MEASURING SURFACE CHEMICAL PROPERTIES USING FLOW ADSORPTION CALORIMETRY: THE CASE OF AMORPHOUS ALUMINUM HYDROXIDES AND ARSENIC (V)
By
NADINE JACK KABENGI
A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
2004
Copyright 2004
by
Nadine Jack Kabengi
To Papy, Mamy, Karine and Sabine
ACKNOWLEDGMENTS
First and foremost, I thank my parents and sisters for their continuous support, and love. They believed and encouraged me during my whole life, sacrificing so much for me to get the best of opportunities. For that and much more, I will be forever indebted. I would have never been able to make it without the unconditional love and friendship of my friends Daad Abi Ghanem, Tania Ghandour, and Nathalie Khattab. I am deeply grateful for them being part of my daily life for the past eleven years in spite of being thousands of miles apart most of the time. Countless hours of conversations, phone calls and emails with my mentor and friend, Dr. Riad Baalbaki, were decisive in shaping the way I look at almost every aspect of life. I am a better person because of his influence.
I would like to thank the Wedgworth family whose generous donation to the
Everglades Research and Education Center supported my assistantship for the entire duration of my work. I will be forever indebted to them for giving me such a precious opportunity. I realize I have been twice as lucky to be advised by two remarkable individuals, Dr. Samira H. Daroub and Dr. R. Dean Rhue, who offered unlimited guidance and expertise during my studies and helped me grow as a scientist and a human being. They have both been more than advisors to me. In addition to sharing his knowledge and genuine interest in science with me on a daily basis, Dr. R Dean Rhue is a true mentor who taught me how to fish. I have enormously enjoyed every fishing trip we had, talking about science, life and anything in between. I thank his wife, Mrs. Janice
Rhue, for treating me like a daughter and cooking me lots of good meals.
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I owe special thanks to Dr. Willie G. Harris not only for his inexhaustible help and for his interest in my project and me, but especially for sharing with me his bohemian love of classical music, art and good conversation. I am deeply grateful to Dr. Nick
Comerford for all his expert help and suggestions equally as a scientist and a graduate coordinator. I am also thankful that Dr. Comerford allowed me to appreciate his witty sense of humor. I owe special thanks to Dr. Michael J. Scott for being on my committee and always being willing to help, but especially for treating me more like a friend than an advisee. I thank Dr. Randall Brown, who agreed to sit on my committee and took genuine interest in my research. I will be forever grateful to Dr. Samuel O. Colgate for introducing me to the magical world of physical chemistry and for making me believe I could actually obtain a minor in it. Dr. Colgate also shared with me his talent in glass blowing and I am proud to have taken this special class with him. Dr. Colgate and his wife treated me as a family friend, inviting me to his home on several occasions and for that I am very touched.
I could have not completed this work without the assistance of several talented individuals in the Soil and Water Science Department at the University of Florida, namely Mr. Keith Hollien, Mrs. Elizabeth Kennelly, Mr. Thomas Luongo and Ms. An
Nguyen. A special place in my heart belongs to Dr. Jaimie Sanchez. He met me at the airport upon my first arrival to the United States and helped me settle down, taking time to guide me through almost everything. I will be forever indebted to Dr. Jaimie Sanchez and his lovely family. Special thanks are due to my friends, Dr. Ziyad Mahfoud, Mr.
Joseph Nguyen, and Dr. Bassem Sabra, without whom I would have never lasted through my first month in Gainesville. They helped me in every aspect of life initially and
v
through the next four years and for that I am thankful. Dr. Rao Mylavarapu, who acted as
a co-chair upon my arrival, is greatly acknowledged for his support and his advice. I also
thank Dr. Hector Castro and Dr. Kanika Sharma for their friendship and support, and
especially for putting up with my multi facetted moods.
Lastly but definitely not least, I would like to thank my friends and coworkers Dr.
Chip Appel and Mr. Bill Reve. I can honestly say that I could not have done it without them and their there-are-no-words-to-describe-it help. I appreciate them, our friendship, and their desire to see me do well. They have made coming to work much more enjoyable, and I have loved having both of them as part of my daily life. I thank them for befriending me and laughing at my jokes.
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TABLE OF CONTENTS
page
ACKNOWLEDGMENTS ...... iv
LIST OF TABLES...... x
LIST OF FIGURES ...... xii
ABSTRACT...... xv
CHAPTER
1 INTRODUCTION ...... 1
2 CALORIMETRY FUNDAMENTALS...... 7
Instrumentation...... 7 Sensitivity and Precision...... 8 Interpreting a Heat Signal...... 10 Operation ...... 13
3 SAMPLES SYNTHESIS AND CHARACTERIZATION...... 14
Materials ...... 14 Synthesis And Preparation of AHO...... 14 Physical Characterization ...... 15 X-Ray Diffraction Analysis...... 15 Thermogravimetric Analysis...... 15 Surface Area Measurement ...... 15 Scanning Electron Microscopy...... 16 Chemical Characterization...... 24 Chemical Composition ...... 24 Anion Exchange Capacity ...... 24 Cation Exchange Capacity ...... 25 Effect of pH on Aluminum Content ...... 26
4 CHEMICAL CHARACTERIZATION USING FLOW ADSORPTION CALORIMETRY ...... 28
Effect of Carbonate...... 28
vii
Ion Exchange ...... 29 Obtaining Heats of Ion Exchange...... 29 Results ...... 30 Anion exchange...... 30 Cation exchange ...... 32 Zero Point of Charge ...... 34 Rationale...... 34 Procedure...... 34 Results ...... 35 Charging of AHO surfaces ...... 38 Additional Observations...... 39 Flip-Flop effect...... 39 Change of anion exchange energetics with pH ...... 40 Reversibility of Surface Charge with pH...... 42 Rationale...... 42 Procedure...... 42 Results ...... 42
5 ENERGETICS OF ARSENATE SORPTION ...... 45
Obtaining Heats of Reactions ...... 45 Results...... 46 Calorimetric Effects...... 46 Arsenate Sorbed...... 51 Heats of Reactions...... 54 Changes In pH With Arsenate Sorption ...... 56 Rationale...... 56 Procedure...... 57 Results ...... 57
6 EFFECT OF ARSENATE SORPTION ON AMORPHOUS ALUMINUM HYDROXIDES SURFACES ...... 59
Effect of Arsenate Sorption on Ion Exchange ...... 59 Rationale...... 59 Procedure...... 60 Results ...... 60 Anion exchange...... 60 Cation exchange ...... 66 Effect of Arsenate Sorption on PZC...... 68 Rationale...... 68 Procedure...... 69 Flow studies...... 69 Batch studies ...... 69 Results ...... 70 PZC shifts: Flow versus batch...... 70 PZC shifts: CEC effects ...... 73
viii
7 THE ARSENATE /AMORPHOUS ALUMINUM HYDROXIDES SORPTION SYSTEM: FINDINGS AND SUGGESTIONS...... 80
Summary of Results...... 80 AHO Physical and Chemical Properties...... 80 Ion Exchange Properties...... 81 Arsenate Sorption Properties...... 81 Effects of Arsenate Sorption ...... 81 Findings and Suggestions ...... 82 On the Structure and Morphology of AHO...... 82 On AHO Surface Chemistry...... 83 On Arsenate Sorption ...... 84
8 SUMMARY AND CONCLUSIONS...... 88
APPENDIX RAW DATA ...... 91
LIST OF REFERENCES...... 104
BIOGRAPHICAL SKETCH ...... 108
ix
LIST OF TABLES
Table page
3-1 Specific surface areas of the amorphous aluminum hydroxide batches...... 16
3-2 Aluminum content of the amorphous aluminum hydroxide batches ...... 24
3-3 Anion exchange capacity of the amorphous aluminum oxides ...... 25
3-4 Average Al:Cl mole ratio calculated for the four AHO batches ...... 25
3-5 Effect of various pH treatments on the percent change in Al content in AHO batches...... 26
4-1 Peak areas associated with Cl/NO3 exchange peaks of batches 1, 2, 3, and 4 shown in figure 4-1...... 32
4-2 Heats of Cl /NO3 exchange measured calorimetrically on different batches of amorphous aluminum oxides...... 32
4-3 Peak areas, CEC and heats of K/Na exchange obtained a pH 10.5...... 33
4-3 Data collected for the determination of the point of zero charge of a B2 column. ..36
5-1 Arsenate loadings and corresponding Al:As mole ratios obtained on all four batches of amorphous aluminum hydroxides...... 52
5-2 ∆H values, amounts of sorbed arsenate and Al:As molar ratios for AHO samples from batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4)...... 54
5-3 Heat production after sequential additions of orthophosphate solution to a Fe(OOH) suspension...... 56
5-4 Effect of arsenate sorption on pH of solution...... 58
6-1 Peak areas and reductions associated with Cl/NO3 exchange peaks before and after arsenate exposure of samples shown in figure 6.1...... 60
6-2 Comparisons between samples that showed an increase in Cation Exchange Capacity (CEC) after As exposure and samples that did not...... 68
x
6-4 Comparisons in zero point of charge shifts and other data of batch 3 samples arsenated in flow and in batch...... 72
6-5 Comparisons of heats of ion exchange and other data from clean and arsenated batch 2 samples at different pHs...... 74
6-6 Comparisons of heats of ion exchange and other data from clean and arsenated batch 3 samples at different pHs...... 75
6-7 Reductions in Cl/NO3 peak areas with increase in solution pH from 5.75 to 7.25 and 8.0 for clean batches 2 and 3 samples...... 77
6-8 Reductions in Cl/NO3 peak areas with increase in solution pH from 5.75 to 8.0 for clean and arsenated batche 2 and 3 samples...... 77
A-1 Ensemble of data collected for columns in arsenate sorption experiments...... 91
A-2 Ensemble of data collected for columns in arsenate sorption experiments...... 92
A-3 Ensemble of data collected for columns in arsenate sorption experiments...... 93
A-4 Ensemble of data collected for columns in arsenate sorption experiments...... 94
A-6 Ensemble of data collected for columns in arsenate sorption experiments...... 95
A-7 Ensemble of data collected for columns of clean AHO in ZPC experiments...... 96
A-8 Ensemble of data collected for columns of clean AHO in ZPC experiments...... 97
A-9 Ensemble of data collected for columns in ZPC experiments after arsenate adsorption...... 98
A-10 Ensemble of data collected for columns in ZPC experiments after arsenate adsorption...... 99
A-11 Ensemble of data collected for columns in ZPC experiments after arsenate adsorption...... 100
A-12 Ensemble of data collected for columns in ZPC experiments after arsenate adsorption...... 101
A-13 Ensemble of data collected for columns in ZPC experiments after arsenate adsorption...... 102
A.14 Ensemble of data collected from columns used in reversibility experiments...... 103
A.15 Ensemble of data collected for columns used in energetics experiments...... 103
xi
LIST OF FIGURES
Figure page
2-1 Schematic of column, thermistors, and calibrating resistors used in flow calorimetry...... 7
2-2 Example of a curve depicting the linear relationship between peak area size and flow rate...... 9
2-3 Peaks obtained with various size heat pulses (A) and the associated calibration curve (B)...... 10
2-4 Pulse mode-physical adsorption and system calibration...... 11
2-5 Saturation mode- physical adsorption and desorption by carrier fluid ...... 12
2-6 Saturation mode-chemisorption...... 12
3-1 XRD pattern for batch 1 AHO A) untreated samples, and B) samples washed with DDI...... 17
3-2 XRD pattern for batch 2 AHO A) untreated samples, and B) samples washed with DDI...... 18
3-3 XRD pattern for batch 3 AHO A) untreated samples, and B) samples washed with DDI...... 18
3-4 XRD pattern for batch 4 AHO A) untreated samples, and B) samples washed with DDI...... 19
3-5 TGA patterns for batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4)...... 19
3-6 Scanning electron micrographs showing AHO batch 1 morphology at 3 different scales...... 20
3-7 Scanning electron micrographs showing AHO batch 2 morphology at 3 different scales ...... 21
3-8 Scanning electron micrographs showing AHO batch 3 morphology at 3 different scales ...... 22
xii
3-9 Scanning electron micrographs showing AHO batch 4 morphology at 3 different scales ...... 23
4-1 Heats of Cl/NO3 exchange obtained at pH 5.7 on batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4)...... 30
4-2 Replicates of Cl/NO3 heats of exchange obtained at pH 5.7 on batch 1 over a period of 6 days...... 31
4-3 Heats of K/Na exchange obtained at pH 10.5 on batch 2 (B2), batch 3 (B3) and batch 4 (B4)...... 33
4-4 A) Cl/NO3 and B) K/Na heats of exchange obtained on AHO batch 1 at different pHs...... 36
4-5 Corrected heats of cation and anion exchange for batches 2 and 3 of AHO as a function of solution pH, measured by flow adsorption calorimetrically...... 37
4-6 Effect of ascending and descending pH on calorimetric peak areas of batches 2 and 3 of AHO measured by flow adsorption calorimetry...... 43
4-7 Effect of ascending and descending pH on the AEC of batches 2 and 3 of AHO samples...... 44
5-1 Heats of reaction of arsenate with batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4)...... 47
5-2 Contrast between peak shapes for Cl/NO3 exchange and arsenate reaction with an AHO batch 1sample...... 48
5-3 Two arsenate cycles on same sample of batch 1 of AHO...... 48
5-4 Four consecutives arsenate cycles and corresponding returns to NaCl solution collected consecutively (A, B, and C) and after 24 hours (D) on same AHO sample...... 49
5-5 Two consecutive arsenate cycles (A, B) and corresponding return to NaCl solution (A) collected on same AHO sample...... 50
5-6 SEM-EDX elemental maps of two distinct particles, A and B, of an arsenated batch 1 AHO sample...... 53
6-1 Cl/NO3 heats of exchange before and after exposure to arsenate for batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4) samples...... 61
6-2 Relationship between heats and magnitude of Cl/NO3 exchange of batch 1 samples throughout various arsenate treatments...... 62
xiii
6-3 Loss of available anion exchange capacity of B1 samples (in µmoles of charge) as a function of amount of arsenate sorbed in µmoles...... 63
6-4 Loss of Cl/NO3 heat in relation to an increasing arsenate surface loading ...... 65
6-5 K/Na heats of exchange at pH 5.75 after exposure to arsenate for batch 1 (B1), batch 2 (B2), and batch 3 (B3) samples ...... 67
6-6 Calorimetric determination of the ZPC of a batch 2 (B2) and batch 3 (B3) AHO sample treated with As(V) in a flow system...... 71
6-7 Calorimetric determination of the ZPC of batch 3 (B3) AHO samples treated with As(V) in a batch system...... 72
xiv
Abstract of Dissertation Presented to the Graduate School of the University of Florida in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy
MEASURING SURFACE CHEMICAL PROPERTIES USING FLOW ADSORPTION CALORIMETRY: THE CASE OF AMORPHOUS ALUMINUM HYDROXIDES AND ARSENIC (V)
By
Nadine Jack Kabengi
August, 2004
Chair: Samira Daroub Cochair: R. Dean Rhue Major Department: Soil and Water Science
Flow adsorption calorimetry provides a direct, quantitative measure of the heat
evolved in a reaction and is ideally suited for measuring reactions occurring at the
liquid/solid interface. This study centers on demonstrating the application of flow adsorption calorimetry as a powerful technique in probing chemical surfaces, thus obtaining information not readily accessible by other methods. Methodologies developed were applied to the study of arsenate sorption onto amorphous aluminum hydroxides
(AHO). The AHO surface was chemically characterized using flow calorimetry. Heats of
-1 exchange determined for Cl/NO3 were 3.6 to 5.8 kJ mol and those for K/Na 1.2 to 1.6 kJ
mol -1. A new technique to measure the zero point of charge (ZPC) calorimetrically was developed and used. Measured ZPC values were consistent with literature values and were around 8.5 and 9.5 for the AHO. The charging of the AHO surfaces supported a 2 pK mechanism as calorimetric data indicated that at pHs around the ZPC, the surface of
xv
AHO is neutrally charged: i.e, neither positive nor negative charges exist. The reaction of
arsenate with AHO was exothermic (40 to 60 kJ mole-1) and decreased with increasing
surface coverage. Most of the arsenate was sorbed on positively charged sites and
involved no increase in surface negative charge. Arsenate sorption exhibited an initial, rapid uptake phase, thought to be ligand exchange with aquo- and hydroxo- groups, followed by a slower secondary reaction resulting from arsenate gaining access to the less
accessible reactive sites that are dispersed throughout the structure. With increasing
amounts of arsenate sorbed, arsenate reacted with the bridging ol groups, creating
negative sites for cation exchange. This bond-breaking mechanism occurred at a slower
rate and was energy consuming. Sorbed arsenate caused the ZPC to shift by around 1 pH
unit. The shift in ZPC resulted from a change in the pK of reactive hydroxyl groups. The
role of polydentate arsenate complexes and deprotonation of sorbed arsenate was also
discussed. The results of this study show that flow adsorption calorimetry is a uniquely
informative yet rapid experimental tool that can be applied to numerous applications in
surface chemistry studies.
