.2. Summary of data used to establish adsorption isotherm of PO4 by sulfate green rust.

SamplePO4 solution

(L) g green rust �M P (end)�M P (start)

Change in �M P

�M P/g green rust

mg PO4/g green rust Avg mg/g SD mg/g SA m2/g

�M/m2

green rustAvg �M/

m2 SD �M/m2

10A 0.015 0.015 1.873 11.494 9.621 9.706 0.922 0.936 0.020 2.560 3.791 3.849 0.08110B 0.015 0.015 1.582 11.494 9.912 9.999 0.950 2.560 3.90625A 0.015 0.015 5.618 25.990 20.372 20.545 1.951 1.919 0.046 2.560 8.025 7.892 0.18825B 0.015 0.015 6.296 25.990 19.694 19.865 1.887 2.560 7.76050A 0.015 0.014 17.079 51.140 34.061 36.715 3.487 3.682 0.276 2.560 14.342 15.144 1.13550B 0.015 0.015 10.557 51.140 40.583 40.823 3.877 2.560 15.947

100A 0.015 0.014 36.967 100.843 63.877 68.905 6.544 6.315 0.324 2.560 26.916 25.975 1.331100B 0.015 0.014 41.551 100.843 59.292 64.086 6.086 2.560 25.034200A 0.015 0.015 75.354 199.749 124.395 125.109 11.882 11.563 0.450 2.560 48.871 47.561 1.851200B 0.015 0.016 74.095 199.749 125.655 118.406 11.245 2.560 46.252500A 0.015 0.015 234.198 506.557 272.359 273.757 25.999 29.252 4.601 2.560 106.936 120.318 18.924500B 0.015 0.014 189.031 506.557 317.526 342.271 32.505 2.560 133.700

1000A 0.015 0.013 533.677 1048.305 514.628 595.900 56.593 55.335 1.779 2.560 232.774 227.599 7.3171000B 0.015 0.015 481.052 1048.305 567.254 569.409 54.077 2.560 222.4251500A 0.015 0.013 739.657 1576.332 836.675 971.894 92.301 83.447 12.521 2.560 379.646 343.229 51.5021500B 0.015 0.015 794.542 1576.332 781.790 785.438 74.593 2.560 306.812

pH 4 A 0.015 0.014 23.601 100.004 76.403 82.139 7.801 7.748 0.075 2.560 32.086 31.867 0.309pH 4 B 0.015 0.013 30.122 100.004 69.882 81.020 7.694 2.560 31.648

pH 5.5 A 0.015 0.015 29.444 100.359 70.915 71.255 6.767 7.351 0.826 2.560 27.834 30.238 3.399pH 5.5 B 0.015 0.013 28.314 100.359 72.045 83.561 7.936 2.560 32.641pH 8 A 0.015 0.013 33.318 100.294 66.976 77.692 7.378 6.498 1.245 2.560 30.348 26.728 5.119pH 8 B 0.015 0.015 41.454 100.294 58.840 59.158 5.618 2.560 23.109pH 9 A 0.015 0.014 35.030 101.844 66.815 72.002 6.838 6.900 0.088 2.560 28.126 28.381 0.361pH 9 B 0.015 0.013 38.710 101.844 63.134 73.308 6.962 2.560 28.636

I/S 0.005 A 0.015 0.012 7.426 103.071 95.645 120.147 11.410 10.455 1.351 2.560 46.932 43.004 5.555I/S 0.005 B 0.015 0.014 9.686 103.071 93.385 100.036 9.500 2.560 39.076

63

Table 2A.3. Summary of data used to establish adsorption isotherm of PO4 by naturally occurring (geologic) siderite.

Sample PO4 solution

(L) g siderite �M P (end)�M P (start)

Change in �M P

�M P/g siderite

mg PO4/g siderite Avg mg/g SD mg/g SA m2/g

�M/m2

sideriteAvg �M/

m2 SD �M/m2

10A 0.100 0.145 10.160 10.330 0.169 0.116 0.011 0.076 0.062 0.270 0.431 2.977 2.40810B 0.100 0.097 8.962 10.330 1.367 1.409 0.134 0.270 5.21710C 0.100 0.093 9.505 10.330 0.825 0.886 0.084 0.270 3.28325A 0.099 0.097 22.635 23.928 1.293 1.321 0.125 0.081 0.055 0.270 4.892 3.164 2.16025B 0.099 0.106 23.714 23.928 0.215 0.200 0.019 0.270 0.74225C 0.100 0.115 22.726 23.928 1.203 1.042 0.099 0.270 3.85750A 0.098 0.099 46.007 48.363 2.357 2.341 0.222 0.220 0.006 0.270 8.672 8.588 0.21650B 0.100 0.102 45.942 48.363 2.421 2.362 0.224 0.270 8.74950C 0.098 0.104 45.974 48.363 2.389 2.252 0.214 0.270 8.342

100A 0.100 0.103 92.336 96.372 4.036 3.931 0.373 0.364 0.018 0.270 14.558 14.185 0.683100B 0.098 0.101 92.659 96.372 3.713 3.617 0.344 0.270 13.396100C 0.099 0.102 92.304 96.372 4.068 3.942 0.374 0.270 14.600200A 0.099 0.102 185.221 191.517 6.296 6.109 0.580 0.530 0.081 0.270 22.625 20.668 3.144200B 0.099 0.104 186.674 191.517 4.843 4.601 0.437 0.270 17.041200C 0.099 0.107 185.027 191.517 6.489 6.031 0.573 0.270 22.338500A 0.099 0.095 476.532 487.509 10.977 11.448 1.087 0.989 0.091 0.270 42.400 38.554 3.546500B 0.100 0.107 476.532 487.509 10.977 10.218 0.970 0.270 37.845500C 0.100 0.101 477.823 487.509 9.686 9.562 0.908 0.270 35.416

1000A 0.100 0.099 975.986 1001.492 25.505 25.803 2.450 2.187 0.228 0.270 95.566 85.309 8.8851000B 0.099 0.105 978.569 1001.492 22.923 21.708 2.062 0.270 80.3991000C 0.099 0.112 976.955 1001.492 24.537 21.590 2.050 0.270 79.9621500A 0.099 0.106 1485.449 1516.927 31.478 29.439 2.796 2.561 0.204 0.270 109.033 99.890 7.9521500B 0.100 0.100 1491.260 1516.927 25.667 25.539 2.425 0.270 94.5881500C 0.100 0.113 1487.709 1516.927 29.218 25.933 2.463 0.270 96.048

pH 4 A 0.099 0.110 92.175 95.581 3.406 3.051 0.290 0.307 0.049 0.270 11.298 11.978 1.908pH 4 B 0.098 0.099 92.724 95.581 2.857 2.836 0.269 0.270 10.503pH 4 C 0.099 0.115 91.142 95.581 4.439 3.816 0.362 0.270 14.133

pH 5.5 A 0.098 0.114 92.853 97.017 4.165 3.597 0.342 0.363 0.029 0.270 13.323 14.151 1.137pH 5.5 B 0.100 0.096 93.466 97.017 3.551 3.695 0.351 0.270 13.684pH 5.5 C 0.100 0.112 92.368 97.017 4.649 4.171 0.396 0.270 15.448pH 8 A 0.099 0.109 93.305 96.226 2.922 2.654 0.252 0.239 0.051 0.270 9.830 9.301 1.977pH 8 B 0.098 0.104 93.079 96.226 3.148 2.959 0.281 0.270 10.960pH 8 C 0.098 0.120 93.886 96.226 2.341 1.921 0.182 0.270 7.114pH 9 A 0.099 0.116 98.115 98.390 0.274 0.234 0.022 0.046 0.055 0.270 0.868 1.809 2.138pH 9 B 0.097 0.096 98.309 98.390 0.081 0.082 0.008 0.270 0.303pH 9 C 0.099 0.093 97.308 98.390 1.082 1.149 0.109 0.270 4.255

I/S 0.005 A 0.098 0.099 95.274 98.357 3.083 3.058 0.290 0.256 0.041 0.270 11.326 9.966 1.590I/S 0.005 B 0.099 0.112 95.855 98.357 2.502 2.219 0.211 0.270 8.219I/S 0.005 C 0.099 0.099 95.565 98.357 2.793 2.795 0.265 0.270 10.354

64

Table 2A.4. Summary of data used to establish adsorption isotherm of PO4 by ferrihydrite (as a slurry).

Sample PO4

solution (L) g HFO �M P (end)�M P (start)

Change in �M P

�M P/g HFO

mg PO4/g HFO Avg mg/g SD mg/g SA m2/g

�M/m2

HFO Avg �M/ m2 SD �M/m2

10A 0.098 0.088 0.010 9.458 9.448 11.622 1.104 1.108 0.006 200 0.058 0.058 0.00010B 0.099 0.088 0.000 9.458 9.458 11.733 1.114 200 0.05910C 0.098 0.088 0.000 9.458 9.458 11.648 1.106 200 0.05825A 0.098 0.088 0.000 21.526 21.526 26.495 2.516 2.531 0.023 200 0.132 0.133 0.00125B 0.100 0.088 0.000 21.526 21.526 26.926 2.557 200 0.13525C 0.098 0.088 0.000 21.526 21.526 26.520 2.519 200 0.13350A 0.099 0.088 0.349 43.350 43.002 53.145 5.047 5.094 0.051 200 0.266 0.268 0.00350B 0.100 0.088 0.000 43.350 43.350 54.200 5.147 200 0.27150C 0.099 0.088 0.000 43.350 43.350 53.554 5.086 200 0.268

100A 0.100 0.088 0.000 87.934 87.934 109.377 10.388 10.306 0.078 200 0.547 0.543 0.004100B 0.098 0.088 0.090 87.934 87.843 107.734 10.231 200 0.539100C 0.098 0.088 0.042 87.934 87.892 108.436 10.298 200 0.542200A 0.099 0.088 0.074 176.718 176.644 217.478 20.654 20.811 0.188 200 1.087 1.096 0.010200B 0.099 0.088 0.100 176.718 176.618 218.608 20.761 200 1.093200C 0.100 0.088 0.000 176.718 176.718 221.325 21.019 200 1.107500A 0.099 0.088 16.052 439.668 423.616 524.965 49.856 49.954 0.087 200 2.625 2.630 0.005500B 0.099 0.088 15.978 439.668 423.690 526.705 50.021 200 2.634500C 0.099 0.087 15.655 439.668 424.013 526.324 49.985 200 2.632