xvi CHAPTER 1 INTRODUCTION
Surface chemical reactions at interfaces between natural solids and aqueous
solutions play a fundamental role in geochemical and environmental contexts. The
importance of interfacial processes is emphasized in the following quotation from Werner
Stumm et al. (1987): “Almost all of the problems associated with understanding the
processes that control the composition of our environment concern interfaces, above all
the interfaces of water with naturally occurring solids.” While we have a fundamental understanding of these processes, which include mineral dissolution and precipitation as well as sorption and desorption of chemical species to surfaces, a molecular scale understanding of the complexity and interplay of the chemical and biological processes has yet to be achieved (Brown et al., 1999). To understand the relation between surface chemical reactivity and its composition, structure, and properties, new experimental techniques, which provide additional molecular-level information must be applied.
Calorimetry provides a direct, quantitative measure of the heat involved in a reaction. This measured heat is related to a change in enthalpy, a fundamental property of the system that can provide information about the chemical and physical processes taking place. Calorimetric measurements can be conducted in various ways, yet they are generally divided into two categories: (1) “batch” adsorption, and (2) “flow” adsorption calorimetry.
Flow adsorption calorimetry is better suited for measuring interactions occurring at the liquid/solid interface. Flow calorimetry has the following advantages over
1 2
conventional batch calorimetry (Steinberg, 1981): (i) flow calorimetry can resolve a
complex series of reactions that occur more or less simultaneously but at different rates;
(ii) multiple adsorption/desorption cycles can be applied to the same sample, allowing
reversible and irreversible processes to be distinguished; (iii) changes that occur in the
surface properties of the solid as a result of specific treatments or aging effects can be
quantified; and (iv) when both the amount of adsorption and its associated heat are
measured, information about surface heterogeneity can be obtained.
Flow adsorption calorimetry has been used to study the surface chemistry of
different types of solids in a number of different applications. Surface area measurements
were one of flow calorimetry’s first applications, the measured heats of adsorption being
directly proportional to the solid surface area and pre-calibrated with a solid of known
area for a particular solute/solvent system (Steinberg, 1981). Using a comparable
rationale, Groszek and Partyka (1993) were able to relate the integral heats of
displacement of 1-butanol from n-heptane and water solutions to hydrophobic and
hydrophilic surface areas for a variety of carbonaceous and mineral solids. Likewise,
Meziani et al. (1997) used the 1-butanol/water system to evaluate the polar surface sites
of the mesoporous silicoaluminates of the MCM-41 in relation to increasing chain length
and Al content. Mathonat et al. (1998) demonstrated the use of flow calorimetry for
determining enthalpies of absorption and the solubility of CO2 in aqueous monoethanolamine solutions. In a technique similar to that for surface area, it is also possible to measure the Lewis acid/base character of solid surfaces. Schneider et al.
(1997) investigated the acid-base character of organic polymer surfaces and were able to obtain good correlation between their calorimetric data and results
3
obtained with X-ray photoelectron spectroscopy. Application of flow calorimetry has
even extended to microbial biodegradation studies in soils. Critter and Airoldi (2001)
argued that the calorimetric method is very sensitive toward changes in the microbial
biomass, which could not be detected by more conventional methods. Rhue et al. (2002)
measured surface chemical properties of soil using flow calorimetry and were able to
obtain novel information regarding phosphate sorption on Ultisol samples.
Despite all of the aforementioned applications, the actual capabilities of flow
adsorption calorimetry remain essentially unexploited in numerous fields, namely in
adsorption science and its implications for soils and aquatic systems, where questions
about some of the basic chemical processes that occur at the surface of soils, clays and
sediments remain unanswered despite decades of investigations and studies. So far, our
knowledge on adsorption systems, such as oxide/electrolyte interfaces, has come almost
exclusively from adsorption isotherms (titrations or individual) and electrokinetics
effects. And while these have yielded useful information on rates and amounts of solute
sorbed as well as equilibrium constants, they give no insight into the mechanism by
which sorption occurs (Sposito, 1984). Moreover, according to Rudziński et al. (1999a),
it is a known fact in adsorption science that calorimetric effects of adsorption are much more sensitive to the nature of an investigated adsorption system than adsorption
isotherms because calorimetric effects are related to temperature and temperature appears
in the common exponential terms in statistical thermodynamics. In fact, these authors
have published calorimetric experiments that decreased by one half the number of
“curve-fitted” parameters, and continue to utilize flow adsorption calorimetry for
theoretical interpretations of simple ion adsorption at the oxide/electrolytes interfaces.
4
This work centers on demonstrating the application of flow adsorption calorimetry as a powerful technique in probing and understanding chemical surfaces, obtaining thus information not readily accessible by other methods. Methodologies developed will be further applied to the study of a classic oxyanion/oxide sorption system in soil science illustrating how, in conjunction with existing technologies, flow calorimetry can greatly improve our understanding of basic surficial processes in soil/clay systems. For that purpose, we have elected to study the case of amorphous aluminum hydroxides and arsenic (V).
In light of soil’s complexity, experimental and theoretical studies investigating the fundamental aspects of these processes are not only difficult to perform on natural samples. Also, results could be rather complicated to interpret on the grounds of the large number of inter-related biotic and abiotic processes. One approach to this problem in soil science has been to use the individual mineral components known to be responsible for the sorption capacity of a particular adsorbate. Aluminum oxides are abundant in natural waters and soils as high surface area minerals, mineral coatings, and colloids. Because of their significant adsorptive properties, they have been recognized, along with other oxides and organic matter of soils, as largely responsible for soil’s sorption capacity.
Furthermore, the predominant Al oxides phases in soils are often the amorphous species that possess greater reactivity than the corresponding crystalline mineral surfaces (Bohn et al., 1979; McBride, 1994). These amorphous minerals are often difficult to isolate in pure form, prompting the usage of synthetic amorphous aluminum oxides as reference materials instead.
Although the choice of adsorbate is not quite central to the proceedings of this
5 research, we had two motives for choosing to investigate arsenate. First, arsenic is the focus of public attention and is receiving special attention from the scientific community.
This scrutiny is mainly due to the almost epidemic-like health problems of hundreds of thousands of people in Bangladesh and West Bengal, India, caused by arsenate- contaminated groundwater. This unsavory reputation achieved by arsenic is also due to the toxic and carcinogenic properties of a number of its compounds. Arsenic is rather ubiquitous in the environment as a result of mineral weathering and dissolution, geothermal activity, and numerous anthropogenic activities. Sources include arsenical pesticides, fertilizers, mine drainage, smelter wastes, dipping vats, tanning industries, wood preservatives and feed additives.
Second, despite an abundance of literature on arsenate, and other oxyanions such as phosphate, chromate, and molybdate and their reactions with oxide surfaces, the exact reaction mechanism is still unclear and no one model has been found to fit all data. There seems to be a general consensus that these inorganic oxyanions adsorb to soil minerals as a two-step ligand exchange reaction. Yet, Sposito (1989) indicates that none of the pieces of experimental evidence is conclusive in itself; they only form a powerful argument for the applicability of oxyanion adsorption if considered all together. Also there have been sprinkled in the literature some examples that do not necessarily challenge the applicability of ligand exchange but that indicate that other reactions might also occur depending on the pH of the solution (Rajan et al., 1974) or on surface coverage (Jain et al., 1999). Calorimetric investigation of the arsenate/oxide system will uncover additional pieces of evidence that no doubt will greatly benefit the existing state of knowledge on inorganic oxyanions sorption in general.
6
The first two chapters (2 and 3) are mainly descriptive in nature: Chapter 2 presents a detailed description of calorimetry fundamentals and instrumentation constructed in our laboratory, while Chapter 3 deals with the aluminum hydroxide synthesis and characterization. Chapter 4 is geared toward demonstrating the applications of flow adsorption calorimetry to surface chemical characterization. In chapter 4, we report first on the direct measurements of the calorimetric effects of ion adsorption. Second, we present new techniques developed to measure the point of zero charge calorimetrically and examine the reversibility of surface charge with pH. Chapters 5 and 6 deal with the arsenate/oxide sorption system; whereby in the former we examine the physical and chemical reactions involved in arsenate sorption, and in the latter the effects of arsenate sorption on the aluminum hydroxides surfaces are identified calorimetrically. The last chapters (7 and 8) constitute a synthesis of the important conclusions learned from each of these experiments, putting forward a proposed arsenate reaction model, as well as suggestions for further research and future applications.
CHAPTER 2 CALORIMETRY FUNDAMENTALS
Instrumentation
Several inexpensive flow calorimeters for measuring heats of adsorption from solution onto solids were constructed in our lab. A calorimeter consisted of a small column assembly (Fig 2-1) sealed inside a 500-mL polyethylene bottle, and the bottle was placed in a 50-L insulated container, filled with water, at room temperature. The water bath provided good thermal stability against ambient temperature changes and generally resulted in baselines with negligible drift.
A pair of thermistors, one upstream and the other downstream from the column, formed one-half of an electronic bridge and sensed temperature changes in the solution as it passed through the column. A change in solution temperature produced a differential output voltage from the bridge. This differential voltage was fed into an instrumentation amplifier, and the amplified signal was fed into a computer for processing. The system possessed high sensitivity, low thermal drift, and a good signal-to-noise ratio.
Figure 2-1. Schematic of column, thermistors, and calibrating resistors used in flow calorimetry.
7 8
Approximately 15 mg of material were placed inside the column and solutions
containing reactive species were forced through the column using a total pressure drop of
about 100 cm of water. Flow rates were controlled with a precision needle valve at the
outlet side of the calorimeter and were generally constant to within a few µl min-1 during a day’s run. Run times varied between 20 min and 1 h, depending on the time required for the signal to return to baseline.
Peak areas were obtained by integrating the signal (volts) numerically over time.
This time-averaged peak area (V min) was converted to a flow rate-averaged peak area
(V ml) by multiplying by the average flow rate. This was measured for each peak by collecting the effluent volume and dividing by the time over which the volume was collected. Peak areas were converted to energy units (Joules) by comparison with peaks generated with a calibrating resistor located within the flow stream and immediately upstream from the column (Fig 2-1). Voltage and current for the heat pulses were measured and the heat input was then calculated from the relation Q(Joules)=V×A× t where V is the voltage, A is the amperage, and t is the time, in seconds, that the resistor was energized.
Sensitivity and Precision
A typical peak-to-peak noise level for the flow calorimeters used in this study was
2 mV or less. The thermistors had a nominal resistance of 10 kΩ at 25°C and a temperature coefficient of resistance of about -400Ω °C-1. Assuming an acceptable
signal-to-noise ratio for peak detection to be about 5, and using the measured gain for the
instrumentation amplifier, the calculated sensitivity for the flow calorimeter corresponded
to a temperature change of about 10-5 °C. This is similar to that for a true
9
microcalorimeter, which can also detect a temperature change on the order of 10-5 °C
(Steinberg, 1981).
Figure 2-2 shows peak areas that were generated by 45 mJ heat pulses and plotted as a function of flow rate. Peak areas obviously depended on flow rate. This dependence was taken into consideration when comparing heat data obtained at different flow rates by applying a correction factor that was based on the linear relationship in Figure 2-2.
60
50
40
30 y = 139.74x - 5.525 R2 = 0.868 20
Peak area (V mL) 10
0 0.25 0.27 0.29 0.31 0.33 0.35 0.37 0.39 0.41 0.43
Flow rate (mL min-1)
Figure 2-2. Example of a curve depicting the linear relationship between peak area size and flow rate.
Heat pulses for calibrating the instrument generally ranged from about 5 mJ to more than 100 mJ in size, corresponding to times of 2 to 45 s for energizing the calibrating resistor. Calibrating resistors were about 75kΩ and were energized at 15.0 V resulting in a power output of about 3 mW. Figure 2-3 shows a series of heat pulses and the associated calibration curve, the latter corrected for differences in flow rate. Linear regression of heat pulse peak areas versus energy input consistently gave R2 > 0.99.
10
0.3 70 A B y = 0.6965x 60 R2 = 0.999 0.2 50
0.1 40
Volts 30 0
Peak Area (V mL) 20 -0.1 10
-0.2 0 0 5 10 15 20 25 30 35 020406080100 Time (min) Energy input (mJ)
Figure 2-3. Peaks obtained with various size heat pulses (A) and the associated calibration curve (B).
Precision for replicated heat pulses was acceptable, with coefficients of variation that were generally less than 5%. Precision for exothermic and endothermic peaks obtained on a single column was of a magnitude similar to that for heat pulses. Precision obtained with replicated samples over periods of several months using different flow calorimeters was generally less than 15%.
Interpreting a Heat Signal
The information carried within a heat signal extends beyond simple thermal measurements. Interpreting and decoding this plethora of information is by no means simple and straightforward. In a paper entitled “What You Can Do with Surface
Calorimetry” (Steinberg, 1981), the author illustrates the kind of information that can be obtained from flow calorimetry. Figure 2-4 shows an exotherm adsorption peak (A) followed by the corresponding endotherm desorption peak (A’). The areas are equal
11
when desorption is complete. Also illustrated in the figure is a calibration peak (2)
obtained by passing a timed current through a calibrating resistor.
Figure 2-4. Pulse mode-physical adsorption and system calibration (Steinberg, 1981).
The author argued the validity of the following statements:
• The initial slope provides some indication of the rate of reaction.
• The peak width and shape is an indication of the uniformity of the surface site energies.
• Peak shape is also affected by flow rate in that higher pumping speeds result in sharper peaks.
• Particles that are too large or loosely packed contribute to poor peak definition.
• Areas under the curves are proportional to the strength of interaction
Steinberg (1981) compares physical adsorption to chemical adsorption or
chemisorption, showing that the heats of adsorption from either physisorbed or
chemisorbed solutes result in curves of similar shape (Figures 2-5 and 2-6), with areas under the curve proportional to the strength of interaction. However, if the solute is
chemisorbed, little or no desorption takes place (Fig 2.6).
12
Figure 2-5. Saturation mode- physical adsorption and desorption by carrier fluid (Steinberg, 1981).
Figure 2-6. Saturation mode-chemisorption (Steinberg, 1981).
13
Operation
Approximately 15 mg of amorphous aluminum hydroxides (AHO) were packed inside each column assembly and the total dry weight of the column piece recorded before placing inside the calorimeter. Our calorimeters were equipped with 4 different input lines merging at a central manifold connected to the column. It was possible to expose the material inside the column to any particular order of four different solutions.
Additionally, a bypass line outputting from the manifold allowed solutions in the lines to be changed without going through the column assembly. Columns were initially saturated with a solution until a steady baseline was obtained. The input solution was then changed and the heat of exchange for that particular reaction was recorded. The return of the signal to baseline indicated the end of the reaction. Depending on the particular design of the experiment, a different solution was then introduced to the column. After the last treatment, the column was removed from the calorimeter. The column was reweighed to account for entrained solution. The AHO sample was then removed from the column for further analysis.
CHAPTER 3 SAMPLE SYNTHESIS AND CHARACTERIZATION
The amorphous aluminum hydroxides, hereafter referred to as AHO, used throughout this study were synthesized in our laboratory.
Materials
Water used in all experiments was double distilled deionized water (DDI) obtained on a Heraess two stage quartz still. The DDI had a conductivity ranging from between 6 and 10 µmhos. All reagents and salts used, unless otherwise specified, were of reagent grade or better and were used as received upon purchase.
Synthesis And Preparation of AHO
Two types of AHO were prepared. The first was obtained by titrating a solution of
AlCl3 with NaOH to pH 6 and allowing the suspension to stand overnight. The gel-like solid was separated from the solution phase by centrifugation and dried in an oven overnight at 70°C. Three different batches were synthesized using this methodology and were designated as batch 1 (B1), batch 2 (B2) and batch 3 (B3). There was no precaution to wash salts prior to drying. NaCl precipitates were identified along with the AHO as described later in this chapter. The second type of AHO was obtained according to the recipe of Sims and Bingham (1968) and was designated as batch 4 (B4). The procedure called for 200 ml of 1.5 M AlCl3 to be neutralized by slowly adding 400 ml of 2 M
NaOH. The precipitate formed was filtered after 1 hour, resuspended in 1 liter of distilled water, and filtered again. The precipitate was then washed 3 times by suspending in 95 % ethanol and filtering. The last ethanol wash was removed by filtering on a buchner funnel
14 15
after which the aluminum hydroxides were dried in an oven at 70 °C for 24 hours. After
removal from the oven, all four batches of AHO samples were crushed, passed through a
150 µm mesh sieve and stored in a glass scintillation vial until future use. The effect of aging on the AHO was established calorimetrically and will be discussed in later chapters.
Physical Characterization
X-Ray Diffraction Analysis
X-ray diffraction (XRD) analyses were acquired on a computer-controlled x-ray diffractometer equipped with stepping motor and graphite crystal monochromator.
Samples were mounted on a low background quartz crystal mount and were scanned at 2°
2Ө per minute using Cu Kα radiation. The XRD patterns for B1, B2, and B3 revealed
both the noncrystalline nature of the AHO and the presence of crystalline NaCl.
Comparisons with patterns obtained on washed AHO samples from the same batches
showed there was no salt occluded inside the solids (Figs 3-1 to 3-3). Batch 4 XRD
showed patterns characteristic of Bayerite, a polymorph of crystalline Al(OH)3. As the
synthesis method for B4 called for three washes with 95 % ethanol to ensure salt
removal, there were no differences between washed and unwashed XRD data (Fig 3-4).
Thermogravimetric Analysis
Thermogravimetry was conducted using a computer-controlled thermal analyzer operating at a heating rate of 20 °C min-1. Data obtained confirmed the hydrated
nature of the AHO (Al(OH)3.xH2O) (Fig 3-5).
Surface Area Measurement
Specific surface areas (Table 3-1) were obtained from a multipoint Brunauer-
Emmet-Teller (BET) N2 adsorption isotherms obtained with a Quantochrome Autosorb
16
1C-Ms surface area analyzer. Samples were degassed prior to measurements for 4 hours
at 70 °C.
Table 3-1. Specific surface areas of the amorphous aluminum hydroxide batches Batch 1 Batch 2 Batch 3 Batch 4 ------m2 g-1 ------S.S.Aa 212 114 64 443 a specific surface areas
Values obtained were in agreement with reported values in the literature for
amorphous aluminum oxides (Goldberg et al., 2001). The variation among batches was
expected as it represents an inherent characteristic of amorphous hydroxides surface areas
that are affected by various factors including drying, reactions in the aqueous solutions,
and synthesis method.
Scanning Electron Microscopy
A scanning electron microscope (SEM) was used to study the surfaces of the AHO.
Samples from each batch were fixed on a carbon stub and coated with a very thin layer of
carbon to prevent electron build up on surfaces. For each batch, micrographs taken at
three different magnifications, 140 (200 µm scale bar), 350 (100 µm scale bar), and
35,000 (1 µm scale bar), are presented (Figs 3-6 to 3-9).
Micrographs obtained were very similar to those published in the literature for amorphous aluminum hydroxides (Goldberg et al., 2001). The crystal appearance, at lower magnification, has been attributed to the existence of dimmers and longer chains organized into a card house structure with internal voids (Bagnall et al., 1990). This phenomenon is due to the dominance of repulsive forces during the specific precipitation
environment. Micrographs obtained at higher magnifications reveal the spongy, porous
nature of the AHO, lacking any significant crystalline organization as proven by the
17
diffraction patterns. Published characterizations of AHO (Goldberg et al., 2001) have
established the extensive internal porosity of these solids via a number of methods such
as mercury intrusion porosimetry and nuclear magnetic resonance spectroscopy. While
we did not conduct any additional experiments, besides SEM, to document and measure
the internal porosity of our solids, the chemical data collected throughout this study
substantiated that it was indeed the case.
Figure 3-1. XRD pattern for batch 1 AHO A) untreated samples, and B) samples washed with DDI.
18
Figure 3-2. XRD pattern for batch 2 AHO A) untreated samples, and B) samples washed with DDI.
Figure 3-3. XRD pattern for batch 3 AHO A) untreated samples, and B) samples washed with DDI.
19
Figure 3-4. XRD pattern for batch 4 AHO A) untreated samples, and B) samples washed with DDI.
Figure 3-5. TGA patterns for batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4).
20
Figure 3-6. Scanning electron micrographs showing AHO batch 1 morphology at 3 different scales.
21
Figure 3-7. Scanning electron micrographs showing AHO batch 2 morphology at 3 different scales
22
Figure 3-8. Scanning electron micrographs showing AHO batch 3 morphology at 3 different scales
23
Figure 3-9. Scanning electron micrographs showing AHO batch 4 morphology at 3 different scales
24
Chemical Characterization
Chemical Composition
The Aluminum (Al) content of the four batches was measured on digested samples.
A known amount of each batch was weighed into a small scintillation vial and around 1 ml of DDI and four drops of concentrated HCl added. The vials were tightly sealed and dropped into a water bath on a heating plate. Samples were left to digest in the boiling
water until all solids disappeared from solution, usually around 30 to 45 minutes. Batch 4 samples consistently took longer to digest with the solution, never totally clearing up
even after more than 12 hours on the hot plate. This occurrence might be caused by the
presence of the more resistant Bayerite in the samples as documented in the XRD
patterns.
Aluminum content was also measured for the AHO in each column used in the
calorimetry experiment. This provided additional replicates for each batch. A summary of
the Al contents of the AHO is Table 3-2. These means were based on the total weight of
sample, which included moisture and, for B1, B2 and B3 NaCl. By comparison, the Al
content of pure, anhydrous Al(OH)3 is 34.6 %.
Table 3-2 Aluminum content of the amorphous aluminum hydroxide batches Batch 1 Batch 2 Batch 3 Batch 4 ------Al % ------Average 12.80 18.9 16.3 19.9 Std Dev a 2.89 3.72 2.3 4.11 n b 12 15 12 6 a Standard deviation b Number of samples included in calculations of average and standard deviation
Anion Exchange Capacity
The anion exchange capacity (AEC) of the AHO was determined by a modified ion
adsorption method using Cl and NO3 anions. The samples were placed inside a column
25
assembly and the total dry weight of the piece recorded. Samples were then leached with
50 mM KNO3 solution for 40 minutes. This time was verified beforehand calorimetrically to be sufficient for the nitrate anions to saturate the exchange sites. The
final wet weight of the piece was obtained and samples emptied into a small scintillation
vial and digested as described in the previous section. Nitrate was measured by ion
chromatography (IC). Nitrate measurements were corrected for entrained solution based on the difference between dry and wet column pieces weights; and the resulting values used to calculate the AEC of the solids (Table 3-3).
Table 3-3 Anion exchange capacity of the amorphous aluminum hydroxides Batch 1 Batch 2 Batch 3 Batch 4 -1 ------cmol(+) kg ------Average 94 (264)a 104 (198) 111 (246) 131 (237) Std Dev b 5.88 10.43 2.54 22.50 n c 6 5 3 3 a Numbers in parenthesis are expressed on the basis of kg of pure Al(OH)3,i.e., corrected for salt and moisture contents b Standard deviation c Number of samples included in calculations of average and standard deviation
Based on averages of aluminum content and anion exchange capacity, the following Al:Cl mole ratios were calculated for the four AHO batches (Table 3-4.)
Table 3-4. Average Al:Cl mole ratio calculated for the four AHO batches Batch 1 Batch 2 Batch 3 Batch 4 ------mole ratio------Al:Cl 5:1 6.75:1 5.4:1 5.6:1
Cation Exchange Capacity
At pH values above the zero point of charge, the AHO surface was negatively
charged and exhibited cation exchange capacity (CEC). When relevant, the CEC was
measured in a similar fashion to the AEC as described in the section above. Potassium
and Na were used as the exchanger cation pair. Potassium was measured using atomic
26
absorption spectrophotometry (AA). Similarly, K values were corrected for entrained
solution and net values used to calculate the CEC. The 50 mM KNO3 leaching solution
was adjusted to match the pH at which the specific experiment was conducted..
In the few instances where the AHO surface exhibited both AEC and CEC
simultaneously, it was possible using the method described above to measure both on the
same sample digest.
Effect of pH on Aluminum Content
Some experiments necessitated the use of various pH treatments, namely pHs 8.0
and 10.5. The objective of this test was to examine the effect of these pH treatments,
specifically the higher ones, on the dissolution of Al from the AHO. A known amount of each batch (B2, B3, and B4) was weighed into a scintillation vial and about 20 ml of
solution added to it. Three pH treatments were tested: 5.7, 8.0 and 10.5, in addition to a
DDI control. Each treatment was replicated three times. The vials were shaken for about
5 minutes and left to stand for one day. The supernatant was filtered and analyzed for Al.
Comparison to the original Al content for each batch (Table 3-2) yielded the percentage
of Al lost. The data are presented in Table 3-5.
Table 3-5 Effect of various pH treatments on the percent loss in Al content in AHO batches Batch 2 Batch 3 Batch 4 average std deva average std dev Average std dev ------% Al ------DDIb 4.00 0.19 4.93 0.02 15.33 0.68 pH 5.7 1.88 0.29 2.54 0.29 10.41 0.19 pH 8.0 0.00 0.00 0.00 0.00 0.00 0.00 pH 10.5 2.95 0.40 2.07 0.27 0.00 0.00 a Standard deviation b Double distilled deionized water.
27
Subjecting the AHO to higher pHs (8.0 and 10.5) did not cause greater Al loss from the solids when compared to either the control or our usual working pH of 5.75. It was hence deemed safe to carry out the experiments with higher pHs as planned.
CHAPTER 4 CHEMICAL CHARACTERIZATION USING FLOW ADSORPTION CALORIMETRY
One purpose of this work was to demonstrate the application of flow adsorption calorimetry (FAC) to the study of surface chemical properties that cannot be obtained readily by other methods. At the same time, by allowing multiple treatment cycles to be applied on the same sample, FAC can be employed to study the enthalpic and physical effects of various treatments on surface properties. In this chapter, we report first on the direct measurement of the calorimetric effects of ion adsorption. Second we present some experiments that were devised to (i) measure the point of zero charge calorimetrically and
(ii) examine the reversibility of surface charge with pH
Effect of Carbonate
All the experiments in this study were conducted in the presence of carbonate.
Solutions and samples were exposed to the atmosphere with no attempt to control or limit the carbonate adsorption that occurred. Numerous studies in the literature have studied the competitive adsorption of carbonate on the surface of the AHO and evaluated its effect on these surfaces (Schulthess and McCarthy, 1990; Su and Suarez, 1997). As described in these papers and other references, however achieving an environment that is absolutely free of carbonate is extremely difficult, and requires instrumentation that was neither available in our lab nor easily adapted to the calorimeters. Hence, it was decided not to attempt to get rid of carbonate; in as much as carbon dioxide and carbonate are
28 29
ubiquitous in soil and aquatic environments, it was deemed more realistic to incorporate
them.
Ion Exchange
Obtaining Heats of Ion Exchange
Approximately 15 mg of AHO were placed inside the calorimeter, and equilibrated
with 50 mM NaCl until a steady baseline was obtained. The input solution was then
changed to 50 mM NaNO3, which resulted in an exothermic heat of exchange as NO3
displaced exchangeable Cl. When the signal returned to baseline, the solution was then
changed back to 50 mM NaCl, and the endothermic peak associated with the
displacement of exchangeable NO3 by Cl was recorded. Several cycles of Cl/NO3
exchange were thus recorded. Similarly, the heats of cation exchange were obtained by
changing between 50 mM NaNO3 and 50 mM KNO3 solutions. K replacing Na on the exchange sites was an exothermic reaction , while Na displacing K resulted in an endothermic heat of exchange.
The ionic strength (I) and pH of each pair of solutions were maintained constant and matching throughout the duration of one experiment. As our calorimeters are equipped with four input lines, it was possible to measure heats of both cation and anion exchange on the same column sample sequentially and repeatedly. In the later case, either the cation or the anion pair was changed between two consecutive solutions. The solution sequence 50 mM NaCl, 50 mM NaNO3, 50 mM KNO3, 50 mM NaNO3 and 50 mM NaCl
would have yielded one cycle of Cl/NO3 exchange and one cycle of K/Na exchange.