1000A 0.098 0.088 215.602 891.956 676.355 833.823 79.188 79.541 0.318 200 4.169 4.188 0.0171000B 0.098 0.087 212.567 891.956 679.389 840.311 79.804 200 4.2021000C 0.099 0.088 215.892 891.956 676.064 838.472 79.630 200 4.1921500A 0.099 0.088 527.865 1365.817 837.952 1037.376 98.520 98.536 0.101 200 5.187 5.188 0.0051500B 0.099 0.087 533.677 1365.817 832.140 1038.693 98.645 200 5.1931500C 0.098 0.087 532.708 1365.817 833.109 1036.580 98.444 200 5.183

pH 4 A 0.098 0.087 0.000 91.940 91.940 114.351 10.860 10.885 0.087 200 0.572 0.573 0.005pH 4 B 0.100 0.087 0.000 91.940 91.940 115.645 10.983 200 0.578pH 4 C 0.099 0.087 0.510 91.940 91.430 113.863 10.814 200 0.569

pH 5.5 A 0.099 0.087 0.000 87.347 87.347 108.736 10.327 10.240 0.085 200 0.544 0.539 0.004pH 5.5 B 0.098 0.087 0.000 87.347 87.347 107.806 10.238 200 0.539pH 5.5 C 0.097 0.087 0.000 87.347 87.347 106.942 10.156 200 0.535pH 8 A 0.098 0.087 0.000 88.110 88.110 109.698 10.418 10.347 0.065 200 0.548 0.545 0.003pH 8 B 0.098 0.088 0.116 88.110 87.993 108.789 10.332 200 0.544pH 8 C 0.097 0.087 0.000 88.110 88.110 108.353 10.290 200 0.542

I/S 0.005 A 0.099 0.088 0.039 90.384 90.346 111.881 10.625 10.649 0.029 200 0.559 0.561 0.002I/S 0.005 B 0.098 0.087 0.000 90.384 90.384 112.049 10.641 200 0.560I/S 0.005 C 0.100 0.088 0.000 90.384 90.384 112.468 10.681 200 0.562

65

Table 2A.5. Summary of data used to establish adsorption isotherm of PO4 by ferrihydrite-coated sand.

Sample PO4

solution (L)g HFO sand �M P (end)

�M P (start)

Change in �M P

�M P/g HFO sand

mg PO4/g HFO sand Avg mg/g SD mg/g SA m2/g

�M/m2

HFO sandAvg �M/

m2 SD �M/m2

10A 0.097 1.094 0.239 10.404 10.165 0.904 0.086 0.094 0.009 2.279 0.397 0.433 0.04110B 0.100 1.070 0.000 10.404 10.404 0.970 0.092 2.279 0.42510C 0.099 0.945 0.000 10.404 10.404 1.086 0.103 2.279 0.47725A 0.099 0.970 0.000 23.678 23.678 2.420 0.230 0.227 0.006 2.279 1.062 1.049 0.02725B 0.098 1.005 0.000 23.678 23.678 2.319 0.220 2.279 1.01725C 0.098 0.950 0.000 23.678 23.678 2.431 0.231 2.279 1.06750A 0.099 0.964 0.604 47.685 47.082 4.813 0.457 0.458 0.004 2.279 2.112 2.116 0.01750B 0.100 0.976 0.843 47.685 46.843 4.786 0.455 2.279 2.10050C 0.099 0.951 1.075 47.685 46.610 4.864 0.462 2.279 2.134

100A 0.098 0.968 26.045 96.727 70.682 7.126 0.677 0.690 0.019 2.279 3.127 3.189 0.087100B 0.099 0.949 24.760 96.727 71.967 7.494 0.712 2.279 3.288100C 0.100 0.980 26.435 96.727 70.292 7.182 0.682 2.279 3.151200A 0.097 0.913 99.729 194.390 94.661 10.035 0.953 0.934 0.017 2.279 4.403 4.313 0.079200B 0.100 1.014 95.080 194.390 99.310 9.761 0.927 2.279 4.283200C 0.099 1.025 93.918 194.390 100.472 9.695 0.921 2.279 4.254500A 0.099 0.952 367.730 483.635 115.904 12.023 1.142 1.140 0.003 2.279 5.275 5.267 0.014500B 0.097 1.009 359.013 483.635 124.621 12.020 1.142 2.279 5.274500C 0.100 1.072 355.462 483.635 128.173 11.966 1.136 2.279 5.251

1000A 0.098 0.982 849.428 981.152 131.724 13.195 1.253 1.215 0.042 2.279 5.790 5.612 0.1941000B 0.098 0.986 857.176 981.152 123.976 12.319 1.170 2.279 5.4061000C 0.099 0.989 852.656 981.152 128.496 12.855 1.221 2.279 5.6411500A 0.098 1.069 1349.851 1502.399 152.548 13.956 1.325 1.367 0.036 2.279 6.124 6.316 0.1671500B 0.099 1.059 1346.299 1502.399 156.100 14.601 1.387 2.279 6.4071500C 0.099 1.072 1343.393 1502.399 159.005 14.624 1.389 2.279 6.417

pH 4 A 0.097 0.993 2.486 101.134 98.648 9.685 0.920 0.921 0.014 2.279 4.250 4.255 0.064pH 4 B 0.099 0.990 5.786 101.134 95.348 9.556 0.908 2.279 4.193pH 4 C 0.099 0.949 6.299 101.134 94.835 9.849 0.935 2.279 4.322

pH 5.5 A 0.099 1.065 8.188 96.081 87.894 8.146 0.774 0.787 0.027 2.279 3.574 3.636 0.126pH 5.5 B 0.098 0.920 15.332 96.081 80.749 8.616 0.818 2.279 3.781pH 5.5 C 0.099 1.004 14.151 96.081 81.931 8.097 0.769 2.279 3.553pH 8 A 0.099 1.029 30.752 96.921 66.169 6.360 0.604 0.625 0.019 2.279 2.791 2.887 0.086pH 8 B 0.099 1.017 28.983 96.921 67.938 6.641 0.631 2.279 2.914pH 8 C 0.099 0.974 30.519 96.921 66.401 6.738 0.640 2.279 2.957

I/S 0.005 A 0.097 0.964 35.514 99.423 63.909 6.427 0.610 0.616 0.013 2.279 2.820 2.846 0.059I/S 0.005 B 0.097 0.978 34.997 99.423 64.425 6.391 0.607 2.279 2.805I/S 0.005 C 0.100 0.924 38.000 99.423 61.423 6.641 0.631 2.279 2.914

66

Table 2A.6. Summary of data used to establish adsorption isotherm of PO4 by goethite-coated sand.

SamplePO4 solution

(L)g goethite

sand�M P (end) �M P (start)

Change in �M P

�M P/g goethite

sand

mg PO4/g goethite

sandAvg mg/g SD mg/g SA m2/g

�M/m2

goethite sand

Avg �M/ m2 SD �M/m2

10A 0.100 0.972 0.216 9.492 9.276 0.954 0.091 0.086 0.007 2.394 0.398 0.380 0.03110B 0.099 0.987 0.000 9.492 9.492 0.951 0.090 2.394 0.39710C 0.100 1.131 0.161 9.492 9.330 0.823 0.078 2.394 0.34425A 0.100 1.206 0.142 23.266 23.124 1.914 0.182 0.199 0.027 2.394 0.800 0.874 0.11825B 0.099 0.923 0.749 23.266 22.517 2.419 0.230 2.394 1.01025C 0.100 1.177 0.449 23.266 22.818 1.946 0.185 2.394 0.81350A 0.100 1.032 12.233 48.751 36.518 3.535 0.336 0.340 0.026 2.394 1.477 1.496 0.11750B 0.099 1.230 7.203 48.751 41.548 3.328 0.316 2.394 1.39050C 0.098 0.971 10.502 48.751 38.248 3.880 0.368 2.394 1.621

100A 0.099 1.169 46.491 98.083 51.592 4.391 0.417 0.420 0.003 2.394 1.834 1.846 0.011100B 0.100 1.134 47.976 98.083 50.107 4.423 0.420 2.394 1.847100C 0.099 0.887 58.081 98.083 40.002 4.446 0.422 2.394 1.857200A 0.100 0.960 151.773 197.909 46.136 4.824 0.458 0.452 0.006 2.394 2.015 1.990 0.024200B 0.099 1.061 147.092 197.909 50.817 4.759 0.452 2.394 1.988200C 0.100 0.990 151.095 197.909 46.814 4.707 0.447 2.394 1.966500A 0.099 1.042 441.986 489.930 47.944 4.548 0.432 0.410 0.020 2.394 1.900 1.805 0.086500B 0.100 0.892 452.963 489.930 36.967 4.143 0.393 2.394 1.731500C 0.099 0.887 451.672 489.930 38.258 4.274 0.406 2.394 1.785

1000A 0.100 1.028 944.669 997.779 53.109 5.155 0.490 0.421 0.061 2.394 2.153 1.852 0.2701000B 0.099 1.021 957.584 997.779 40.195 3.907 0.371 2.394 1.6321000C 0.099 1.192 946.606 997.779 51.172 4.241 0.403 2.394 1.7711500A 0.099 1.006 1450.904 1501.915 51.011 5.013 0.476 0.449 0.034 2.394 2.094 1.975 0.1491500B 0.097 1.011 1457.038 1501.915 44.877 4.327 0.411 2.394 1.8071500C 0.100 1.052 1450.904 1501.915 51.011 4.841 0.460 2.394 2.022

pH 4 A 0.100 0.986 47.976 99.972 51.996 5.272 0.501 0.499 0.001 2.394 2.202 2.197 0.006pH 4 B 0.099 0.951 49.429 99.972 50.543 5.246 0.498 2.394 2.191pH 4 C 0.099 1.089 41.842 99.972 58.130 5.259 0.499 2.394 2.197

pH 5.5 A 0.099 1.020 50.301 99.584 49.284 4.770 0.453 0.459 0.006 2.394 1.993 2.018 0.027pH 5.5 B 0.099 0.986 51.301 99.584 48.283 4.827 0.458 2.394 2.016pH 5.5 C 0.099 0.990 50.849 99.584 48.735 4.897 0.465 2.394 2.046pH 8 A 0.099 1.066 57.855 99.600 41.745 3.878 0.368 0.370 0.009 2.394 1.620 1.629 0.040pH 8 B 0.099 1.055 58.759 99.600 40.841 3.816 0.362 2.394 1.594pH 8 C 0.100 0.884 64.183 99.600 35.417 4.004 0.380 2.394 1.672

I/S 0.005 A 0.099 1.051 59.179 103.814 44.634 4.200 0.399 0.385 0.014 2.394 1.754 1.691 0.063I/S 0.005 B 0.098 1.040 61.052 103.814 42.762 4.050 0.385 2.394 1.692I/S 0.005 C 0.098 0.817 71.318 103.814 32.495 3.897 0.370 2.394 1.628

67

Table 2A.7. Summary of data used to establish adsorption isotherm of PO4 by lepidocrocite.