30
Results
Anion exchange
Figure 4-1 shows the results for Cl/NO3 exchange obtained on all four batches of our AHO. The two peaks were always equal in area, as expected for a reversible ion- exchange reaction. In addition, the reaction was rapid as it took around 20-30 minutes for the signal to return to baseline, indicating the end of the reaction.
Figure 4-1. Heats of Cl/NO3 exchange obtained at pH 5.7 on batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4).
The results obtained were reproducible over time and samples. Figure 4-2 presents
replicates of Cl/NO3 exchange collected on a B1 sample over the span of one week. As
shown, the quantity of heat associated with anion exchange remained relatively constant.
This also attests to the lack of significant effect of aging on the AHO surface over the
31
period of weeks. Similar results (data not shown) were documented on other AHO batches for periods up to 2 or 3 weeks.
Peak areas replicated over calorimeters and samples differed due to differences in solution flow rates, electronic gains of different calorimeters, and sample size. These factors were corrected for, resulting in a precision of less than 10%.
s 0.15 Volt day 2 day 3 0 day 5 day 6
-0.15
-0.3
-0.45 0153045 Time (mn)
Figure 4-2. Replicates of Cl/NO3 heats of exchange obtained at pH 5.7 on batch 1 over a period of 6 days.
Integrated peak areas in units of V ml are presented in Table 4-1. These peak areas were converted to energy units (mJ) by comparison with peak areas generated using heat pulses of known size.
Using the measured anion exchange capacity for each batch (Table 3-3) the Cl/NO3
-1 heats of exchange were converted to kJ molC . Approximately 3 cycles of Cl/NO3
were obtained for each calorimeter column. An average of the six peak areas (3 exotherms + 3 endotherms) were used to calculate the ∆H for a specific column. Table 4-
2 shows the ∆H values for each batch averaged over all columns.
32
Table 4-1. Peak areas associated with Cl/NO3 exchange peaks of batches 1, 2, 3, and 4 shown in figure 4-1. Batch 1 Batch 2 Batch 3 Batch 4 Cl NO3 Cl NO3 Cl NO3 Cl NO3 P.Aa -53.13 55.43 -65.20 64.66 -58.73 55.84 -38.93 42.85 a Peak areas integrated in V ml
Table 4-2. Heats of Cl /NO3 exchange measured calorimetrically on different batches of amorphous aluminum oxides. Batch 1 Batch 2 Batch 3 Batch 4 -1 ------kJ molC ------Average 5.00 5.81 3.55 3.84 Std Dev a 0.82 1.49 0.55 0.11 n b 15 14 8 2 a Standard deviation b Number of columns included in calculations of averages and standard deviations
Cation exchange
There was no detectable heat of K/Na exchange on samples from any batch at pH
5.75, indicating that the AHO surfaces bore no negative charge at this pH,i.e., all valence-
1/2+ unsatisfied terminal hydroxyls groups at the surface are protonated Al-OH2 . The
surfaces of the AHO exhibit cation exchange capacity only at pH values close to or above
the ZPC. Therefore, K/Na exchange data was only available when the pH of the solutions
was adjusted to appropriate values. When it occurred, the K/Na exchange exhibited
similar characteristics to that of the AEC (Fig 4-3) in that it was rapid, reversible and
reproducible over time and samples.
The areas of the peaks, the CEC and the associated heats of exchange are
presented in Table 4-3. These were expressed in kJ mol-1 when measurement of the
cation exchange capacity for the column was possible. While reproducible peak areas for
the K exotherm and Na endoderm were measurable, it was not always possible to obtain
the corresponding analytical measures of exchangeable K and Na.
33
Figure 4-3. Heats of K/Na exchange obtained at pH 10.5 on batch 2 (B2), batch 3 (B3) and batch 4 (B4).
This is because flow calorimetry is more sensitive than either atomic absorption
spectrophotometry or ion chromatography. Heats of K/Na exchange measured at pH 10.5 were less than those for Cl/NO3 exchange measured at pH 5.75.
Table 4-3. Peak areas, CEC and heats of K/Na exchange obtained a pH 10.5. Peak Areas CECa ∆H -1 -1 -1 Vml cmol(+) kg kJ molC K exotherm Na endotherm Batch 2 1.42 -1.49 7.20 1.15 Batch 3 4.04 -1.8 n.ab n.a Batch 4 5.23 -6.59 38.35 1.61 a Cation exchange capacity b not available
34
Zero Point of Charge
Rationale
Measurements of the Zero Point of Charge (ZPC) of AHO traditionally have been conducted by electrokinetics experiments or potentiometric titrations. Electrokinetics experiments are based on the mobility of charged particles in an electrical field, referred to as electrophoretic mobility (EM) (Findlay et al., 1996; O'Brien et al., 1995).The ZPC is usually obtained by determination of the pH at which particles exhibit zero EM (Sposito,
1989). Potentiometric titrations theoretically determine the point of zero salt effect
(PZSE). However, for an AHO immersed in 1:1 electrolyte solution composed of non- specifically adsorbing ions, which do not form either inner-sphere or outer sphere surface complexes, PZSE=ZPC.
The rationale behind the calorimetric determination of the ZPC is that the heats of
ion exchange are directly related to surface charge. When heats of both anion and cation
exchange (AEC/CEC) are measured at different solutions pHs, the ZPC of the AHO is
obtained. This method offers several advantages over existing methods. It is quicker, gives better control over the ionic strength of the solution (I), and allows the ZPC to be
determined on a single 15 mg sample. While flow adsorption calorimetry has been used
in the past to determine surface areas, to our knowledge this is the first report of
experiments using flow calorimetry to determine the ZPC.
Procedure
Approximately 15 mg of AHO were packed inside the calorimeter. The heats of
Cl/NO3 and K/Na exchange were measured as described in the section above with pH 5.7
solutions. The same measurements were repeated with pHs 8.0 and 10.5. Solutions in the
lines were changed using the bypass line while the sample remained in the column. The
35
NaCl and NaNO3 solutions were adjusted to pH 8.0 using 1mM NaHCO3 while 1mM
KHCO3 was used for the KNO3 solution. Similarly, pH 10.5 was obtained using either a 1
mM Na2CO3 or K2CO3 buffer. Changes in the background solution ionic strength (I) due
to the pH adjustments were the same for all input solutions and amounted to ≤ 2 %. The
pH of the solutions collected at the outlet side of the calorimeter after passing through the
column was monitored until desired pH values were reached. This ensured the AHO
surfaces were at the target pHs. At the end of the pH 10.5 runs, the AHO sample was
recovered from the column for AEC and CEC determinations. These measurments were
obtained according to the digestion method described in Chapter 3.
Results
Figure 4-4 shows the Cl/NO3 and K/Na peaks obtained at the different pH values.
As the pH is increased to pH 8.0, the Cl/NO3 peaks decreased in size, indicating a
reduced number of positively charged sites available on the surface. At pH 10.5, the absence of peaks for Cl/NO3 exchange suggested that the surface no longer bore positive
charge. The behavior of the K/Na exchange completes the picture. At the starting pH of
5.75, there was no cation exchange as witnessed by the flat line.
The K/Na exchange peaks increased with increasing pH. The surface of the AHO
was thus becoming gradually more negatively charged as pH increased. It is worthwhile
mentioning that while we were able to measure K/Na exchange at pHs 8.0 and 10.5, we
encountered more difficulties obtaining well-defined returns to baselines, with a higher
incidence of baselines shifts for K/Na exchange.
36
Figure 4-4. A) Cl/NO3 and B) K/Na heats of exchange obtained on AHO batch 1 at different pHs.
Table 4-3 summarizes the effect of pH on AEC and CEC for a B2 column at the end of a ZPC determination. Similar data were obtained for other B2 columns and for columns of the other batches.
Table 4-3. Data collected for the determination of the point of zero charge of a B2 column. pH AECb CECc AEC CEC AEC CEC -1 -1 -1 Vml cmol(+/-) kg kJ mol 5.7 64.95 0 92.5 0 3.5 -- 8.0 11.77 0.8 --a ------10.5 0 4.7 0 7.2 -- 1.2 a not determined b average values of NO3 exotherms and Cl endotherms peak areas c average values of K exotherms and Na endotherms peak areas
37
It is apparent from the data in Table 4-3 that the heats of exchange for AEC and
CEC were different, i.e., a given amount of AEC at pH 5.75 released more energy than
the same amount of CEC at pH 10.5. If the ZPC is to be determined from calorimetric
data, the peak areas for either AEC or CEC must be corrected to account for these energy
differences. Therefore, the peak areas for CEC were corrected for differences in heats of
exchange, guaranteeing that a unit area associated with cation exchange was equivalent to
that for anion exchange. Corrected peak areas for B2 and B3 are plotted against pH in Fig
4-5. The ZPC was taken as the pH at which the two corrected heats of reaction were
equal. The ZPC for B2 and B3 fell at 9.5 and 9.4 respectively and are indicated in the
figure by the diamonds. These ZPC values are in agreement with reported values in the
literature for amorphous AHO (Goldberg and Johnston, 2001; McBride, 1994).
Figure 4-5.Corrected heats of cation and anion exchange for batches 2 and 3 of AHO as a function of solution pH, measured by flow adsorption calorimetrically. The diamonds indicate the ZPC.
38
Charging of AHO surfaces
Our experiments on determining the ZPC calorimetrically revealed valuable information about the charging mechanisms of these surfaces. Traditionally two models have been formulated to describe the charging of oxides surfaces, and although they were
developed based on crystalline surfaces, they are widely adopted as responsible for
charging of amorphous oxide surfaces as well. In the 2 pK model (Schindler and Stumm,
1+ 1987), the surface S–OH2 is subject to two pK values, controlling the formation of the neutral S–OH group and the negative S–O-group as follows:
0 + + S–OH + H ↔ S–OH2 Ksa1
S–O- + H+ ↔ S–OH0 Ksa2
According to this model, a “charge neutral” surface can exist, i.e. S—OH0, one that has
neither positive nor negative charges. It is now accepted, however, that it is the dangling
valence unsatisfied OH groups (i.e., bonded to a single Al) that are the most reactive and
responsible for surface charge. The mean charge per Al–O bond according to Pauling’s
valence rule and assuming Al in octahedral coordination with 6 oxygens is +3/6 = +1/2,
resulting in a charge of +1/2 or -1/2 on this type of group . Consequently, some authors
(Bolt and Van Riemsdijk, 1982; Hiemstra et al., 1989) have argued for a different scheme
of charge development: a “one-pK” model that uses one affinity constant instead of two to describe the charging behavior of a metal oxide. The basic charging equation of the one-pK model is:
1/2 - + 1/2+ S–OH + H ↔ S–OH2 KH
As follows from this reaction, one or the other of these two chemical forms exist on the
surface at all times depending on the solution pH. An electrically neutral surface is one
with equal numbers of positive and negative charge, but a “charge neutral” surface, i.e.
39 one that carries no charge, is not possible with this model (Hiemstra et al., 1999;
McBride, 1994).
As evident from the ZPC curves in Figure 4-5, the decline in AEC was not matched by an increase in CEC, which is required if the one pK model accurately describes surface charging of the AHO. Considering the detection limit of our calorimeters, the absence of CEC at pH 5.75 or AEC at pH 10.5 is a definite sign that S–
OH 1/2 + sites are not converting to S–OH 1/2 -. Instead, the calorimetric data indicate that the surface of AHO is nearly charge neutral, i.e., neither positive or negative charges exist at pHs around the ZPC. This finding is consistent with the 2 pK model.
Additional Observations
In the process of determining the ZPC of different AHO batches calorimetrically, we observed several unexplained phenomena. Although no additional experiments were devised to specifically investigate these trends, they were replicated enough over time and samples to believe they were not artifacts. Reporting these observations not only exposes the nature of information accessible exclusively by flow adsorption calorimetry, but also opens up new dimensions for researching and understanding AHO surfaces.
Flip-Flop effect
As discussed earlier in this chapter, we measured cation exchange capacity of AHO surfaces by looking at K/Na exchange. On all batches, K exchanging with Na was generally exothermic and Na replacing K endothermic (Fig 4-4). With batch 4, however, the heats of exchange reversed sign at pH 8.0: K replacing Na was endothermic while Na replacing K became exothermic. Although reproducible over multiple K/Na exchange cycles at pH 8.0, the heats of exchange returned to their original signs at pH 10.5. While
40
the heat for K/Na exchange “flip-flop” was observed only for batch 4, an AEC “flip-flop”
was observed for one B3 column at pH 8.0 after an arsenate treatment.
Calorimetric measurements capture the difference in surface preference toward a
pair of ions. For ions of equal charge, the energy of exchange is largely determined by the
effective radius of the ion and the energy of hydration. The surface behavior toward ion
exchange is often divided into two general cases (McBride, 1994): (i) the weak field case
when the surface charge is located beneath the surface and the difference in hydrated
radius between two ions controls the energy of exchange, and (ii) the strong field case when the charge is near or at the surface, ions are in direct contact with the surface and the order of selectivity for ions, and hence the energy of exchange, depends on the ionic
(or dehydrated) radius. While weak field is accepted to describe the behavior of the permanent-charge expandible phyllosilicates, it is recognized that some oxides exhibit strong field behavior. Moreover, it is recognized that the behavior depends on the specific oxide, type of electrolyte and state of hydration of the electrolyte ions (Sverjensky, 2001).
The “flip-flop” effect seems to suggest that the surface of the AHO changes from a weak field to a strong field exchanger in the vicinity of the ZPC. Understanding this phenomenon required research that was beyond the scope this study.
Change of anion exchange energetics with pH
Another trend we observed was the change of anion exchange energetics with pH on B2 columns. Obtaining the ∆H for an exchange reaction on a kJ mol-1 basis required
knowledge of the quantity of heat involved and the magnitude of the exchange. For every
column on which Cl/NO3 peak areas were measured, it was possible to calculate the heat
of anion exchange assuming the AEC values for clean batches reported in Table 3-2.
41
These were later averaged and presented in Table 4-2. These ∆H values were initial heats
of reaction as they were obtained on “clean” columns, i.e, before any treatment
application. After exposing the AHO to various treatments in the calorimeter, Cl/NO3 peak areas were re-measured calorimetrically, the AHO recovered and AEC measured.
This yielded final heats of reaction.
Throughout a number of experiments that required raising the pH of the input solutions, we collected initial and final heats of anion exchange. The anion exchange values obtained on B2 columns exhibited a consistent trend whereby the ∆H values decreased on average from a value of 7.00 kJ mol-1 at pH 5.75 to 2.60 kJ mol-1 at pH 7.5.
The decrease was significant and indicated that the energetics of Cl/NO3 exchange were
affected by pH.
Since calorimetric measurements capture differences in preference between a pair of ions,
the pH increase is definitely changing ion preferences. Determining the underlying
mechanism whereby this occurs will require extensive research but will undoubtedly
reveal valuable information on the AHO surface charging behavior. Again, this was
beyond the objectives of this work.
For the calorimetric determination of the ZPC, the change in the heats of anion
exchange was not accounted for. This failure to account for heats of anion exchange,
however, did not affect ZPC values since the change in heats of exchange did not alter
the value of the pH at which the surface becomes neutrally charged; at this point there
will be no detectable heats of exchange. It will however affect the slope at which the
anion curves approach the ZPC value.
42
Reversibility of Surface Charge with pH
Rationale
Demonstrating the reversibility of surface charge of an AHO is a difficult task if traditional methods are employed. The amounts of acid and/or base required to adjust the pH of one sample in one direction and then the other significantly affect the ionic strength (I) of the background solution. The charge development on these amphoteric surfaces is dependent on the electrolyte level of the bathing solution. It would be improper to draw conclusions on the reversibility of the charge while I is not constant.
In the experiments to measure ZPC the pH of the solutions was increased while maintaining the background I constant. The underlying principle in this study was to reapply the pH treatments on the same AHO samples in the reverse order and substantiate the reversibility of AHO surface charge with pH. To our knowledge, such a measurement has never been reported in the literature.
Procedure
Approximately 15 mg of AHO was packed inside the calorimeter and taken through the procedure described earlier for the ZPC determination. At the end of the pH
10.5 measurements, the pH treatments were reapplied in the reverse order and heats of exchange re-measured. The column was flushed with a particular pH solution until the pH measured at the outlet side corresponded to the desired value. After each calorimeter experiment, the AHO sample was recovered for AEC determination.
Results
As illustrated by Fig 4-6, the surface charge as measured by calorimetric peak areas for B2 and B3 samples was completely reversible with pH. The peak areas for Cl/NO3 exchange obtained at different pH values in both the ascending and descending order
43
were equal. Note that the pH only reached a value of around 7.0 on the descending leg.
Despite flushing the sample with the unbuffered pH 5.75 solutions for several days, the pH in the effluent did not drop below 7.0. Lowering the pH of the AHO columns using unbuffered salt solutions is apparently difficult and time consuming.
Actual AEC measurements are shown in Fig 4-7 and confirm the reversibility measured calorimetrically. The AEC values in Fig 4-7 were obtained on AHO samples
from calorimetry runs, insuring that the AHO was treated identically in both experiments.
The correspondence of the values, i.e. calorimetric peak areas and measured AEC, not only supports the reversibility of the surface charge, but also attests to the stability of the
AHO surfaces when exposed to hundreds of pore volumes of varying pH solutions.
Figure 4-6. Effect of ascending and descending pH on calorimetric peak areas of batches 2 and 3 of AHO measured by flow adsorption calorimetry.
44
Figure 4-7. Effect of ascending and descending pH on the AEC of batches 2 and 3 of AHO samples. .
CHAPTER 5 ENERGETICS OF ARSENATE SORPTION
In previous chapters, we demonstrated the use of flow adsorption calorimetry to the study of surface chemical properties of amorphous aluminum hydroxides (AHO). This chapter describes experiments using flow adsorption calorimetry to examine the physical and chemical reactions involved in arsenate sorption by AHO. According to Rudzinski et al. (1999b), the calorimetric effects of adsorption are much more sensitive to the nature of an adsorption system than adsorption isotherms. They state that knowledge of the enthalpy effects accompanying the adsorption of ions, may bring more light onto fundamental features of these adsorption systems. Traditionally, sorption isotherms have been used to obtain information about the rate and amount of solute sorbed over time.
They give no insight, however, into the mechanism by which sorption occurs (Sposito,
1984). According to Steinberg, (1981) flow calorimetry can (i) resolve a complex series of reactions that occur more or less simultaneously but at different rates, and (ii) be applied to multiple adsorption/desorption cycles on the same sample, allowing reversible and irreversible processes to be distinguished.
Obtaining Heats of Reactions
Columns containing about 15 mg of AHO were placed inside the calorimeter, and equilibrated with 50 mM NaCl until a steady baseline was obtained. The input solution was then changed to one in which 2 mmol L-1 of the NaCl had been replaced with
-1 1 mmol L of Na2HAsO4.7H20 i.e., keeping the total Na concentration at 50 mM and As
(V) at 1mM. The NaCl/Na2HAsO4 solution had been adjusted to the same pH (5.75) as
45 46
the NaCl. The AHO was exposed to the arsenated solution for a fixed period of time at
which point the solution was switched back to 50 mM NaCl. Further treatments,
depending on the experimental design for the particular sample, awaited the signal return
to baseline. This procedure allowed us to expose AHO to various durations of 1 mM
As(V), thereby varying the surface loading. At the end of the experiment, the AHO was
usually recovered from the column and digested. Arsenate was determined by ion
chromatography and/or graphite furnace.
Results
Calorimetric Effects
Heat signals generated by As adsorption by AHO are shown in Fig 5-1. The
reaction of arsenate with samples from all four batches of AHO was strongly exothermic.
Returning to the original NaCl solution after a 20 minute Na2HAsO4 treatment produced
no corresponding heat signal, indicating that the peak observed with Na2HAsO4 was not
caused by reversible Cl/HASO4 exchange. It is clear from the peak shapes that arsenate
reaction with the AHO differed kinetically from Cl/NO3 exchange (Fig 5-2) in that heat
evolved at a much slower rate than Cl/NO3 exchange and the reaction lasted considerably longer. The arsenate reaction was still going strong after 20 min as evidenced by the continuous evolution of heat. However, this heat evolution did not occur at the same rate throughout the reaction: it slowed down as the reaction progressed. If allowed to continue, the arsenate sorption reaction would take as much as an hour to return to baseline.
47
Figure 5-1. Heats of reaction of arsenate with batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4). The arrows indicate a change from arsenate solution back to NaCl.
This behavior was observed with all four batches of AHO. This change in rate of heat evolution could indicate separate reactions that occur at different rates. Similarly, phosphate sorption by soils has been found to exhibit an initial, rapid uptake phase, followed by a much slower rate of uptake. Several explanations have been offered to explain this slow uptake; they encompass (i) diffusion to less accessible sites within pores of solid aggregates (Lookman et al., 1994; Parfitt, 1989), (ii) penetration into the amorphous Fe and/or Al oxides by solid-state diffusion (Barrow, 1983; Van Riemsdijk et al., 1984), and (iii) precipitations with metals derived by dissolution of the soil matrix
(Lookman et al., 1994; Pierzynski et al., 1990).
48
The heat signature for arsenate sorption was repeated in additional arsenate
sorption cycles (Fig 5-3), even ones administered back to back and separated only with one NaCl treatment. If there are indeed different reactions associated with arsenate sorption, these must occur every time the surface is exposed to an arsenate sorption cycle.
Successive cycles of the same duration were often of similar size or only slightly smaller.
Figure 5-2. Contrast between peak shapes for Cl/NO3 exchange and arsenate reaction with an AHO batch 1sample. Identical results were obtained for all other batches. The arrows indicate a change from arsenate solution back to NaCl.
Figure 5-3. Two arsenate cycles on same sample of batch 1 of AHO. The arrows indicate a change from arsenate solution back to NaCl.
Figure 5-4 shows calorimetric results obtained for col 6B1 (see Appendix) that was exposed to 4 cycles of arsenate, each of which was allowed to run uninterrupted until the
49
signal returned to baseline, indicating the end of the arsenate sorption reaction. This
differed from our usual protocol in which we controlled the timing of the arsenate
exposure by switching to the 50 mM NaCl solution after 20 mn.
Figure 5-4. Four consecutives arsenate cycles and corresponding returns to NaCl solution collected consecutively (A, B, and C) and after 24 hours (D) on same AHO sample.
The sample in Fig 5-4 was initially equilibrated with a solution of 50 mM NaCl.
The first cycle of arsenate (Fig 5-4. A) returned to baseline in about 45 minutes and returning to a NaCl solution generated a very small endotherm. Figure 5-4. B and C show the resulting curves of the second and third arsenate cycles and the corresponding returns to the NaCl solution. The fourth arsenate cycle presented in figure 5-4. D was obtained after allowing the AHO to sit in 50 mM NaCl for 24 hours before applying the arsenate
50
The Cl endotherms obtained after the arsenate cycles were consistently smaller in
size than the arsenate exotherms, indicating a high degree of irreversibility in arsenate
sorption. Their presence, however, suggests a small reversible component in the arsenate
reaction with the AHO. The return to baseline during each arsnate cycle apparently
indicated that the arsenate sorption capacity had been exhausted. The consecutive
arsenate cycles had bigger peak areas than the corresponding Cl endotherms, signifying
that the heat signatures corresponded to additional arsenate sorption. While the first 3
arsenate cycles declined in size, the fourth cycle, obtained after a 24 hours rest, had a
bigger peak area than in the previous arsenate cycle.
Results obtained on another AHO sample (Col 25 B2) were in agreement with
above findings.
Figure 5-5. Two consecutive arsenate cycles (A, B) and corresponding return to NaCl solution (A) collected on same AHO sample. The arrow in B indicates a change from arsenate solution back to NaCl.
The first exposure to arsenate resulted in a peak (Fig 5-5. A) that returned to baseline in around 30 minutes when uninterrupted. A small endotherm was obtained as the solution was returned to 50 mM NaCl (Fig 5-5. A). A second arsenate exposure had a heat signature with a peak area larger than the previous endotherm.
51
The reoccurrence of an arsenate heat signal after the previous peak has returned to baseline suggests the existence of a mechanism by which a reactive surface is regenerated
perhaps due to a spatial rearrangement of arsenate on the AHO surface. That is, the
adsorbed arsenate species either diffuse along the surface to less accessible sites or into
the interior of the AHO. If this occurs, these latter changes yield no measurable heat of
reaction. The absence of a heat signature could indicate that the reactions involved (i) are
entropy driven, (ii) occur at a very slow rate whereby the heat generated is dissipated and
lost in the baseline noise, or (iii) consist of endothermic and exothermic reactions whose
heat signals cancel each other. Over time, sites that react with arsenate exothermically are
regenerated, explaining the peaks in the second and third cycles. The larger arsenate peak
obtained after 24 hours in NaCl suggests a time dependency with a greater extent of
redistribution happening as time goes by.
The ability of the arsenate to detach from the site at which it is attached and move
along the surface has never been explicitly stated in studies of oxyanion adsorption on
AHO, although it is a requirement for some suggested mechanisms to be achievable.
Whether it is solid state diffusion (Van Riemsdijk et al., 1984) or change from a
monodentate to a bidentate surface species with increasing surface coverage (Suarez et
al., 1998), arsenate mobility was the underlying implied mechanism. Many conceptual
and mathematical representations of multidentate adsorption on crystalline (hydr)oxides
surfaces encompass similar conceptual approaches (Benjamin, 2002; Hiemstra and Van
Riemsdijk, 1996).
Arsenate Sorbed
The amounts of sorbed arsenate on each batch varied with the duration, mode and
number of arsenate cycles to which the sample was exposed. Adsorbed arsenate
52 concentrations in relation to the above-mentioned factors are tabulated in Appendix A, along with all other measurements collected for each column. Ranges of arsenate loadings for each batch are presented in Table 5-1. Also included in the table are the corresponding Al:As mole ratios.
There was considerable variation in the amount of arsenate sorbed even holding the arsenate solution concentration and time of exposure constant. This finding highlights the amorphous nature of the surfaces, which seem to uptake and distribute arsenate differently from batch to batch and at each exposure.
Table 5-1. Arsenate loadings and corresponding Al:As mole ratios obtained on all four batches of amorphous aluminum hydroxides. Minimum Maximum As Al:As As Al:As µg g-1 mole ratio µg g-1 mole ratio Batch 1 6,000 24.50:1 31,000 38.70 Batch 210,200 35.30:1 39,100 13.90 Batch 311,700 36.40:1 67,300 8.29 Batch 4 22,200 24.20:1 --a -- a not available; one batch 4 sample arsenated.
Despite the high amounts of arsenate sorbed, the Al:As ratios suggest that the arsenate sorption capacity of the AHO has not been saturated yet. An indication of the total sorption capacity of amorphous Al and Fe oxides has been given by Schoumans and
Groenendijk, (2000). In their work related to modeling soil phosphorus (P) levels and leaching from agricultural lands in the Netherlands, they say the maximum P sorption capacity of noncalcalareous sandy soils is 0.5 times the oxalate extractable of Al and Fe, i.e., the amorphous components, on a mole basis. This translates into Al:As mole ratios as high as 2:1 which are far higher than the ratios we obtained for arsenate, believed to adsorb onto AHO by very similar, if not exact, mechanisms as phosphate. This
53
observation will be further examined in upcoming sections, as the effect of arsenate on
ion exchange will be studied.
Figure 5-6 shows an SEM-EDX elemental dot map of an arsenated B1 sample.
Light areas in the dot maps indicate elevated concentrations of a particular element. As it
can be seen on the figure, the arsenate rich area (lower left corners) mirrored the image of the particle studied shown in the top left corner. So did the Cl enriched zones presented in the lower right corner. While the elemental dot maps are not meant to provide evidence as to the mechanism of sorption, they are presented as a confirmation of the presence of
the arsenate on the AHO surfaces along with Cl, the Cl being present on the anion
exchange sites. It is also important to note the greater intensity of Cl dot maps compared
to the arsenate. This suggests that while high arsenate coverages were obtained, they do
not saturate the totality of possible exchange sites whose magnitude is indicated by the Cl
exchange capacity.
Figure 5-6. SEM-EDX elemental maps of two distinct particles, A and B, of an arsenated batch 1 AHO sample.
54
Heats of Reactions
Using the amounts of arsenate sorbed, the heats of reaction obtained were
expressed on a kJ mol-1 basis (Table 5-2). The reaction of arsenate with samples from all
four batches of AHO was strongly exothermic: ∆H values ranged from 3 to 60 kJ mol-1 of arsenate adsorbed with the majority of the values falling between 20 and 40 kJ mol-1. The
∆H values in relation to the amount of sorbed arsenate and the corresponding Al:As ratio for all columns in this study are presented in Table 5-2. There is no correlation between either the amount of sorbed arsenate or the Al:As molar ratio and ∆H. A decline in ∆H values, however, in each batch was always associated with decreasing Al:As ratios. This observation suggests that the heat of adsorption changes with the amount of arsenate adsorbed.
Table 5-2. ∆H values, amounts of sorbed arsenate and Al:As molar ratios for AHO samples from batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4). ∆H As sorbed Al:As Column Name kJ mol-1 µg.g-1 mole ratio Col 1 B1 64.0 24,200 15.33 Col 3 B1 63.5 6,920 44.98 Col 4 B1 66.3 9,657 21.08 Col 6 B1 38.7 31,060 7.24 Col 8 B1 37.4 15,536 28.51 Col 9 B1 43.4 19,548 22.0 Col 11B1 18.0 21,053 17.93 Col 12 B1 24.5 6,000 69.90 Col 13 B1 35.6 8,519 46.92 Col 14 B1 40.7 10,721 37.05 Col 15 B1 34.3 19,586 20.30 Col 17 B2 48.9 10,178 35.51 Col 25 B2 15.9 11,030 37.54 Col 26 B2 6.8 39,142 13.86 Col 2 B3 31.9 21,933 21.73 Col 11 B3 33.0 11,667 36.35 Col 13 B3 3.0 29,066 19.73 Col 14 B3 6.2 39,656 14.99 Col 15 B3 4.7 67,290 8.29 Col 1 B4 37.7 22,246 24.19
55
For example, Rajan et al. (1974) have shown that the type of surface hydroxyls that
react with phosphate change from aquo-groups at low coverage, to hydroxo-groups at
medium coverage, to ol-groups at high coverage. Since these require increasing amounts
of energy to displace, the heat of reaction with arsenate most likely would decrease with
surface coverage. The ∆H values tabulated in Table 5-2 are integral heats and, therefore,
are averages over all reactions that occurred throughout the course of arsenate sorption.