SamplePO4 solution

(L) g lepidocrocite �M P (end) �M P (start)Change in �M P

�M P/g lepidocrocite

mg PO4/g lepidocrocite Avg mg/g SD mg/g SA m2/g

�M/m2

lepidocrociteAvg �M/

m2 SD �M/m2

10A 0.099 0.090 0.100 8.125 8.025 8.799 0.836 0.837 0.002 75.212 0.117 0.117 0.00010B 0.099 0.091 0.000 8.125 8.125 8.811 0.837 75.212 0.11710C 0.098 0.090 0.000 8.125 8.125 8.840 0.839 75.212 0.11825A 0.098 0.088 0.071 24.377 24.306 26.956 2.560 2.470 0.130 75.212 0.358 0.346 0.01825B 0.098 0.090 0.000 24.377 24.377 26.620 2.528 75.212 0.35425C 0.098 0.098 0.055 24.377 24.322 24.437 2.321 75.212 0.32550A 0.099 0.088 0.688 48.735 48.047 54.319 5.159 4.850 0.267 75.212 0.722 0.679 0.03750B 0.099 0.097 0.000 48.735 48.735 49.517 4.703 75.212 0.65850C 0.098 0.097 0.000 48.735 48.735 49.385 4.690 75.212 0.657

100A 0.098 0.089 13.444 97.324 83.881 92.493 8.784 8.643 0.235 75.212 1.230 1.210 0.033100B 0.098 0.099 8.494 97.324 88.830 88.145 8.371 75.212 1.172100C 0.098 0.090 12.404 97.324 84.920 92.372 8.773 75.212 1.228200A 0.099 0.095 90.141 195.633 105.492 109.506 10.400 10.294 0.141 75.212 1.456 1.441 0.020200B 0.099 0.104 82.973 195.633 112.660 106.704 10.134 75.212 1.419200C 0.098 0.098 87.138 195.633 108.495 108.977 10.349 75.212 1.449500A 0.098 0.091 373.219 483.957 110.739 119.626 11.361 11.187 0.281 75.212 1.591 1.566 0.039500B 0.099 0.091 373.864 483.957 110.093 119.376 11.337 75.212 1.587500C 0.099 0.100 368.053 483.957 115.904 114.384 10.863 75.212 1.521

1000A 0.098 0.093 883.004 995.519 112.514 118.456 11.250 11.371 0.344 75.212 1.575 1.592 0.0481000B 0.097 0.100 868.476 995.519 127.043 123.829 11.760 75.212 1.6461000C 0.098 0.095 882.036 995.519 113.483 116.926 11.104 75.212 1.5551500A 0.097 0.100 1384.396 1511.762 127.366 124.120 11.788 11.760 0.050 75.212 1.650 1.646 0.0071500B 0.098 0.108 1376.325 1511.762 135.437 123.219 11.702 75.212 1.6381500C 0.098 0.099 1386.979 1511.762 124.783 124.132 11.789 75.212 1.650

pH 4 A 0.099 0.106 0.839 97.615 96.775 90.065 8.553 8.773 0.525 75.212 1.197 1.228 0.073pH 4 B 0.098 0.095 2.076 97.615 95.539 98.677 9.371 75.212 1.312pH 4 C 0.099 0.109 0.504 97.615 97.111 88.380 8.393 75.212 1.175

pH 5.5 A 0.099 0.101 3.238 96.743 93.505 91.369 8.677 8.901 0.198 75.212 1.215 1.246 0.028pH 5.5 B 0.098 0.095 4.752 96.743 91.991 95.345 9.055 75.212 1.268pH 5.5 C 0.097 0.094 5.108 96.743 91.636 94.458 8.971 75.212 1.256pH 8 A 0.099 0.104 10.829 97.728 86.899 82.533 7.838 7.754 0.095 75.212 1.097 1.086 0.013pH 8 B 0.099 0.103 12.330 97.728 85.398 81.859 7.774 75.212 1.088pH 8 C 0.098 0.107 9.414 97.728 88.313 80.560 7.651 75.212 1.071

I/S 0.005 A 0.098 0.099 13.447 99.277 85.831 85.131 8.085 8.059 0.117 75.212 1.132 1.128 0.016I/S 0.005 B 0.099 0.101 11.258 99.277 88.020 85.926 8.160 75.212 1.142I/S 0.005 C 0.099 0.105 10.609 99.277 88.668 83.512 7.931 75.212 1.110

68

Table 2A.8. Summary of data used to establish adsorption isotherm of PO4 by synthetic siderite.

Sample PO4 solution (L) g siderite �M P (end)�M P (start)

Change in �M P

�M P/g siderite

mg PO4/g siderite Avg mg/g SD mg/g SA m2/g

�M/m2

sideriteAvg �M/

m2 SD �M/m2

10A 0.098 0.080 0.000 9.400 9.400 11.542 1.096 1.056 0.078 63.050 0.183 0.176 0.01310B 0.099 0.080 0.000 9.400 9.400 11.650 1.106 63.050 0.18510C 0.097 0.090 0.000 9.400 9.400 10.178 0.967 63.050 0.16125A 0.098 0.080 0.000 23.652 23.652 28.995 2.754 2.678 0.162 63.050 0.460 0.447 0.02725B 0.100 0.090 0.000 23.652 23.652 26.241 2.492 63.050 0.41625C 0.099 0.080 0.000 23.652 23.652 29.370 2.789 63.050 0.46650A 0.098 0.080 0.000 47.750 47.750 58.470 5.553 5.352 0.369 63.050 0.927 0.894 0.06250B 0.098 0.090 0.000 47.750 47.750 51.872 4.926 63.050 0.82350C 0.098 0.080 0.000 47.750 47.750 58.715 5.576 63.050 0.931

100A 0.099 0.070 1.521 95.952 94.431 133.944 12.721 11.636 0.941 63.050 2.124 1.943 0.157100B 0.098 0.080 1.414 95.952 94.538 116.258 11.041 63.050 1.844100C 0.098 0.080 0.381 95.952 95.571 117.373 11.147 63.050 1.862200A 0.097 0.090 11.839 188.837 176.998 190.410 18.083 17.743 0.565 63.050 3.020 2.963 0.094200B 0.098 0.090 24.166 188.837 164.671 179.968 17.092 63.050 2.854200C 0.098 0.080 33.512 188.837 155.325 190.118 18.055 63.050 3.015500A 0.098 0.090 30.532 464.586 434.054 472.251 44.850 46.492 3.302 63.050 7.490 7.764 0.551500B 0.098 0.090 34.997 464.586 429.589 466.820 44.334 63.050 7.404500C 0.098 0.080 33.028 464.586 431.558 529.576 50.294 63.050 8.399

1000A 0.098 0.090 196.263 921.101 724.838 787.497 74.789 74.225 1.053 63.050 12.490 12.396 0.1761000B 0.099 0.090 223.059 921.101 698.042 768.776 73.011 63.050 12.1931000C 0.100 0.090 210.178 921.101 710.923 788.414 74.876 63.050 12.5051500A 0.098 0.090 508.817 1402.476 893.659 977.067 92.792 90.937 2.038 63.050 15.497 15.187 0.3401500B 0.099 0.100 457.806 1402.476 944.669 934.561 88.755 63.050 14.823

pH 4 A 0.099 0.090 0.672 95.113 94.441 103.927 9.870 9.496 0.625 63.050 1.648 1.586 0.104pH 4 B 0.098 0.100 0.455 95.113 94.657 92.386 8.774 63.050 1.465pH 4 C 0.099 0.090 0.517 95.113 94.596 103.646 9.843 63.050 1.644

pH 5.5 A 0.098 0.100 0.497 95.678 95.180 93.467 8.877 10.348 1.275 63.050 1.482 1.728 0.213pH 5.5 B 0.099 0.080 1.282 95.678 94.396 116.355 11.050 63.050 1.845pH 5.5 C 0.099 0.080 1.385 95.678 94.293 117.064 11.118 63.050 1.857pH 8 A 0.098 0.100 0.539 94.951 94.412 92.816 8.815 9.128 0.562 63.050 1.472 1.524 0.094pH 8 B 0.098 0.100 0.313 94.951 94.638 92.584 8.793 63.050 1.468pH 8 C 0.099 0.090 1.094 94.951 93.857 102.940 9.776 63.050 1.633pH 9 A 0.097 0.093 0.000 98.390 98.390 102.428 9.728 9.698 0.534 63.050 1.625 1.620 0.089pH 9 B 0.099 0.101 0.000 98.390 98.390 96.349 9.150 63.050 1.528pH 9 C 0.097 0.089 0.149 98.390 98.241 107.576 10.217 63.050 1.706

I/S 0.005 A 0.097 0.080 0.000 99.988 99.988 121.773 11.565 10.387 1.136 63.050 1.931 1.735 0.190I/S 0.005 B 0.098 0.090 0.000 99.988 99.988 108.431 10.298 63.050 1.720I/S 0.005 C 0.098 0.100 0.000 99.988 99.988 97.908 9.298 63.050 1.553

69

Table 2A.9. Summary of data used to establish adsorption isotherm of PO4 by naturally sulfate green rust.

SamplePO 4 solution

(L) g magnetite�M P (end)

�M P (start)

Change in �M P

�M P/g magnetite

mg PO 4/g magnetite Avg mg/g SD mg/g SA m2/g

�M /m 2

magnetiteAvg �M /

m 2 SD �M /m 2

10A 0.099 0.107 0.978 9.648 8.670 7.997 0.760 0.687 0.067 1.743 4.588 4.151 0.40510B 0.100 0.146 0.000 9.648 9.648 6.602 0.627 1.743 3.78810C 0.098 0.133 0.000 9.648 9.648 7.107 0.675 1.743 4.07825A 0.100 0.107 3.203 25.152 21.949 20.486 1.946 1.793 0.169 1.743 11.753 10.829 1.02325B 0.099 0.116 2.683 25.152 22.469 19.179 1.821 1.743 11.00425C 0.101 0.138 1.866 25.152 23.286 16.961 1.611 1.743 9.73150A 0.099 0.220 15.365 48.428 33.063 14.869 1.412 1.965 0.487 1.743 8.531 11.872 2.94350B 0.099 0.135 17.602 48.428 30.826 22.669 2.153 1.743 13.00650C 0.100 0.125 17.738 48.428 30.690 24.538 2.330 1.743 14.078

100A 0.098 0.154 53.691 99.826 46.136 29.278 2.780 2.674 0.110 1.743 16.797 16.157 0.663100B 0.099 0.138 62.117 99.826 37.709 26.969 2.561 1.743 15.473100C 0.100 0.153 56.790 99.826 43.036 28.236 2.682 1.743 16.200200A 0.098 0.133 157.230 197.489 40.260 29.643 2.815 2.704 0.111 1.743 17.007 16.335 0.672200B 0.098 0.144 155.841 197.489 41.648 28.473 2.704 1.743 16.335200C 0.100 0.138 159.748 197.489 37.742 27.300 2.593 1.743 15.663500A 0.099 0.155 446.184 488.477 42.294 27.119 2.575 2.583 0.148 1.743 15.559 15.601 0.894500B 0.099 0.165 445.861 488.477 42.617 25.673 2.438 1.743 14.729500C 0.099 0.178 436.821 488.477 51.657 28.788 2.734 1.743 16.516

1000A 0.099 0.148 939.181 989.546 50.365 33.551 3.186 3.168 0.360 1.743 19.249 19.138 2.1721000B 0.100 0.179 936.598 989.546 52.948 29.478 2.800 1.743 16.9121000C 0.098 0.109 948.221 989.546 41.325 37.044 3.518 1.743 21.2531500A 0.099 0.142 1487.709 1542.917 55.208 38.623 3.668 3.282 0.378 1.743 22.159 19.830 2.2851500B 0.100 0.101 1511.923 1542.917 30.994 30.661 2.912 1.743 17.5911500C 0.101 0.158 1489.000 1542.917 53.917 34.407 3.268 1.743 19.740

pH 4 A 0.100 0.106 69.478 97.631 28.153 26.656 2.531 2.564 0.032 1.743 15.293 15.490 0.191pH 4 B 0.099 0.119 65.184 97.631 32.447 27.022 2.566 1.743 15.503pH 4 C 0.099 0.107 68.154 97.631 29.477 27.320 2.595 1.743 15.674

pH 5.5 A 0.100 0.131 63.086 96.775 33.690 25.845 2.454 2.540 0.139 1.743 14.828 15.346 0.841pH 5.5 B 0.101 0.143 56.370 96.775 40.405 28.440 2.701 1.743 16.316pH 5.5 C 0.100 0.107 69.123 96.775 27.652 25.961 2.466 1.743 14.895pH 8 A 0.100 0.090 78.163 98.551 20.388 22.586 2.145 2.143 0.009 1.743 12.958 12.945 0.054pH 8 B 0.099 0.107 74.224 98.551 24.327 22.461 2.133 1.743 12.886pH 8 C 0.099 0.105 74.450 98.551 24.101 22.644 2.150 1.743 12.991

pH 8 A2 0.099 0.163 60.470 100.311 39.840 24.245 2.303 2.267 0.156 1.743 13.910 13.697 0.943pH 8 B2 0.099 0.100 74.870 100.311 25.441 25.300 2.403 1.743 14.515pH 8 C2 0.098 0.108 76.000 100.311 24.311 22.076 2.097 1.743 12.666pH 9 A 0.099 0.095 98.115 98.390 0.274 0.286 0.027 0.047 0.052 1.743 0.164 0.283 0.314pH 9 B 0.099 0.099 98.309 98.390 0.081 0.081 0.008 1.743 0.046pH 9 C 0.099 0.096 97.308 98.390 1.082 1.115 0.106 1.743 0.640

IS 0.005 A 0.098 0.126 66.863 102.393 35.530 27.644 2.625 2.540 0.084 1.743 15.860 15.343 0.510IS 0.005 B 0.099 0.099 76.484 102.393 25.909 25.869 2.457 1.743 14.841IS 0.005 C 0.098 0.145 62.795 102.393 39.598 26.718 2.537 1.743 15.329

71

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Strauss, R., G.W. Bruemmer, and N.J. Barrow. 1997. Effects of crystallinity of goethite

II: Rates of sorption and desorption of phosphate. Eur. J. Soil Sci. 48:101-114. Stumm, W., and J.J. Morgan. 1981. Aquatic Chemisty: An Introduction Emphasizing

Chemical Equilibria in Natural Waters John Wiley & Sons, Toronto. Torrent, J., V. Barron, and U. Schwertmann. 1990. Phosphate adsorption and desorption

by goethites differing in crystal morphology. Soil Sci. Soc. Am. J. 54:1007-1012. Torrent, J., U. Schwertmann, and V. Barron. 1994. Phosphate sorption by natural

hematites. Eur. J. Soil Sci. 45:45-51. Torrent, J. 1997. Interactions between phosphate and iron oxide. Adv. GeoEcology

30:321-344. Willett, I.R., and R.B. Cunningham. 1983. Influence of sorbed phosphate on the stability

of ferric hydrous oxide under controlled pH and Eh conditions. Aust. J. Soil Res. 21:301-308.

Williams, J.D.H., J.K. Syers, S.S. Shukla, and R.F. Harris. 1971. Levels of inorganic and

total phosphorus in lake sediments as related to other sediment parameters. Environ. Sci. Technol. 5:1113-1120.

Zachara, J.M., J.K. Fredrickson, S.M. Li, D.W. Kennedy, S.C. Smith, and P.L. Gassman.

1998. Bacterial reduction of crystalline Fe3+ oxides in single phase suspensions and subsurface materials. Am. Mineral. 83:1426-1443.

75

CHAPTER 3

Phosphate Adsorption by Heterogeneous Mixtures of Ferrous and Ferric Oxides Under Hydrodynamic Conditions

INTRODUCTION

I. Importance and cycling of phosphate in the natural environment

Phosphate is an essential component of ATP (adenosine tri-phosphate), DNA

(deoxyribonucleic acid), RNA (ribonucleic acid), and phospholipids as a result of its

unique chemical properties. The critical roles occupied by phosphate in combination

with relatively low availability leads to very rapid biological cycling and a high level of

phosphate conservation in undisturbed ecosystems (Schlesinger, 1991). Freshwater and

estuarine ecosystems are often nutrient limited and are particularly sensitive to increases

in phosphate (Reddy et al., 1999; Manahan, 1994). While phosphate is not directly toxic

to aquatic organisms, eutrophication and toxic algal blooms do cause fish kills.

Eutrophication occurs when a limiting nutrient is added to a water body. The resulting

algal blooms consume oxygen and prevent light infiltration. As the algae begin to die

and decompose the biological oxygen demand spikes, resulting in anoxia (Brady and

Weil, 1999). Recent studies in the Everglades document that small additions of

phosphate, due to increased runoff from agriculture, appear to be altering the ecology of

the area (DeBusk et al., 2001; Noe et al., 2001; DeBusk et al., 1994).

The prominent correlation between phosphorus and iron minerals in estuaries

(Pant and Reddy, 2001; Roden and Edmunds, 1997), lake and ocean sediments (Young

and Ross, 2001; Slomp et al., 1998; Olila and Reddy, 1997; Olila and Reddy, 1995;

Moore and Reddy, 1994), and soils (Golterman, 1999; DeMello et al., 1998; Szilas et al.,

1998; Willett and Higgens, 1978) has led to numerous examinations of P adsorption

capacities and mechanisms by iron oxides and hydroxides. The most commonly studied

ferric minerals have been hematite and goethite since they are the most stable and most

abundant, although not necessarily the most reactive, iron minerals found in soils

(Cornell and Schwertmann, 1996; Sparks, 1995; Torrent et al., 1994). The mechanism of

phosphate adsorption onto ferric oxides is dominated by ligand exchange in which two

singly coordinated hydroxyl groups or water molecules are replaced by a single

phosphate anion resulting in the formation of a bidentate, binuclear complex (Reddy et

76

al., 1999; Torrent, 1997; Colombo et al., 1994; Goldberg and Sposito, 1985; Parfitt and

Russell, 1977; Parfitt et al., 1975). The phosphate surface complexes are very stable and

result in slow exchange rates and an apparent irreversibility (hysterisis) of phosphorus

adsorption. The strength of these complexes leads to long-term phosphorus storage in

soils, sediments, and wetlands (Reddy et al., 1999; Goldberg and Sposito, 1985; Parfitt et

al., 1975). Phosphate sorption by soils and pure mineral phases occurs in two distinct

steps - rapid adsorption taking place during the first few minutes of the reaction followed

by a slower stage (Reddy et al., 1999; Strauss et al., 1997; McLaughlin et al., 1977).

Slow adsorption increases with specific surface area, porosity, and ferrihydrite impurities

(Strauss et al., 1997; Torrent 1997; Torrent et al., 1994), and it has also been attributed to

diffusion of phosphate into the structural pores and defects of soil minerals (Ryden et al.

1977a). The slow sorption step has also been associated with the irreversibility of

phosphate retention by soils and pure iron oxides (Madrid and DeArambarri, 1985;

Ryden and Syers, 1977). The phosphate adsorption capacity of iron oxides generally

averages 2.5 µM P/m2, but ranges from 1.5 to 3.5 µM/m2 (Reddy et al., 1999; Torrent

1997; Enyard, 1994; Torrent et al., 1990; Borggaard, 1983

Solution pH will determine the extent, rate, and mechanism of adsorption by iron

oxides. In general, it appears that phosphorus sorption by ferric oxides decreases with

increasing pH (Geelhoed, 1997; Willett and Cunningham, 1983; Ryden et al., 1977b;

McLaughlin et al., 1977; and Hingston et al., 1972). Strauss et al. (1997) and Shang et al.

(1992) noted that pH was a particularly important control on P adsorption kinetics during

the initial rapid reaction. Ionic strength and competition with other ions also influence

phosphate retention. Ryden et al. (1977b) and McLaughlin et al. (1977) found P

adsorption onto ferric hydroxide gel to increase with ionic strength. Ryden et al. (1977b)

also noted that the ionic strength effects on chemisorption were predominantly kinetic

and that the effects on physical adsorption were absolute. Arai and Sparks (2001) found

that at pH greater than 7.5, P adsorption by ferrihydrite increased with ionic strength.