Because of this fact, it is difficult to explain unequivocally, based on these data alone, what caused the decline in the ∆H values as the Al:As ratio decreased.
A decline in the heat of adsorption would mean that the AHO surface is not energetically uniform and/or that the sorption mechanism changes with surface coverage.
The energies of adsorption could vary from site to site on the AHO surface or adsorbed arsenate could affect the energy of sorption of arsenate on adjacent sites, resulting in a heterogeneous distribution of energy sites. Alternatively, the decline in heats of adsorption with surface coverage could be due to a change in the reaction mechanism.
Similar findings were obtained by Miltenburg and Golterman (1998) who found that the heat of the adsorption of orthophosphate onto iron hydroxide decreased with decreasing ratios of Fe/P (Table 5-3). Although obtained with a different metal oxide/anion system, the authors concluded that phosphate adsorption does not take place on energetically equal sites. Additionally, the enthalpies of reactions and Fe:P mole ratios presented were similar to our findings, suggesting a common underlying mechanism in the sorption of comparable oxyanions on iron and aluminum hydroxides.
Similarly, to explain data obtained on adsorption of arsenate and arsenite on amorphous iron hydroxide, Pierce and Moore (1981) suggested a heterogeneous site
56
model for the oxide surfaces which intrinsically have different types of sites with
different affinities for adsorbate ions. This is consistent with the findings of Rajan et al.
(1974).
Table 5-3. Heat production after sequential additions of orthophosphate solution to a Fe(OOH) suspension. Adapted from Miltenburg and Golterman, 1998. mg P Joules kJ per Fe:P pH added produced Mol mole ratio 0.0 5.9 2.5 1.63 40.5 50 6.1 5 1.16 14.4 25 7.5 1.16 9.6 16.5 6.6 10 0.995 6.2 12 12.5 0.485 2.4 10 6.5 15.0 0.202 0.8 8.2 6.3 17.5 0.151 0.5 7 6.3
Changes In pH With Arsenate Sorption
Rationale
Adsorption of arsenate on AHO is believed to occur predominantly by ligand
+1/2 -1/2 exchange with surface structural S-OH2 and/or S-OH . Possible reactions of arsenate adsorption on similar surfaces abounds in the literature (Goldberg and Johnston,
2001; Jain et al., 1999; Suarez et al., 1998). Based on the proposed effect on pH, the equations can be grouped into three categories; those that propose (i) an increase in pH or an OH- release (equation 1), (i) a decrease in pH or an H+ release (equation 2) or (iii) no change in pH (equation 3). For example, for arsenate forming bidentate-binuclear complexes, the equations are as follows:
- 0 - 2SOH + H2AsO4 ↔ S2HAsO4 + H2O + OH Equation 1
+ - 0 + 2SOH2 + H2AsO4 ↔ S2HAsO4 + 2H2O + H Equation 2
- - 2SOH + H2AsO4 ↔ S2AsO4 + 2H2O Equation 3
57
The purpose of this experiment was to investigate the change in pH, if any, resulting from
the sorption of arsenate on our AHOs. Conducting a correct quantitative analysis would
have required a CO2-free environment as well as a larger sample size than we could have
used. Therefore, the results of this experiment are qualitative in nature.
Procedure
It was not possible to accurately detect a pH change at the outlet side of the
calorimeter on either flowing or collected solution. Alternatively, this study was
conducted in batch as follows. To an Eppendorf 250 µl capped centrifuge tube 200 µl of
the NaCl/Na2HAsO4 solution that was used in flow experiments was added. After an
initial pH determination, a small amount of AHO was added and the tube capped and
shaken. The pH was re-measured after 5 min and again after 2 days of equilibration. All
measurements were replicated three times. A microcombination flexible pH electrode
with a 0.5 mm tip was used.
Results
The initial response to the addition of an AHO sample to a NaCl/Na2HAsO4
solution was an increase in pH for all batches except B4 (Table 5-4), for which the pH
decreased. The pH values obtained at the end of a 2-day period during which the samples
sat in the arsenate solution showed a significant decrease relative to the pH of the NaCl/
Na2HAsO4 solution.
Although qualitative, our results suggest the occurrence of at least two reactions of
arsenate with the AHO surface: an initial reaction that releases OH- from the surface and
a slower reaction that releases H+. It can also be inferred that there is an order to the arsenate reactions: the OH- is released at an earlier time than the H+. Based on the H+/OH-
release stoichiometry of arsenate adsorption on ferrihydrite as well as on the surface
58
charge changes, Jain et al. (1999) suggested a changing mechanism of adsorption with increasing surface coverage.
Table 5-4. Effect of arsenate sorption on pH of solution. AHO weight pH values mg initial after 5 mn after 2 days Batch 1 17.1 (1.27)a 5.93 (0.14) 7.05 (0.07) --b Batch 2 6.60 (0.36) 5.37 (0.10) 5.98 (0.24) 4.80 (0.28) Batch 3 4.27 (0.31) 5.48 (0.10) 6.23 (0.41) 4.81 (0.29) Batch 4 1.77 (0.55) 5.03 (0.07) 4.35 (0.42) 4.30 (0.04) a Number in parenthesis are standard deviations of the means. b Not measured at the time of the experiment.
CHAPTER 6 EFFECT OF ARSENATE SORPTION ON AMORPHOUS ALUMINUM HYDROXIDES SURFACES
Flow adsorption calorimetry will not only give a direct measure of the heat effect
but will also identify and quantify changes that occur in the properties of a surface as a result of the specific treatments or aging effects (Steinberg, 1981). Indeed, the changes in the heats and extent of ion exchange before and after arsenate treatment on a sample of
AHO can be used as a probe of the surface and the mechanisms by which arsenate interacted with it. In this chapter, we report first on the effect of arsenate sorption on ion exchange. Second, we investigate the resulting changes in the ZPC of the AHO.
Effect of Arsenate Sorption on Ion Exchange
Rationale
The evaluation of changes in oxide surface properties has been especially useful in adsorption studies of phosphate and sulfate in the past as well as arsenate and arsenite more recently (Goldberg and Johnston, 2001; Jain et al., 1999). Properties mostly investigated were H+/OH- release stoichiometry, changes in surface charge, Infra Red
absorption spectra as well as electrophoretic mobility.
This section describes experiments that were designed to study the changes in heat
of reaction and magnitude of ion exchange resulting from arsenate sorption. Collected in
a “flow” system, the data captured effects of multiple arsenic treatments to the same
AHO sample, offering thus a unique perspective into the arsenate sorption system that
cannot readily be obtained with other analytical techniques.
59 60
Procedure
Columns containing approximately 15 mg of AHO were packed inside the
calorimeter. Heats of ion exchange (Cl/NO3 and K/Na) were measured as detailed in
Chapter 4. Samples were then exposed to cycles of 1 mM As (V) of varying duration
according to the procedure described in Chapter 5. The heats of ion exchange were re-
measured after each arsenate cycle. When exposure to all arsenate cycles had ended,
columns were recovered and digested. Sorbed arsenate was determined by ion
chromatography and/or graphite furnace. AEC and CEC were quantified as described in
Chapter 2.
Results
Anion exchange
Arsenate sorption on all four batches of AHO resulted in loss of AEC (Fig 6-1).
There was no corresponding increase in CEC. Arsenate sorption on AHO is believed to
½ + occur predominantly by ligand exchange with the dangling valence unsatisfied S–OH2
groups. These same groups are the source for anion exchange as well. The areas of the
peaks in Fig. 6-1 and the percent reduction after arsenate sorption are presented in Table
6-1. A reduction in AEC was observed for all columns exposed to arsenate (see also
Appendix).
Table 6-1. Peak areas and reductions associated with Cl/NO3 exchange peaks before and after arsenate exposure of samples shown in figure 6.1 Batch 1 Batch 2 Batch 3 Batch 4 ------peak areas in V ml a------Before As 55.30 89.60 44.50 19.9 After As 34.40 69.30 32.10 5.94 % reduction 37.80 22.70 38.30 70.2 a average values of NO3 exotherm and Cl endotherm peak areas
61
Figure 6-1. Cl/NO3 heats of exchange before and after exposure to arsenate for batch 1 (B1), batch 2 (B2), batch 3 (B3) and batch 4 (B4) samples.
Exposure to additional arsenate cycles continued to reduce the AEC as evidenced
by the reductions in peak areas for Cl/NO3 exchange (data tabulated in Appendix).
The loss in heats of exchange was proportional to the decrease in AEC of the
sample. The AEC remaining after arsenate sorption was measured on the column digest
at the end of each calorimetric run. The decrease in AEC matched fairly well the loss in
heats of anion exchange as shown in Fig 6-2, indicating that the heat of Cl/NO3 exchange
remained constant as the AEC decreased due to arsenate loading. Thus, arsenate sorption
had no effect on the heat of anion exchange, and in turn, no effect on the energetics of ion
exchange.
62
Figure 6-2. Relationship between heats and magnitude of Cl/NO3 exchange of batch 1 samples throughout various arsenate treatments.
The correspondence between the heat and magnitude of Cl/NO3 exchange
suggests that arsenate adsorption does not substantially alter the enthalpy of anion
exchange reactions where the ions involved interact with the surface via electrostatic interactions. While the arsenate reacts with the same surface functional groups that are responsible for Cl/NO3 exchange, the nature of the electrostatic interaction between the surface and either NO3 or Cl remains essentially the same. This is in contrast to the effect
of pH where increasing the pH reduced AEC but also decreased the heat of exchange between Cl and NO3. Further investigations would be required to fully understand the causes for this difference in behavior on amorphous AHO surfaces. Although some insight into this problem could be attained via calorimetric experiments, such was beyond
the scope of this study.
The loss of available AEC as a function of the amount of arsenate adsorbed is
depicted in Fig 6-3. For each mole of arsenate adsorbed, between one and two moles of
AEC were lost, with an average of 1.61 moles (standard deviation ± 0.68).
63
Figure 6-3. Loss of available anion exchange capacity of B1 samples (in µmoles of charge) as a function of amount of arsenate sorbed in µmoles.
The surface charge of the AHO is determined by protonation and deprotonation of the surface hydroxyls. Since it has already been found that the 2 pK model describes the
AHO surfaces better, the surface charge in this model will be as follows:
+ + + - H 0 -H > S--OH > S--OH > S--O - 2 + + +H +H
It has been postulated that arsenate forms three surface species with iron and
aluminum oxides: a monodentate and two bidentate (mononuclear and binuclear) species
(Waychunas et al., 1993), with the mode of bonding related to the extent of surface coverage (Fendorf et al., 1997). Using the 2 pK model, it is easy to account for the 1:1 relation between moles of arsenate sorbed and moles of AEC lost with either monodentate or bidentate arsenate complexes. For example, reactions with aquo-groups could be written as follows:
+ - 0 —SOH2 + H2AsO4 → —S--OH2AsO3 ] + H2O monodentate
+ - 1+ —(SOH2 )2 + H2AsO4 → —(S--OAsOH )2] + 2 H2O bidentate
64
These equations obviously depict only a part of the whole picture as will be demonstrated
later. Yet none of the reaction mechanisms proposed in the literature so far, with the 2pK
model, were able to account for the stoichiometry of charge lost observed in this
experiment, i.e., a 2:1 mole ratio between AEC lost and arsenate sorbed, while explaining the other observations, i.e., the OH- release and the lack of conferred negative charge
discussed later in this chapter. Furthermore, Figs 6-3 and 6-4 demonstrate that while an
increase in arsenate loading reduces the AEC, some anion exchange sites remain
available even after very high amounts of arsenate loading.
In modeling bidentate adsorption on pristine surfaces, Benjamin (2002) argued
that while the number of available sites decreases as a direct result of bidentate sorption,
sites immediately adjacent to the adsorbed molecule become unavailable to further
bidentate sorption. Depending on their position on the surface, these excluded sites could
be available for formation of monodentate surface species. The number of these excluded
sites is a function of surface coverage and pattern of surface occupation i.e. proportion of
monodentate to bidentate species. This model could explain the observation that, for a
specified duration of arsenate exposure, AHO from the same batch can differ in the
amount of arsenate sorbed. Arsenate loading may depend on whether the occupied sites
are populated with monodentate or bidentate surface complexes, and the relative
proportion of the two types of complexes. Since the population of mono- and bidentate
complexes on the AHO surface is controlled by statistical probabilities (Benjamin, 2002),
the surface loading in any given exposure is somewhat like “throwing the dice”.
65
Figure 6-4. Loss of Cl/NO3 heat in relation to an increasing arsenate surface loading
Additionally, our data show that arsenate sorption on the AHO surface: (i) neutralizes positively charged sites, explaining the decrease in AEC, and (ii) limits further arsenate reactions with sites, leaving them available for Cl/NO3 exchange even at high surface loadings. One explanation offered by Benjamin (2002), is that sorbed molecules, in this case arsenate, physically block other exchange sites due to steric crowding.
If arsenate were blocking sites, however, the positive charge would still exist and have to be neutralized by Cl/NO3. A physical blockage would increase the distance separating the positive charge and the Cl/NO3 and should reduce the heat of exchange.
Since it has been established that sorbed arsenate does not affect the heat of exchange
(Fig 6-3.), it is difficult to argue that a similar physical blockage took place.
A possible mechanism that would account for losses > 1:1 is through polydentate, namely tridentate complexes, which are attached to the surface by three bonds. The equation for the reaction would take the following form:
+ 3+ - 1+ + —(SOH2 )3 ] + H2AsO4 → —S3HAsO4] + H + 3 H2O tridentate
66
This reaction will account for a 2:1 stoichiometry between AEC lost and arsenate sorbed
and could explain the H+ release observed in earlier experiments. Tridentate complexes
have been very rarely invoked in oxyanion sorption studies, which tended to concentrate on monodentate and bidentate (mononuclear and binuclear) complexes. Brown et al.
(1999) mentions the possibility of oxyanions, such as selenate and phosphate, forming
tridentate complexes with mineral surfaces. Yet other authors, (Tejedor-Tejedor and
Anderson, 1990) discarded the possibility of tridentate complexes of phosphate on goethite because of incompatibility with surface area properties. The feasibility of
tridentate complexes on AHO surfaces will depend on distances between neighboring Al
ions and As-O bonds. Yet without contradictory evidence, it seems at least theoretically
logical to expect such polydentate complexes on a three dimensional highly porous
surface.