Celi et al. (2000) observed increased phosphorus sorption by goethite with ionic strength

in agreement with Ryden et al. (1977b).

Phosphate transport is a concern in many areas, particularly locations where

sewage is primarily treated using septic system infiltration beds. Since P is a limiting

77

nutrient in most freshwater systems understanding the factors influencing P retention,

such as adsorption and mineral precipitation, are important considerations (Robertson et

al., 1998; Zanini et al., 1998; Robertson, 1995). Additionally, understanding the factors

inhibiting phosphate adsorption is necessary to predict the fate of P and for the design

and operation of remediation systems.

Walter et al. (1999) examined P transport in a glacial aquifer on Cape Cod; excess

P was derived from sewage disposal on a military base. They noted that P adsorption by

Al and Fe oxides has greatly retarded the movement of the plume relative to groundwater

flow rates. Phosphorus sorbs very strongly to surfaces of Fe and Al oxides, which

commonly coat sediment grains in glacial aquifers, thus retarding P transport; in many

areas, 90 to 99% of the P in a given volume of aquifer is contained on sediment surfaces.

Even though sorption processes are inhibiting P transport, continued loading is creating a

large P reservoir that might be released later. Increased loading also allows for P

breakthrough and inhibits retention. It is possible, however, that P desorption may occur

with the influx of fresh water.

Stollenwerk (1996) simulated P transport through the Cape Cod aquifer and

observed that the amount of P in solution is dependant upon P concentration and pH. He

also observed that even though P input had discontinued, a large reservoir remained in

the sediments, which could desorb later and affect nearby water bodies. During column

experiments, higher feed concentration resulted in more rapid saturation of adsorption

sites and earlier breakthrough. At lower concentrations, the breakthrough curve was

more gradual and complete saturation took longer.

P transport has also been observed in oxic and suboxic conditions. Oxic and sub

oxic sediments have the capacity to completely sorb inflowing P for long periods of time;

however this is dependant on pH and upon how long the column has been loaded (or how

near the adsorption capacity it is) (Walter et al., 1999).

Phosphorus transport is commonly associated with the reductive dissolution of

Fe(III) (hydr)oxides (Jensen et al., 1999; Szilas et al., 1998; Olila and Reddy, 1997).

Jensen et al. (1999) noted two steps of transport for reactive solutes through structured

soils. The initial loading step involves the transfer from source (solid phase) to solution

(mobile or percolating phase). The second step is translocation, or transport with the

78

mobile phase. They postulated that P associated with, Fe(III), a redox sensitive species,

under reducing conditions is released, and if surrounding solution is mobile, the initial

step is favored. During reductive dissolution the released Fe(II) was resorbed; however,

soil had less affinity for iron than phosphate.

II. Need and Research Objectives

The goal of this study is to specifically evaluate the interaction of iron and

phosphate in several controlled, dynamic systems. Since iron is a predominant sorbent of

phosphate in many systems (Pant and Reddy, 2001; Young and Ross, 2001; Golterman,

1999; DeMello et al., 1998; Slomp et al., 1998; Szilas et al., 1998; Olila and Reddy,

1997; Roden and Edmunds, 1997; Olila and Reddy, 1995; Moore and Reddy, 1994;

Willett and Higgens, 1978), the results from this study will help determine how nutrients,

specifically phosphate, interact with iron phases in dynamic, suboxic environments. This

research will also help to characterize interaction between iron and phosphate in riparian

and groundwater systems, which will lead to a better understanding of the fate and

transport of phosphate in the freshwater environment. These dynamic experiments are

necessary in order to evaluate phosphate retention under more realistic conditions. The

specific objectives of this project were to compare the phosphate adsorption of minerals

in static and flow conditions. The retention time and capacity of the columns were also

examined; in addition to assessing which iron phases were responsible for phosphate

retention within the columns. These experiments were also used to determine what other

factors are likely to influence phosphate retention in hydrodynamic systems.

A total of six hydrodynamic column experiments were run; three were anoxic and

three were oxic. The three oxic columns all contained only one iron phase, ferrihydrite-

coated sand, and two of the columns were replicates that were poorly buffered while the

third column was well buffered. The three anoxic columns consisted of heterogeneous

mixtures of iron oxides and hydroxides. In two of the anaerobic columns, the iron

reducing bacterium Shewanella putrefaciens (strain CN 32) was used to facilitate the

conversion of ferrihydrite to magnetite, lepidocrocite, and goethite. Ferrous chloride was

used to drive magnetite precipitation, as well as conversion of ferrihydrite to goethite and

lepidocrocite, in the third anoxic column.

79

MATERIALS AND METHODS

A total of six column experiments (ranging in complexity) were conducted to

evaluate the influence of iron oxides and hydroxides on phosphate retention and

transport. Three experiments utilized only ferrihydrite-coated sand and were carried out

under aerobic conditions at ambient temperature. Two of these columns were poorly

buffered while the third was buffered at pH 7 using MOPS. The remaining three

experiments were conducted in an anaerobic glovebox (Coy Products, Ann Arbor, MI)

having O2(g) levels less than 5 ppm. The first of the anaerobic experiments examined

phosphate transport through a FeCl2 reacted ferrihydrite column in which magnetite was

generated as a principal sorbent. The two remaining anaerobic columns, each with a

different feed solution composition, were inoculated with an iron reducing bacterium to

facilitate production of mixed and reduced iron phases such as siderite, magnetite, and

green rust, as well as conversion of ferrihydrite to goethite and lepidocrocite.

I. Ferrihydrite-coated sand preparation

Ferrihydrite (2-line) was made using the method described by Cornell and

Schwertmann (1996). A ferric chloride solution (2.3 L of 0.062 M FeCl36H2O, pH near

2) was titrated with 0.2 M NaOH to pH 7. The solution was then stirred and allowed to

stabilize for several hours, with pH maintained by base addition. Once the solution pH

was stable, the precipitate was allowed to settle and rinsed three times with deionized

water. Ferrihydrite coated sand was made by drying the slurry onto 660 g of Iota 6

Quartz Sand (Umimin Corporation). The sand was dried at room temperature and was

rinsed several times with and stored in deionized water.

II. General column set-up and sampling

Columns, made from acrylic, had dimensions of 25 cm length and 3.5 cm

diameter, with eight sampling ports located along the length of the column (Figure 3.1).

Sampling ports consisted of 16-guage needles held in place by Swage-Loc® fittings.

Prior to insertion, needles were blunted and filled with glass fiber to prevent sand from

escaping. A peristaltic pump was used to introduce feed solution to the bottom of the

column. Effluent was collected in a sampling vessel and excess solution was diverted to

80

a waste container (Figure 3.2). Influent, effluent, and side port samples were collected

and monitored for parameters specific to each column. Side port solution was collected

by closing the valve on the effluent end of the column and directing liquid into a test tube

or syringe.

III. Aerobic columns (ferrihydrite-coated sand only)

During all three aerobic flow experiments the columns were loaded such that

ferrihydrite-coated sand was situated between uncoated silica sand. Coated sand filled

the column between ports 5 through 7, a total length of ~7.5 cm. The remainder of the

column was packed with uncoated silica sand. The large amount of uncoated sand helped

to evenly disperse solution introduced to the column.

Two aerobic column experiments utilized a feed solution containing 100 µM

NaH2PO4 adjusted to pH 7 using NaOH. A third aerobic column was different only in

that the 100 µM NaH2PO4 solution was buffered (pH 7) using 10 mM MOPS (3-(N-

Morpholino) propane sulfonic acid). In all three aerobic experiments, a flow rate of ~0.4

mL/min was used. Flow was discontinued once complete, or near complete (>95%),

phosphate breakthrough was achieved. Initially, feed, effluent and port samples were

collected daily to measure pH and determine phosphate concentration. After 20 days,

sampling was conducted every other day. Samples for elemental analysis were

preserved with concentrated HCl and refrigerated until processed.

Upon experiment termination aliquots, of coated sand were reacted with 6 M HCl

overnight. The extract was diluted and analyzed for P and Fe.

81

Figure 3.1. Schematic diagram of column apparatus used to evaluate PO4

3- transport through packed mineral beds. Ports were numbered 1 (bottom) through 8 (top).

Sampling ports 25 cm 3.5 cm

valve

valve

flow

82

IV. Anaerobic columns (multiple iron phases coating sand)

An initial anaerobic column experiment was conducted involving the abiotic synthesis of

magnetite. A 2 mM FeCl2 solution buffered (pH 7) with 10 mM PIPES (Piperazine-

N,N'-bis-[2-ethanesulfonic acid]) was used as the influent solution into ferrihydrite-

coated sand for the first 120 h to generate magnetite. After generation of magnetite, the

FeCl2 feed was discontinued and the column flushed with deionized water for 20 minutes.

A 100 µM NaH2PO4 solution adjusted to pH 7 (with NaOH) was then introduced and

flow continued until complete phosphate breakthrough was observed. A flow rate of

~0.45 mL/min was maintained throughout the experiment. Influent, effluent, and side

port samples were collected daily for Fe, P, and pH measurement.

When complete breakthrough was observed the column was disassembled and divided

based on obvious color differences. Four aliquots from each section were reacted with 6

M HCl overnight (total extraction). Two of these were used to determine Fe(II) by the

ferrozine method and the other two were used for total Fe and P analysis. Surface area

and solid phase analysis, using extended x-ray absorption fine structure spectroscopy,

were also performed.

Two additional anaerobic columns were used to evaluate phosphate retention

dynamics during the biologically induced reductive dissolution of ferric hydroxide. The

columns were packed so uncoated sand at ports 1 and 8 bracketed ferrihydrite-coated

sand. The coated sand was inoculated with the iron reducing bacterium Shewanella

putrefaciens (strain CN32) during the loading process (see Benner et al., 2002). A

solution representative of groundwater was used on one column and a more heavily

bicarbonate-buffered feed used on the other. Both feed solutions contained ~3.2 mM

lactate as an electron donor. Average feed compositions are listed in Table 3.1. These

two columns are referred to as GW (groundwater) and BC (bicarbonate buffered)

throughout the discussion. The flow velocity was equivalent to 0.5 m/day and the

columns were run for 50 d.

Influent, effluent, and side port samples were collected daily. Alkalinity and pH were

measured by standard procedures. Ion chromatography was used to analyze samples for

lactate, acetate, Cl-, and SO42-. Ferrous iron was determined using the ferrozine method,

83

and elemental concentrations of Ca, Fe, K, Mg, Na, P, S, and Si were measured by ICP-

OES.

Upon termination, material from the packed mineral bed was removed and

divided according to position within the column relative to the side ports. Sand was

homogenized and split into aliquots for bacterial cell counts, EXAFS analysis, total (6 M

HCl) and partial (0.5 M HCl) acid extractions, and wet/dry ratios. Acid extracts were

diluted and analyzed for Ca, Fe, K, Mg, Na, P, S, and Si using ICP-OES and Fe2+ using

the ferrozine method.

V. Static adsorption experiments (on sand)

Static adsorption experiments were conducted on the ferrihydrite-coated sand in

order to determine the adsorption maximum. Solutions used ranged from 500 µM to

1500 µM Na2HPO4. The adsorption maxima for both batches of sand were determined

with solutions having the ionic strength adjusted to 0.1 M with NaCl as well as with

solutions containing only Na2HPO4. The pH was adjusted to 7 using 0.05 M HCl or 0.05

M NaOH. All samples were run in triplicate. Upon completion, 8 mL of solution were

filtered through a 0.45µm filter and acidified using two drops of HNO3. Starting and

ending solutions were analyzed using inductively coupled plasma-optical emission

spectroscopy (ICP-OES). The amount of phosphate adsorbed was determined from the

difference between the initial and final concentrations.

VI. Surface Area Analysis

Specific surface area was determined by N2 adsorption, using the BET equation

(Brunauer et al., 1938). The adsorption curves were obtained using a Coulter SA 3100

surface area analyzer. Samples were degassed by heating for 3 h at 85˚C. Contact with

the atmosphere was avoided by maintaining positive pressure (helium) on the sample

containers.

84

Table 3.1. Average feed solution compositions used during biotic column experiments.

GW Feed BC Feed Parameter Average (mM) Average (mM)

pH 7 7.1 Alkalinity 2.8 meq 9 meq

Ca 2.9 0.73 Fe 0.00 0.002 K 0.09 1.9

Mg 0.42 1.1 Na 1.3 15 P 0.007 0.26 S 0.43 1.1 Si 0.04 0.03 Cl 0.08 51

SO4 0.04 1 Lactate 3.2 3.2

85

RESULTS

I. Aerobic columns

Incomplete phosphate breakthrough resulted in both columns even after 120 h of

reaction and was correlated to a spike in effluent pH (Figure 3.2). After the initial

phosphate pulse, the effluent exhibits several plateaus that appear to be related to pH

fluctuations. Nearly complete phosphate breakthrough was observed in the two poorly

buffered columns (PB1 and PB2) after ~900 h. Upon termination, effluent phosphate

concentration was approximately 95% of the influent. Data collected from the side ports

reflect the trends seen in the effluent; phosphate concentration in the side ports drops

rapidly shortly after flow began (23 h) (Figure 3.3). The pH initially drops when flow is

started but quickly increases to near 9.5 (Figure 3.3).

The extraction data collected upon termination are shown in Figure 3.4. Overall,

there was little difference in iron and phosphate concentrations throughout the columns.

Phosphate is partitioned evenly and the ratio of iron to phosphate ranges from 5.2 to 6.7,

averaging 5.8 and 5.6 for PB1 and PB2, respectively.

The well-buffered experiment yielded a breakthrough curve with markedly

different characteristics than the poorly-buffered columns. An initial breakthrough of

phosphate occurred at approximately 260 h (Figure 3.2). While the breakthrough was

not complete at this time, the early breakthrough was much more rapid than in the PB

columns. Between 240 and 330 h, the curve is quite steep; however, after 330 h the rate

of breakthrough decreases and upon termination at 500 h the phosphate concentration of

the effluent is ~91% of the influent. The pH did not vary more than 0.1 units throughout

this experiment. Side-port phosphate concentrations (Figure 3.5) also exhibit different

behavior than the poorly buffered columns. In the influent portions of the column,

phosphate is rapidly removed from solution and is reflected by a retarded phosphate front

moving downfield. The total extractions (Figure 3.4) yield very similar results as

observed in the two poorly buffered columns. The iron to phosphate ratio is between 6.1

and 6.5 throughout the ferrihydrite-coated sand.

86

Time (h)

0 100 200 300 400 500 600

pH

3

4

5

6

7

8

mM

P

0.0

0.1

0.2

0.3

0.4

0.5

pH (in) pH (out) P (in) P (out)

Time (h)

0 200 400 600 800 1000 1200

pH

3

4

5

6

7

8

9

10

mM

P

0.0

0.1

0.2

0.3

0.4

0.5

pH (in) pH (out) P (in) P (out)

a

b

Figure 3.2. Average influent and effluent P concentration and pH for aerobic columns PB1 and 2 (a) and aerobic well-buffered column (b).

87

P concentration (mM)

0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14

Dis

tanc

e al

ong

colu

mn

(cm

)

-2

0

2

4

6

8

10

pH

6.5 7.0 7.5 8.0 8.5 9.0 9.5 10.0-2

0

2

4

6

8

10

P concentration (mM)

0.00 0.02 0.04 0.06 0.08 0.10 0.12

Dis

tanc

e al

ong

colu

mn

(cm

)

-2

0

2

4

6

8

10

pH

6.5 7.0 7.5 8.0 8.5 9.0 9.5 10.0-2

0

2

4

6

8

10

23 h120.5 h262.5 h550 h884 h1210 h

a

c

b

d

Figure 3.3. Side-port solution P concentration and pH through time for aerobic columnsPB1 (a-b) and PB2 (c-d). The Y-axis is distance along the column from the beginning ofthe ferrihydrite- coated sand. Data points plotted at –1 and 9 cm represent solution collected from uncoated silica mineral bed. Ferrihydrite-coated sand was present along the column from 0 to 7.5 cm.

88

Concentration (mM/g sand)

0.00 0.02 0.04 0.06 0.08

1.5

4

6.5

0.00 0.02 0.04 0.06 0.08

Dis

tanc

e al

ong

colu

mn

(cm

)

1.5

4

6.5

0.00 0.02 0.04 0.06 0.08

1.5

4

6.5

Fe/g sandPO4/g sand

Figure 3.4. Elemental concentration (mM) of P and Fe on the solid-phase after 50 d reaction for PB 1 and 2 (a-b) and after 21 d of reaction for the well-buffered column (c).

a

c

b

89

P concentration (mM)

0.00 0.02 0.04 0.06 0.08 0.10 0.12

Dist

ance

alo

ng c

olum

n (c

m)

-2

0

2

4

6

8

1017.5 h 67 h164 h 284 h 334 h 498 h

Figure 3.5. Aqueous P concentrations along the flow-path of the well-buffered, aerobic column. Data points plotted at –1 and 9 cm represent solution collected from uncoated silica mineral bed. Ferrihydrite-coated sand was present along the column from 0 to 7.5 cm.

90

II. Anaerobic columns

An abiotic experiment was conducted to generate mixed-valent iron oxides

common to reduced systems without complications induced by microbial activity.

Previous experiments (Benner et al., 2002) illustrate the similarity in reaction products

between abiotic (Fe(II)) and biotic systems. In order to generate reduced and mixed-

valent iron oxides, an FeCl2 solution was reacted with the ferrihydrite-coated sand for

120 h. At that time the column had changed color due to magnetite precipitation. The

lower section of the column, from the beginning of the ferrihydrite coated sand to 3.5 cm

downfield, was dark black, while the remainder of the column (3.5 cm to 7.5 cm

downfield) was a medium gray (Figure 3.6a). Upon introduction of the phosphate

solution a downfield migration of magnetite was observed, and after 6 d a small (1-2 mm)

band at the bottom of the column had returned to a red-brown color. The band continued

to expand, reaching a length of ~6 mm at experiment termination (Figure 3.6b);

additionally, the darkened band migrated downfield and the top of the column became

darker in color throughout the experiment. Upon termination, the column was sectioned

according to distinct color differences. The small band of red-brown sand was designated

as sample A. Sample B consisted of the darkest band of material from the center portion

of the column. The remaining discolored material was designated sample C. The dashed

white line in Figure 3.6b illustrates the division between samples B and C.

Phosphate breakthrough in the anoxic, anaerobic column was observed after 1 d.

The breakthrough curve was gradual and appeared to be related to a corresponding pH

increase (Figure 3.7a). Phosphate appears to be retained in the lowest portion of the

column, where magnetite had been removed (Figure 3.7b). Additionally, dissolved iron

was detected in the port solutions after the introduction of phosphate and increased with

time (Figure 3.8). Once the P front surpassed 3.5 cm, nearly complete breakthrough

proceeded rapidly. The pH increased up to near 9.5 as soon as the phosphate solution

was introduced into the column (Figure 3.7c). When the ferrihydrite-coated sand was

saturated with P, the pH dropped back to near the influent value. Indeed, phosphate was

sequestered within the small band of ferrihydrite-coated sand at the bottom of the column

(0 - 0.6 cm) (Figure 3.9), consistent with the observed breakthrough characteristics. The

91

remaining sections of the column appear to have an appreciable amount of sorbed Fe(II).

The ratios of Fe(III) to Fe(II) in Samples B and C were 1.3 and 3.1, respectively.

EXAFS (x-ray absorption fine structure) spectroscopic analysis was used to

determine the approximate mineralogical compositions in the top (Sample C), middle

(Sample B), and bottom (Sample A) of the column (Figure 3.10) following procedures

described in Benner et al. (2002). Table 3.2 lists the approximate mole percentages of the

iron minerals in each section of the column. The error is +/- 7% and values below 10%

are estimated. The predominate iron phase in each section of the column is different.

Ferrihydrite dominates the bottom of the column while magnetite and lepidocrocite are

the primary phases in the middle and top of the column, respectively.

Table 3.2. Distribution of iron (mole percent) in abiotic, anaerobic column upon experiment termination. A represents the influent end of the column (0 – 0.6 cm), B represents 0.6 o 5 cm, and C represents 5 to 7.5 cm.

mole % (wt/wt)

Section Ferrihydrite Magnetite Goethite Lepidocrocite

C 15 18 13 54

B 13 67 11 8

A 90 10 0 0

Phosphate retention within the packed mineral beds was determined during the total

extraction. Sand at the bottom (0 – 0.6 cm) of the column contained 5.69 µmols PO4/g

sand, while that in the middle (0.6 – 5 cm) contained 0.63 µmols PO4/g sand and the top

section (5 – 7.5 cm) contained 0.74 µmols PO4/g.