Cation exchange
Arsenate sorption on all samples studied, with the exception of two, did not result
in any detectable increase in CEC at pH 5.75 (Fig 6-5).
This phenomenon suggests that arsenate does not confer any negative charge to the
surface. Considering that the detection limit of our calorimeters is < 1 mJ and estimating
the heats of K/Na exchange to be around 2.0 kJ mol-1 (conservative assumption based on
actual values measured, see Appendix ), the detection limit for CEC is < 0.5 µmol (+), a value less than half the number of µmoles of sorbed arsenate. This is clearly shown in the appendix for all arsenated columns, namely col 1 B1, col 17 B2 and col 10 B3 presented in Fig 6-5.
67
Figure 6-5. K/Na heats of exchange at pH 5.75 after exposure to arsenate for batch 1 (B1), batch 2 (B2), and batch 3 (B3) samples
Table 6-2 presents results for the two samples (cols 14B3 and 15B3) that showed an increase in CEC after exposure to arsenate. Also included for comparison are data for two samples that did not exhibit any CEC despite similar experimental conditions. For reasons unknown, arsenate exposure for columns 14B3 and 15B3 resulted in higher arsenate loadings than usual; actually, these two values were the highest arsenate loadings that were ever measured in our calorimeter experiments. It is therefore suggested that the appearance of CEC in these two samples was related to the higher arsenate loadings.
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Table 6-2. Comparisons between samples that showed an increase in Cation Exchange Capacity (CEC) after As exposure and samples that did not. As cycle As As sorbed Al:As As CEC CEC Column minutes µg/g µmol ratio kJ/mol cmolc/kg kJ/mol 1 B4 20 22,200 7.80 37.70 24.20 0 --a 12B1 10 6,000 1.20 69.9 24.5 0 -- 14B3 10 39,700 8.50 15.0 6.19 1.97 2.36 15B3 15 67,300 13.90 8.30 4.69 3.98 3.87 a not determined as no CEC is measured
Effect of Arsenate Sorption on ZPC
Rationale
It is fairly well established in soil chemistry textbooks and literature that arsenate,
and other oxyanions such as phosphate, carbonate, and chromate, form inner-sphere
complexes with the surface of oxides through what is termed specific adsorption.
Extended evidence has been presented in support of this arsenate sorption mechanism and
included (i) little ionic strength dependence as a function of solution pH (Goldberg and
Johnston, 2001), (ii) Fourier Transform Infrared (FTIR)spectroscopy and electrophoretic mobility (EM) data (Suarez et al., 1998), and (iii) Extended X-ray Absorption Fine
Structure (EXAFS) spectra (Waychunas et al., 1993).
Specific adsorption, or ligand exchange, differs from anion exchange (characterized by outer-sphere complexes) in that it may change the charge on the oxide-particle. The change in surface charge results in shifts in the ZPC and reversal of EM, observations considered conclusive evidence for specific adsorption (Hunter, 1981). Nonetheless,
Suarez et al. (1998) and Goldberg and Johnston (2001) caution that a shift in ZPC need not occur since not all inner-sphere complexes are accompanied by a change in surface
charge.
69
In view of our results demonstrating that in the majority of the cases, arsenate does not confer a charge on the surface, it was of interest to investigate whether arsenate sorption had an effect on the ZPC. Comparisons of heats of anion and cation exchange for clean and arsenated AHO at different pHs can show not only the effects of arsenate on
ZPC but also information about underlying mechanisms for that change.
Procedure
Flow studies
Columns containing approximately 15 mg of AHO were packed inside the calorimeter. Heats of ion exchange (Cl/NO3 and K/Na) were measured and samples
exposed to a different number and/or duration of 1 mM As (V) cycles. The heats of ion
exchange were re-measured after each arsenate cycle. The solutions were then changed
according to the procedure described for the ZPC determination, i.e., increasing pH from
5.75 to pHs 8.0 and 10.5 with heats of ion exchange determined after each pH change.
When all necessary data were collected, columns were recovered and digested. Sorbed
arsenate was determined by ion chromatography and AEC and CEC measured as
described in Chapter 2.
Batch studies
Some B3 AHO samples were arsenated in a batch process prior to packing in the
calorimeter for PZC determination as follows. To a 100 mg sample weighed in a capped
centrifuge tube were added 40 ml of 1mM As(V) solution (50 mM NaCl in which 2
-1 -1 mmol L of the NaCl had been replaced with 1 mmol L Na2HAsO4.7H20, i.e., keeping
the total Na concentration at 50 mM and As (V) at 1mM). The NaCl/Na2HAsO4 solution had a pH of 5.75.
70
Samples were shaken manually for a couple of minutes and left to stand for another
15 minutes, after which tubes were centrifuged and supernatant solution decanted. This
procedure was repeated in total three times and followed by 3 additional washes with 50
mM NaCl, i.e., without the arsenate, to wash any entrained NaCl/Na2HAsO4 solution .
Samples were then dried in an oven at 70 °C, crushed, sieved through a 150 µm mesh and
stored in a glass scintillation vial. Samples of about 15 mg were then packed inside the
calorimeter and taken through the procedure of the ZPC determination. At the end of the
experiments, columns were recovered and digested. Sorbed arsenate was determined by
ion chromatography, and AEC and CEC were measured as described in Chapter 2.
Samples of batch arsenated B3 that did undergo neither calorimetric nor pH treatments were also digested and analyzed to determine their arsenate content. This arsenate concentration was compared with that obtained from batch arsenated B3 samples recovered at the end of the calorimetry experiment.
Results
PZC shifts: Flow versus batch
There was no difference in arsenate sorbed between batch arsenated B3 samples that were recovered at the end of the PZC experiments and those not subjected to pH treatments, indicating that there was no loss of arsenate from the AHO as a result of the pH 8.0 and 10.5 treatments in the calorimeter experiments. Additionally, there was no detectable arsenate measured in solutions collected at the outlet side of the calorimeter during the pH 10.5 exposure. Higher pH treatments thus did not cause any arsenate loss
from the AHO. This finding is consistent with findings described in Chapter 3 on AHO
dissolution at higher pHs.
71
For clean B3 AHO that was arsenated in the flow system, arsenate sorption resulted in relatively small shifts in ZPC, from 9.5 to 8.5 on a B2 sample and from 9.4 to 9.0 on a
B3 sample (Fig 6-6).
Figure 6-6. Calorimetric determination of the ZPC of a batch 2 (B2) and batch 3 (B3) AHO sample treated with As(V) in a flow system. The diamonds indicate the ZPC.
Our results showed lower shifts than most findings reported in the literature.
Goldberg and Johnston (2001) using electrophoretic mobility data, demonstrated shifts of
up to 4 pH units in the ZPC of AHO samples in a 1.0 mM As(V) solution suspended in a
0.01 M NaCl. Similarly, Jain et al. (1999) reported a reduction of ZPC by as much as 2.4
pH units on batch arsenated ferrihydrite samples.
The ZPC obtained from batch arsenated AHO in this study differed markedly from
that where clean B3 AHO was arsenated in the flow calorimeter, and were more in
agreement with shifts in ZPC that have been reported by other authors using batch
arsenated oxides. For B3 samples exposed to arsenate in a batch system prior to the
72
calorimetry, shifts in ZPC from an initial value of 9.4 to 8.0 and 5.5 were measured (Fig
6-7). Despite this variation, these shifts were considerably greater than those obtained on
B3 samples that were arsenated in flow.
Figure 6-7. Calorimetric determination of the ZPC of batch 3 (B3) AHO samples treated with As(V) in a batch system. The diamonds indicate the ZPC.
In an attempt to reconcile the two contrasting results, we compared two samples,
one from the batch arsenated study and one from the flow arsenated study (Table 6-4).
Table 6-4. Comparisons in zero points of charge shifts and other data of batch 3 samples arsenated in flow and in batch. ZPC shift As sorbed K/Na peak areas in V/ml after As Final CEC Column in pH units µg g-1 pH 5.75 pH 8.0 pH 10.5 cmol(-)/kg 11B3-flow 0.4 11,700 0 0.65 8.65 3.67 A2B3-batch 3.9 25,800 1.80 3.06 14.3 12.90 a averages values of K exotherms and Na endotherms peak areas
As shown in the table, several differences occured between columns A2B3 and
11B3. Column A2B3 (i) sorbed more arsenate, 25,842 µg g−1 versus 11,667 µg g−1, (ii) had a measurable heat of cation exchange at pH 5.75 and bigger peak areas at pHs 8.0 and 10.5, and (iii) displayed almost four times as much final CEC. The discrepancy in the
73
magnitude of ZPC shifts between flow- and batch- arsenated samples can be explained on
the grounds of differences in arsenate coverage and its effect of surface charge. In fact,
Goldberg and Johnston (2001) mentioned that the ZPC of their amorphous Al oxide was
not shifted in the presence of a lower arsenate concentration, 0.01 mM As(V) as opposed to a 4 pH unit shift with a higher arsenate concentration, 1.0 mM As (V) solution.
Although they did not report the amount of arsenate actually adsorbed on the samples, they suggested that at the lower arsenate concentration, arsenate forms inner-sphere surface complexes that do not create negative surface charge. However, this explanation accounts for only part of the observations as we will discuss in the next section and final chapter. The difference in behavior between batch- and flow- arsenated AHO demonstrates the unpredictability of these amorphous surfaces, and questions the relevance of batch experiments to field cases where solutions are percolating through the soil at high solid: solution ratios, similar to conditions obtained in the flow calorimeters.
ZPC shifts: CEC effects
In the following section, we examine closely the heats of ion exchange obtained while determining the ZPC of both arsenated and clean samples and the information calorimetry reveals about the mechanisms involved. Data for clean samples and flow-arsenated B2 and B3 AHO are presented in Tables 6-5 and 6-6.respectively.
As observed in previous experiments, arsenate sorption on samples 17B2 and 25
B2 reduced AEC and did not result in any detectable CEC (Table 6-5). The AEC dropped further as the solution pH was raised to 8.0 and eventually disappeared at pH 10.5; CEC was observed at pH 10.5. The ZPC dropped slightly from 9.0 on clean B2 to around 8.2 and 8.5 on arsenated AHO.
74
Table 6-5. Comparisons of heats of ion exchange and other data from clean and arsenated batch 2 samples at different pHs. Column 11B2 17B2 25B2 Description ZPC ZPC after ZPC after As exposure As exposure Al content in % 12.6 14.7 14.9 As sorbed in µmoles 0 1.99 2.19 a Cl/NO3 peak in V ml initial 64.90 94.20 57.3 after As --c 57.0 29.8 pH 8.0 11.80 5.90 2.29 pH 10.5 0 0 0
K/Na peakb in V ml initial 0 0 0 after As --c 0 0 pH 8.0 0 0 0 pH 10.5 1.65 2.85 4.32 ZPC 9.0 8.5 8.2 Final CEC in cmol(-)/kg 1.15 11.20 6.74 a averages values of NO3 exotherms and Cl endotherms peak areas b averages values of K exotherms and Na endotherms peak areas c does not apply
Similar findings were obtained for B3 clean and arsenated samples (Table 6-6) with the only difference being a small but detectable heat of K/Na exchange at pH 8.0. The
ZPC shifted from pH 9.5 on the clean sample to an average of 8.8 on ones exposed to
arsenate. Additionally, B3 samples exposed to arsenate had a higher CEC at pH 10.5 than
did B2 samples.
These results provide valuable insights into the mechanisms whereby As sorption shifts the ZPC. Data collected on these samples show explicitly that arsenate sorption did
not confer a negative charge to the surface at pH 5.75 yet it caused a measurable shift in
the ZPC. This shift was a result of a greater drop in AEC and a greater increase in CEC
for arsenate treated AHO as pH was raised.
75
Table 6-6. Comparisons of heats of ion exchange and other data from clean and arsenated batch 3 samples at different pHs. Column 13B3 9B3 10B3 11B3 Description ZPC ZPC after ZPC after ZPC after As exposure As exposure As exposure Al content in % 14.4 15.5 15.6 15 As sorbed in µmoles --c n.a.d n.a.d n.a.d a Cl/NO3 peak in V ml initial 56.40 65.30 62.30 64.20 after As --c 32.20 42.70 38.20 pH 8.0 19.0 7.38 4.86 8.8 pH 10.5 0 0 0 0
K/Na peakb in V ml initial 0 0 0 0 after As --c 0 0 0 pH 8.0 0 0.93 0.23 0.65 pH 10.5 7.62 6.30 3.05 8.65 ZPC 9.5 8.8 8.6 9 Final CEC in cmol(-)/kg 2.47 6.69 8.43 3.67 a average values of NO3 exotherms and Cl endotherms peak areas b average values of K exotherms and Na endotherms peak areas c not relevant d data not available
Despite a reduction in the number of reactive SOH groups caused by arsenate
sorption, treated samples had a higher CEC at pH 10.5 than did the clean samples. When
compared to clean AHO, the arsenate treated AHO generated more CEC at pH 10.5 and
did it with fewer SOH groups. The clean AHO should have had a higher number of
reactive SOH groups available for deprotonation as pH is raised. This is in sharp contrast
with the generally accepted view that arseante sorption shifts the ZPC by putting negative
charge on the surface. Other mechanisms can be suggested to explain the reported
behavior of the AHO surfaces upon arsenate sorption and pH change.
As mentioned previously in Chapter 4 (see section: Charging of AHO Surfaces),
the AHO surfaces owe their charge to amphoteric surface hydroxyls that undergo
protonation and deprotonation reactions. Depending on the charging model, one or two
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acidity constants (K1 and K2) are thus defined. For the sake of the argument and for the duration of this section, we will refer to only one K value that corresponds to a pH range
of 5.75 to 8.0, regardless of it being K1 of the two pK model or just K in the one pK
model.
Calorimetrically, the K value is manifested through the magnitude of the heats of
Cl/NO3 exchange, i.e., the size of the peak areas. In that sense, a reduction in the size of
the peak area upon increase in pH translates into a decrease in the number of protonated
surface hydroxyls, with this number being a function of K. In other words, if the pK = 6,
+ 0 then at pH 6, half of the SOH2 groups will be deprotonated to SOH , and this would be
detected calorimetrically as a loss of half of the AEC at pH 6.0 compared to, say, the
AEC at a pH of around 4, where essentially 100% of the SOH are protonated. Thus, if at
a given pH, this fractional change in AEC changes with arsenate sorbed, this can be
interpreted as a change in pK. In fact, throughout our experiments, the decrease in peak
areas observed as the solution pH is raised from 5.75 to 7.25 or 8.0 was consistently
uniform within a given batch of clean AHO (Table 6-7). Therefore, it was appropriate to
consider the change in percent reduction in peak area (and AEC) as a result of arsenate
adsorption an indication of a change in K.
An examination of the percent reductions in AEC on the clean and arsenated samples should indicate whether arsenate sorption is affecting the pK. In Table 6-8 are
shown the reductions in Cl/NO3 peak areas for the B2 and B3 sample data in Tables 6-5
and 6-6. For batches 2 and 3, arsenated samples had a significantly higher reduction in
the Cl/NO3 exchange peak areas upon exposure to pH 8.0 than the clean samples. In
+ terms of pK, that means the —SOH2 groups have become more acidic, losing a proton to
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the solution. The shift in the pK of the surface is one possible mechanism by which the
ZPC of the AHO shifts in cases where no negative charge is conferred to the surface by
arsenate sorption.