92

9

6.5

4.0

1.5

-1

Figure 3.6. Schematic illustration of ripened anoxic, abiotic column before (a) andafter (b) introduction of phosphate solution. The dashed white line indicates thetransition from black to gray. In (b) A, B, and C indicate where samples were collected at experiment termination.

a. b.

9

6.5

4.0

1.5

-1

Dis

tanc

e (c

m) f

rom

beg

inni

ng o

f fer

rihyd

rite-

coat

ed sa

nd

C

A

B

93

Time (h)

0 50 100 150 200 250 300

pH

2

4

6

8

P co

ncen

tratio

n (m

M)

0.0

0.1

0.2

0.3

0.4

0.5pH (in) pH (out) P (in) P (out )

P concentration (mM)

0.02 0.04 0.06 0.08 0.10 0.12

Dis

tanc

e al

ong

colu

mn

(cm

)

-2

0

2

4

6

8

10

pH

5.5 6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5-2

0

2

4

6

8

10

17.5 h 42.5 h 72 h 88 h 136 h 250.5 h

Figure 3.7. Influent and effluent P concentration and pH data (a) for abiotic, anoxicexperiment and side-port solution P concentration and pH through time. The Y-axis is distance along the column from the beginning of the ferrihydrite- coated sand. Data points plotted at –1 and 9 cm represent solution collected from uncoated silica mineralbed. Ferrihydrite-coated sand was present along the column from 0 to 7.5 cm.

a

b c

94

Fe concentration (mM)

0.00 0.01 0.02 0.03 0.04

Dist

ance

alo

ng c

olum

n (c

m)

-2

0

2

4

6

8

1042.5 h72 h88 h136 h250.5 h

Figure 3.8. Side-port solution Fe concentration for the anaerobic, abiotic experiment. The Y-axis is distance along the column from the beginning of the ferrihydrite- coated sand. Data points plotted at –1 and 9 cm represent solution collected from uncoated silica mineral bed. Ferrihydrite-coated sand was present along the column from 0 to 7.5 cm.

95

Concentration (mM/g sand)

0.00 0.02 0.04 0.06 0.08

Dis

tanc

e al

ong

colu

mn

(cm

)

0

2

4

6

8

PO4/g sand Fe2+/g sand total Fe/g sand

Figure 3.9. Elemental concentration of the solid-phase after 11 d of reaction for the anoxic, abiotic column. FeT is total iron concentration and A, B, and C represent locations along the flow path. A = 0 - 0.6 cm, B = 0.6 - 5 cm, and C = 5 - 7.5 cm.

A

C

B

96

Concentration (mole % Fe)

0 20 40 60 80 100-1

0

1

2

3

4

5

6

7

FerrihydriteMagnetiteGoethiteLepidocrocite

Figure 3.10. Mineralogical composition (mole % Fe, wt/wt) as determined by EXAFS analysis upon termination of anoxic, abiotic column.

Dis

tanc

e al

ong

colu

mn

(cm

)

97

The biotic columns behaved differently from the abiotic experiments. A detailed

discussion of these experiments can be found in Benner et al. (2002). In the groundwater

(GW) column (Figure 3.11a), phosphate was originally present in the effluent due to the

nutrient enriched media used to grow the bacteria. However, after that initial phosphate

spike, no further phosphate was observed in the effluent (i.e., complete retention of

phosphate was observed). The decrease in phosphate concentration in the GW feed

corresponds to a drop in calcium concentration, indicating possible precipitation within

the feed vessel. Phosphate breakthrough in the bicarbonate (BC) column was observed

after 430 h (Figure 3.11b). The initial phosphate concentration in the effluent peaked at

0.1 mM at 550 h and then dropped to ~0.07 mM for the remainder of the experiment.

The fluctuations in the BC effluent may, in part, be a reflection of changes in feed

chemistry. Phosphate levels in the feed were near 0.4 mM during the first 315 h of the

experiment and were then lowered to near 0.1 mM.

The extraction data on the sand also yielded differing trends. Within the GW

column (Figure 3.12a), phosphate was sequestered within the bottom 3–4 cm of the

ferrihydrite-coated sand and was removed only during the total extraction (6 M HCl).

Ferrous iron was also removed throughout the column, increasing downfield. Phosphate

was distributed throughout the BC column (Figure 3.12b), with the greatest amounts

retained from 8 to 12 cm downfield. Phosphate was removed both during the total and

partial extractions.

Upon termination, the porewater of the two columns also varied (Figures 3.13).

Phosphate was present in the BC porewater throughout the column and ranged in

concentration from 0.05 mM to 0.1 mM; however, phosphate was only detected in one

port about 6 cm downfield in the GW column. This observation may be due in part to the

different concentrations of phosphate in the respective feed solutions. The BC feed

contained much more phosphorus than did the GW feed (Table 3.1).

98

The mineralogical composition of the packed mineral beds varied markedly with

distance along the column (Table 3.3, Figure 3.14). Magnetite increases in abundance

with distance downfield and, not unexpectedly, the amount of ferrihydrite decreases

downfield. Goethite was also observed in the bottom 1 cm of the BC column and the

bottom 9 cm of the GW column, with decreasing amounts with distance downfield.

Table 3.3. Mineralogical composition (in terms of mole percent iron) with distance downfield in GW and BC columns. Flow direction is from 1 to 13.

GW BC

cm downfield Ferrihydrite Goethite Magnetite Ferrihydrite Goethite Magnetite

13 na na na 31 0 69 11 33 1 67 32 0 68 9 36 1 64 20 0 80 7 28 16 56 25 0 75 5 58 12 30 34 0 66 3 52 26 22 69 0 31 1 67 23 10 95 5 0

99

Time (h)

0 200 400 600 800 1000 1200

pH

3

4

5

6

7

8

P co

ncen

tratio

n (m

M)

0.00

0.01

0.02

0.03

0.04

0.05

Time (h)

0 200 400 600 800 1000 1200

pH

3

4

5

6

7

8

9

P co

ncen

tratio

n (m

M)

0.0

0.1

0.2

0.3

0.4

0.5

pH (in) pH (out)

mM P (in)

mM P (out)

pH (in)

pH (out) mM P (in)

mM P (out)

Figure 3.11. Influent and effluent P concentration and pH for the GW (a) and BC (b) columns.

a

b

100

Concentration (mM/g solid)

0.00 0.05 0.10 0.15 0.20 0.25

Dis

tanc

e al

ong

colu

mn

(cm

)

0

2

4

6

8

10

12

14

16

Concentration (mM/g solid)

0.00 0.05 0.10 0.15 0.20 0.25

Dis

tanc

e al

ong

colu

mn

(cm

)

0

2

4

6

8

10

12

14

16

mM PO4 (p) mM Fe (p)mM PO4 (t)mM Fe (t)

mM PO4 (p)mM Fe (p)mM PO4 (t) mM Fe (t)

Figure 3.12. Total (6 M HCl) and partial (0.5 M HCl) extraction data from GW (a) and BC (b) anaerobic, biotic columns. Concentrations are given in mM/g solid.

a

b

101

Concentration (mM)

0.00 0.02 0.04 0.06 0.08

Dis

tanc

e al

ong

colu

mn

(cm

)

0

2

4

6

8

10

12

14

16

Concentration (mM)

0.0 0.1 0.2 0.3 0.4

Dis

tanc

e al

ong

colu

mn

(cm

)

0

2

4

6

8

10

12

14

16

PO4 Fe

PO4 Fe

Figure 3.13. Fe and PO4 concentrations (mM) in pore water in GW (a) and BC (b) biotic columns at experiment termination.

a

b

102

Mole % Fe

0 20 40 60 80 100

1

3

5

7

9

11

13

15di

stan

ce a

long

col

umn

(cm

)

Mole % Fe

0 20 40 60 80 100

1

3

5

7

9

11

13

15

flow

dire

ctio

n

ferrihydritemagnetitegoethite

ferrihydritemagnetitegoethite

Figure 3.14. Mineralogical composition (in terms of mole % iron) of packed mineral beds within GW (a) and BC (b) columns upon experiment termination.

a b

103

DISCUSSION

The distinct difference in the breakthrough curves of the poorly- and well-

buffered aerobic columns corresponds to changes, or lack thereof, in pH. As phosphate

sorbs to the ferrihydrite, hydroxyl ions are released to solution (Reddy et al., 1999;

Torrent, 1997; Colombo et al., 1994; Goldberg and Sposito, 1985; Parfitt and Russell,

1977; Parfitt et al., 1975); thus, as phosphate migrates through the column, hydroxyl ions

are displaced and a high pH front ensues. In the poorly buffered columns, the hydroxyl

release subsequently leads to an increase in pH, which may diminish both the extent and

rate of sorption (Geelhoed, 1997; Strauss et al., 1997; Willett and Cunningham, 1983;

Ryden et al., 1977a; McLaughlin et al., 1977; and Hingston et al., 1972). In the poorly

buffered columns (PB 1–2), P breakthrough occurs after 4 d of flow and is directly

related to a pH increase to near 9.5. In contrast, breakthrough is not observed until 11 d

in the well-buffered system. The influence of pH is also apparent within the pore-water

along the column. A well-defined P front, delineated by the steep gradient, is observed in

the well-buffered column (Figure 3.5). The P front observed in the PB columns is more

diffuse and becomes increasingly so as the pH increases with time. It therefore appears

that the transient pH spike delays P sequestration.

Buffering the pH did not have any appreciable influence on the amount of

phosphate retained by the ferrihydrite-coated sand; however, reaction times differed

within the columns. The well-buffered column was reacted for ~500 h, while the poorly-

buffered columns were reacted for over 1200 h. The effluent P concentration from the

poorly-buffered columns exhibits several plateaus prior to complete breakthrough,

indicating that sorption is still occurring, albeit at a slower rate than what was observed

within the well-buffered column. The effects of the increased pH within the poorly-

buffered columns appear to be counteracted by the rapid flow rate through the column.

As the pH increases, P sorption will be inhibited, but continuous influx of feed solution

(pH = 7) will moderate this effect.

Phosphate breakthrough was achieved when near-equilibrium conditions were

reached between the solution and ferrihydrite-coated sand. However, in all the aerobic

columns the effluent concentrations at the end of the experiment were between 90% and

95% of the influent. The continued removal of P from solution is likely the result of slow

104

diffusion into micro-pores of the iron oxides (Ryden et al. 1977a). The final solid–phase

phosphate concentrations in both the poorly- and well-buffered columns were less than

the maximum adsorption capacity determined during batch study, ~12.6 µmols PO4/g

sand. The PO4 concentrations on the solid within the three columns ranged from 3.9 to

7.1 µmols PO4/g and averaged 6.7 µmols PO4/g for the PB columns and 4.6 µmols PO4/g

for the well-buffered column. Extraction data illustrate that phosphate is partitioned

similarly in all three aerobic columns and the average iron to phosphate ratios in the three

columns were between 5.6 and 6.3 (Figure 3.4). The lower overall iron and phosphate

concentrations in the well-buffered column are likely due to the use of sand that was

coated in a different batch than that used in the poorly buffered columns

The anoxic, abiotic column behaved quite differently from the aerobic columns

due predominantly to different mineralogical compositions. When the ferrous chloride

solution was discontinued, a distinct dark band was observable in the section of iron-

coated sand nearest the influent end of the column. Upon introduction of the phosphate,

further mineralogical change occurred that resulted in a ~6 mm band composed

predominantly (~90%) of ferrihydrite (defined as Section A). The concentration of Fe(II)

in the pore-water solution increases after introduction of the phosphate solution (Figure

3.8), suggesting magnetite dissolution or the displacement of sorbed Fe(II) by phosphate.

The phosphate breakthrough curve is similar to that observed in the poorly

buffered aerobic columns, indicating that pH is one controlling factor of phosphate

retention within the column. The initial breakthrough is concurrent with an increase in

pH (Figure 3.7), though the pH increase in the anoxic column is not as dramatic as that in

the poorly buffered aerobic columns. Additionally, a well-defined phosphate front does

not propagate through the anoxic column (Figure 3.7). Pore water concentrations of P

drop very rapidly during the first day of reaction; however, after that initial period,

phosphate retention decreases dramatically. The immediate drop in retention and

incomplete removal of phosphate corresponds to pH increases to greater than 9.0 within

the column. On the basis of temporal trends in pore-water within the column, it appears

phosphate is being retained predominantly in the lowest section (A) of the column. Once

the material between 0 and 2 cm reaches equilibrium with the influent solution, nearly

complete phosphate breakthrough is observed.

105

The mineralogical composition of the anaerobic, abiotic column at the end of the

experiment varied markedly downfield from the lowest portion of the column (Table 3.2).

The change in iron mineralogy occurred during the initial 5 d period when FeCl2 was

reacted with the ferrihydrite-coated sand. Upon termination, the central, darkest section

(B) of the column was mostly magnetite (~67%) with lesser amounts of ferrihydrite,

goethite, and lepidocrocite. The section of the column furthest downfield (C) contained

~54% lepidocrocite and between ~13% and ~18% (each) ferrihydrite, goethite, and

magnetite. In the presence of low concentrations of Fe(II), ferrihydrite converts to

lepidocrocite or goethite; chloride will favor lepidocrocite formation, while its absence or

the presence of CO32- will favor ferrihydrite conversion to goethite. Magnetite, in

contrast, is formed at higher (>0.4 mM) Fe(II) concentrations at pH ~7 (Benner et al.,

2002; Zachara et al., 2002).

The total extraction data also reflect the variation in mineralogy throughout the

column. Most of the phosphate retained within the column is sequestered in Section A

(~90% ferrihydrite), having a final phosphate concentration of 5.89 µmols PO4/g solid,

slightly lower than the equilibrium value of 9.8 µmols PO4/g determined by batch

experiment for ferrihydrite-coated sand reacted with 200 µM PO4 (I = 0.1 M), which

reached an equilibrium solution concentration of ~100 µM PO4 after 2 d. The remaining

sections of the column retained comparatively little phosphate, 0.62 and 0.69 µmols

PO4/g solid for sections B and C, respectively; these values are similar to the equilibrium

adsorption of phosphate by geologic magnetite reacted with 100 µM PO4 (~0.7 µmols

PO4/g, equilibrium solution concentration ~95 µM PO4). Even though most of the iron is

sequestered within lepidocrocite in the effluent end of the column, the adsorption

characteristics are more similar to that of geologic magnetite. This may be the result of a

layer of magnetite coating the solid within the packed mineral bed or it may be due, in

part, to the sorption of ferrous iron, which would passivate reactive surfaces of

lepidocrocite with respect to phosphate (Roden et al., 2000). The increased iron in

column pore-water indicates that iron is being lost from the system, either from

dissolution of magnetite and/or from desorption of ferrous iron. In any case, sorption of

ferrous iron is likely to hinder phosphate retention within the column. The extraction

data from Section B support this idea. Within this section of the column the ratio of

106

ferric to ferrous iron is ~1.3, indicating that either the magnetite present is non-

stoichiometric or, more plausible, that ferrous iron is sorbed to the solids.

Batch experiments (Chapter 2) reveal that adsorption of P on magnetite is

particularly sensitive to pH, resulting in negligible sorption at pH 9. However, the lack of

P retention within sections B and C is not due entirely to pH restraints. While increased

pH may have prohibited P sorption in the early stages of the experiment, only minimal

quantities of P were removed downfield from Section A even after the pH had returned to

near neutral. The rapid breakthrough downfield from Section A indicate that only minor

amounts of P, if any, are retained within Sections B and C.

The surface area (SA) of the material in the column decreased over the reaction

period. The surface area of the solid matrix was initially ~4.6 m2/g; after reaction,

material in Section A had a SA of 2.65 m2/g, material from B had a SA of 0.741 m2/g,

and material from Section C had a SA of 0.488 m2/g. The low surface area of the

secondarily mineralized iron phases is consistent with the large grain-size observed for

abiotic transitions (Hansel and Benner, unpublished data).

Several important observations must be addressed regarding the biotic

experiments; an even more elaborate discussion can be found in Benner et al. (2002).

The effluent phosphate concentration trends are nearly opposite in the two columns. In

the GW column, P in the effluent immediately spikes to ~0.02 mM and rapidly drops to

below detection levels. The initial high P concentration in the effluent is an effect of the

nutrient enriched media used to grow the bacterial cultures. Conversely, P in the BC

effluent is not initially present, but after ~18 d of reaction the effluent concentrations

increase to ~0.1 mM and then appear to reach a stable value of ~0.07 mM. These

differing trends are, at least in part, due to the different feed solution chemistries. The

average P concentration in the BC feed is nearly 40 times that in the GW influent, which

would lead to saturation of available sorption sites and P breakthrough in the BC column

before the GW column. Phosphorus is partitioned along the length of the BC column,

whereas in the GW column P is retained only in the bottom 2 cm. In the BC column, the

maximum P retention observed is ~10.5 µmols PO4/g solid. However, magnetite is the

major iron-bearing mineral in the downfield section of the column (Figure 3.14). The

characteristics of P adsorption within the BC column are predominantly the result of

107

biomineralization processes. Biogenic minerals are likely to have very high surface areas

and may also incorporate phosphate during precipitation.

In the bottom section of the GW column the P concentration is ~4.2 µmols PO4/g

solid. The unexpectedly high adsorption capacity in this section of the GW column may

be explained by a variation in iron content of the solid, precipitation of a calcium

phosphate mineral phase, or a combination of the two. Comparison of the influent and

effluent solution chemistries from this column shows that appreciable amounts of

calcium and phosphate are removed from the influent during the first two days of

reaction.

In general, P sorption capacities within the aerobic columns are in agreement with

the experimental sorption capacities determined during static experiments. The solids in

the anaerobic columns behaved quite differently than predicted based on adsorption

isotherms. Possible reasons for this include passivation of surfaces by ferrous iron and

the conversion of ferrihydrite to goethite and lepidocrocite (the latter transition would

thus decrease surface area and micro-porosity). Magnetite precipitated within the

columns may also have different properties than that used to establish the isotherm.

Within the biotic column, incorporation into solid phases, precipitation as Ca-phosphates,

and competition with other ions, are possible factors affecting P retention. Finally, P

adsorption by ferrous minerals is, in general, much more sensitive to fluctuations in pH.

Even though ferrous minerals have relatively high capacities for phosphate retention, the

affinity for phosphate is much lower than for the ferric minerals. Given the lower

affinities, one would postulate that sorption by ferrous minerals would be fairly easily

reversed if the surrounding chemical conditions changed slightly (i.e. pH, ionic strength,

and concentration of P in solution).

CONCLUSIONS

Buffering the influent pH of the aerobic columns did not have any appreciable

influence on the amount of phosphate retained by the ferrihydrite-coated sand over a long

reaction period; however, pH was a determining factor in the characteristics of the

breakthrough curve. The anoxic, abiotic column behaved quite differently from the

aerobic columns predominantly due to different mineralogical composition. Most of the

108

phosphate sequestered within that column was sorbed by a narrow (0.6 cm) band of

ferrihydrite-coated sand near the influent end. The remainder of the anoxic column

exhibited sorption capacities similar to that of geologic magnetite, which may be

explained by passivation of reactive surfaces due to Fe(II) sorption or the precipitation of

a layer of magnetite upon the ferrihydrite.

The pH of a system can influence both the rate and extent of PO4 sorption — a

factor dependent on the specific mineral. The presence of elements such as Ca and Fe(II)

may promote the precipitation of phosphate minerals, which increases the apparent

adsorption capacity of a sorbent; however, they may equally serve as inhibitors of

phosphate adsorption.

Reductive dissolution of ferric oxides leading to the precipitation secondary

ferrous minerals create complicated mixtures of sorbents, each of which will exhibit

different sorption capacities and responses to changes in pH. Sorption of ferrous iron

may passivate reactive sites on the primary and secondary phases that could otherwise

sorb phosphate. This work has demonstrated that phosphate retention in dynamic

systems is dependant upon many interrelated factors, each of which must be considered

in order to fully understand and predict P retention, even in these relatively simple

geochemical systems.

The stability and capacity of phosphate retention in systems prone to transitions

from aerobic to anaerobic conditions will be determined by the characteristics of the

secondary ferrous minerals and the presence of other sorbates. In general, P sorption by

anoxic systems will be lower and less stable than that in aerobic systems. Secondary

ferrous minerals have lower surface areas and are easily passivated by the sorption of

ferrous iron or other ions. Additionally, all ferrous phases studied exhibited relatively

low affinities for phosphate, indicating that even small changes in geochemical

conditions would destabilize the system and lead to phosphate desorption.

109

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