Table 6-7. Reductions in Cl/NO3 peak areas with increase in solution pH from 5.75 to 7.25 and 8.0 for clean batches 2 and 3 samples. a Cl/NO3 peak areas in V ml Reduction Columns pH 5.75 pH 7.25 pH 8.0 in % Batch 3 3B3 56.40 19.0 66.30 16B3 48.40 15.80 66.70 5B3 59.60 22.60 62.20 7B3 51.90 16.2 68.80 Batch 2 11B2 67 11.80 81.9 20B2 102 12 88.20 22B2 102 15.70 84.60 a averages values of NO3 exotherms and Cl endotherms peak areas
Table 6-8. Reductions in Cl/NO3 peak areas with increase in solution pH from 5.75 to 8.0 for clean and arsenated batches 2 and 3 samples. Reduction in Cl/NO3 peak area in % columns clean arsenated Batch 3 3B3 66.30 9B3 77.0 10B3 88.60 11B3 77.0 Batch 2 11B2 81.90 17B2 89.70 25B2 92.30
A similar pK shift could be invoked to explain how surfaces that have lost reactive
SOH groups to arsenate sorption exhibit a higher CEC at pH 10.5 than samples that have
not been exposed to arsenate. The remaining reactive groups, which deprotonate into the
neutral species SOH0 as the pH is raised, could have also become more acidic, losing a
78
proton quicker and hence forming a higher number of S--O- sites as compared to clean
surfaces.
While it would have been possible to corroborate this mechanism calorimetrically,
the high ZPC of the AHO was not conducive for such an investigation. CEC was not
always detectable at pH 8.0, and therefore changes in surface negative charge with pH for
clean and arsenated AHO was generally not available. Nevertheless arsenated B3 samples
did exhibit heats of K/Na exchange at pH 8.0 while clean B3 samples did not;
circumstantial evidence in favor of a shift in pK at pH values above the ZPC.
Another possible mechanism could however be put forward to explain the higher
CEC obtained on arsenated samples at pH 10.5. The pK values for arsenic acid in
solution are 2.3, 6.8 and 11.6, thus the dominant aqueous species at pH 5.75 is the
1- monovalent H2AsO4 . If we assume for the sake of simplicity bidentate binuclear
complex formation, the arsenate on the surface will have the following configuration:
Al O OH As Al O OH
As the solution pH is raised to 8.0 and then to 10.5, the adsorbed arsenate could itself undergo a deprotonation reaction, creating sites with negative charge. It is very important to remember, as noted by Anderson and Malotky (1979) that acidity constants for oxyanions adsorbed on the surface may not equal their solution counterparts, due to the chemical influence of the central metal ion. In that sense, the pK values for arsenic acid need not apply exactly to the adsorbed species, but probably indicates an expected
range. The dissociation of the adsorbed arsenate at the surface could account for the extra
CEC observed. Anderson and Malotky (1979) have suggested a similar mechanism
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whereby specifically adsorbed ions provide new surface acidity groups that undergo
deprotonations reactions. At this point, and without further investigations, it is not
possible to support either mechanism conclusively. To do that, it is necessary to partition
the negative charge between reactive SO- groups and the adsorbed arsenate. This was not
possible because some of the AEC that was lost at pH 5.75 was a result of arsenate
+ sorption, and some was a result of the shift in pK. Therefore the number of SOH2 groups on the surface after arsenate sorption is unknown.
Regardless of whether the adsorbed arsenate deprotonates, a change in the pK of the reactive SOH+ sites not only accounts for (i) a shift in the ZPC after arsenate sorption
in the absence of negative charge, and (ii) a higher CEC on arsenated than on clean
samples at pH 10.5, but also explains a stoichiometry > 1:1 between AEC lost and
+ arsenate sorbed. A change in the pK of the SOH2 sites must occur as an immediate result
of arsenate sorption and most likely is a function of the amount of arsenate adsorbed. A
possible mechanism for shifting the pK is that the more electronegative arsenate attracts
electrons away from the surface. It is then possible that these sites become more reactive
towards arsenate, which is able now to neutralize a higher number of sites. The exact
mechanisms underlying this process are not known at that stage, and will require much
deeper investigations to elucidate.
Why these two mechanisms, shift in pK and deprotonation of sorbed arsenate, have
not been evoked more frequently, may in part be due to the methods often adopted in
arsenate and phosphate sorption experiments, which do not allow changes in surface
charge to be easily measured on the same sample.
CHAPTER 7 THE ARSENATE /AMORPHOUS ALUMINUM HYDROXIDES SORPTION SYSTEM: FINDINGS AND SUGGESTIONS
Over the course of this study, we have collected information on the nature of AHO surfaces and their interactions with arsenate. Any mechanism suggested for arsenate sorption should be in agreement with all experimental findings. We will, therefore, present in the first part of this chapter a brief summary of the results obtained and discuss in the second part how these observations interplay to reveal new aspects of the
AHO/arsenate sorption system.
Summary of Results
AHO Physical and Chemical Properties
The AHO synthesized and used in this study had the following chemical and physical characteristics:
• Was amorphous and hydrated in nature, with a porous, sponge like structure that was apparent only below a 1µm scale.
• Possessed a high surface area.
• Did not significantly crystallize over the time span of experiments.
• Had 13 – 20 % Al content with no occluded salt.
• Was capable of high anion exchange capacities, on the order of 94 to 131 cmol(+) -1 -1 kg of solid or 198 to 264 cmol(+) kg of pure Al(OH)3 (approximately 1 mole of positive charge per 6 moles of Al).
• Exhibited a ZPC around pH 9.5.
• Was consistent with a 2 pK model of surface charging based on the existence of the neutral species.
80 81
Ion Exchange Properties
The following characteristics of ion exchange on AHO were revealed through
calorimetric studies:
• Was rapid, reversible and reproducible over time and samples
• Had a heat of exchange between 3.6 and 5.8 kJ mol-1 for anion exchange and between 1.1 and 1.6 kJ mol -1 for cation exchange.
• Possessed unusual characteristics around the ZPC, such as the “flip flop effect” and change in energies of exchange.
Arsenate Sorption Properties
Using flow adsorption calorimetry in conjunction with other methods, we have
established the following characteristics of arsenate sorption on AHO :
• Exhibited a markedly different peak shape than anion exchange indicating kinetically different reactions.
• Was exothermic with heats of adsorption from 40 to 60 kJ mole -1 of arsenate sorbed.
• Was a much slower reaction than ion exchange.
• Heat evolution occurred with every AHO exposure to arsenate, even following cycles where the arsenate heat signature returned to baseline.
• Molar As:Al ratios were always lower than exchangeable Cl:Al ratios, indicating that the AHO maximum sorption capacity was not attained.
• The heats of adsorption decreased with increasing arsenate surface coverage (increasing As:Al mole ratios).
• Arsenate sorption resulted in OH- release followed by H+ release.
Effects of Arsenate Sorption
The effects of arsenate sorption on the AHO were as follows:
• Loss of AEC with no effect on the heat of Cl/NO3 exchange.
• No change in CEC except at very high arsenate loadings
• One mole of arsenate sorbed eliminated about 1.61 mole of anion exchange.
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• Some anion exchange sites remained available even after the highest arsenate loading.
• In flow systems the ZPC shifts by a maximum of 1 pH unit while in batch it decreased by up to 4 units.
• A shift in the pK was observed which explains the ZPC shift in the absence of an increase in surface negative charge.
• The adsorbed arsenate could undergo deprotonations above the ZPC thus creating new negative charged sites on the surface
Findings and Suggestions
On the Structure and Morphology of AHO
Since a working theory of the structure of the AHO is necessary for a systematic discussion of their reactions with arsenate, we shall address first issues of structure and morphology. Collected evidence on their physical and chemical properties have indicated the porous spongy-like structure of these AHO. Gels of hydrous oxides, which lose their elasticity and become powdery upon drying, are called rigid or non-elastic gels (Weiser,
1926). These are considered a 2 phase solid-liquid system in which there is a network or cellular arrangement of solid phase permeated by liquid, forming thus an irregular mesh or interlacing network in which liquid is entrained.
It is mandatory for any proposed AHO structure to accommodate all characteristics of the arsenate sorption reactions observed in this, as well as other, studies. Therefore the
AHO morphology should allow sorbed arsenate (i) to undergo kinetically different reactions with the slower reaction accounting for a slower heat evolution at every exposure, (ii) penetrate inside an AHO structure just after one day of exposure
(unpublished data by K.C.Makris); and (iii) leave some anion exchange sites available for exchange at the surface.
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One possible spatial geometry that would reconcile all of the above observations is
one in which strands of AHO polymer are twisted and folded into an open structure. In
other words, there is no external surface per se, but rather an interconnected network of
pores and conduits of various sizes forming a cotton-like structure, much like that of a
protein, with reactive functional sites dispersed throughout. It is a loose, highly hydrated
structure, which is permeable to hydrated ions. In fact, the SEM-EDX elemental map
shows a distribution of Cl all over the particle. The mole ratios of 1 anion exchange site
per 6 Al ions provide further evidence for such a structure.
On AHO Surface Chemistry
The chemistry of the AHO is such that the terminal surfacial sites are occupied by
aquo (Al-H2O) and hydroxo (Al-OH) groups. Internally, the aluminum atoms are linked
by ol and oxo groups
OH O olAl Al and oxo Al Al OH O
While it is possible for researchers to verify with a high level of confidence these structures on crystalline surfaces with exact location, sites densities, and coordination geometry, it is not the case for amorphous surfaces such as the AHO used in this study.
This becomes problematic when assigning charges on surface reactive OH- as well as
depicting the coordination environment of the metal ion, in our case Al; in fact we have
outlined earlier some of the discrepancies that exist regarding surface charging models ( 1
pK versus 2 pK). Existing literature differs even in its representation of 1 positive charge,
as shown below:
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1+ 1+ OH2 Al OH2 Al or OH2 Al OH
Schematic illustrations of the structure of the arsenate on AHO surfaces, such as monodentate or bidentate species, have often been based on this limited two dimensional surface figure that offer no information as to the neighboring sites, charge distribution or coordination environment. The difficulty in obtaining such information for amorphous surfaces casts some doubt on the relevance of modeling arsenate sorption by AHO. For
AHO, findings should be interpreted by resolution of the experimental data rather than by imposing a particular model on the data. However, without further detailed investigations that are beyond the scope of this study, we can only use similar terminology, which comes with the same limitations, to describe arsenate reactions. Nevertheless, having said that, any model describing the behavior of these surfaces should allow for the existence of a neutral specie as confirmed experimentally, and it is therefore more appropriate to use the 2pK charging model for now.
On Arsenate Sorption
Sorption of arsenate on aluminum hydroxide can be interpreted in terms of physical and chemical processes. Findings in this study were best reconciled on the grounds of proposed work by Rajan et al. (1974) and Rajan (1975; 1976; 1978; 1979) on divalent phosphate and sulfate adsorption by allophanic clays and hydrous aluminum oxides, in which he identifies the different reactive sites from proton consumption and changes in net surface charge. Based largely on Rajan’s work, the following is proposed as a model for As reaction with AHO.
Arsenate sorption exhibits an initial rapid uptake phase, followed by a much
85 slower rate of uptake that may continue for hours or days. The initial, rapid uptake phase is probably a ligand exchange reaction whereby the arsenate displaces aquo groups (Al—
OH2) at the positive sites and hydroxo groups (Al—OH) at the neutral sites. The proposed equations would be:
1+ - 0 Al—OH2] + H2AsO4 ↔ Al—H2AsO4] + OH2
0 - 0 - Al—OH] + H2AsO4 ↔ Al—H2AsO4] + OH
The first reaction neutralizes the positive charge, whereas the second leaves the surface charge unaltered but releases OH-. This explains the loss of AEC with arsenate sorption without an increase in negative charge or CEC. The slower rate of heat evolution probably corresponds to arsenate gaining access to the less accessible reactive sites that are dispersed throughout the structure. This is not diffusion in the classical sense since the accessed sites are not strictly inside the structure as much as they are part of an open porous network. The Cl/NO3 exchange, on the other hand was shown to be rapid; with the anions experiencing no impendence in their movement. This could be attributed to these anions being held in an outer-sphere or electrostatic manner, which allows them to move more freely inside the hydrated AHO. Arsenate, however, forms inner sphere complexes with the “surface” and should move along the surface much more slowly.
Nevertheless, results demonstrated (Chapter 5) that arsenate undergoes a spatial rearrangement that regenerates reactive sites, explaining not only the peaks in successive cycles, but also the recurrence of the initial rapid phase in heat evolution at each exposure to arsenate. Arsenate does not saturate all possible sites as demonstrated by the much higher Al:As than Al:Cl molar ratios obtained, and the fact that some anion exchange sites remain available even after the highest arsenate loading. The return to baseline
86 observed despite the arsenate reactions with the AHO being unfinished could indicate that the reactions involving the redistribution of arsenate(i.e., the regeneration of sorption sites) are either entropy driven or occur at a very slow rate whereby the heat generated is dissipated and lost in the baseline noise.
These explanations constitute only a part of the complete picture. Some of the results remain unaccounted for, namely (i) in some cases, cation exchange capacity has been observed after arsenate sorption indicating a negative charge was conferred to the
AHO “surface”, and (ii) the decreasing heats of adsorption with increasing arsenate surface coverage and As:Al mole ratios on some samples. These observations indicate that at a higher fractional saturation, there is a change in the arsenate sorption mechanism. With increasing sorption, arsenate begins to attack the ol and oxo groups, which internally bridges the Al metal centers. The AHO polymers “break up”, opening further the structure to more arsenate sorption. Arsenate is now incorporated into the structure through structural rearrangement that produces a new arsenate/AHO solid. The proposed equation for reactions with the ol-groups is presented below and results in a net negative surface charge.
1/2- Al - Al--H 2 AsO 4 0 OH + H2 AsO 4 Al--OH1/2- Al
Contrary to the “ligand-exchange” reactions, this bond breaking mechanism may not be exothermic, but rather consumes energy as witnessed by the overall net decrease in the heats of arsenate sorption measured calorimetrically. While there must be a value of sorbed arsenate at which the attack on the ol groups begins, we can not at this point elaborate more on its nature. Most of the samples investigated apparently fell below this
87 value and did not show such a behavior. Moreover, as emphasized throughout this study the amorphous and unpredictable nature of these AHO may very well render it impossible to determine if there is an orderly sequence to these processes or if they are controlled by statistical probabilities.
Nevertheless, these results strongly favor a theory of arsenate sorption that is more consistent with a heterogeneous site model that takes into account different surface saturation values as well as a new visualization of the AHO surfaces.
CHAPTER 8 SUMMARY AND CONCLUSIONS
The main objective of this work was to demonstrate the application of flow adsorption calorimetry as a powerful technique in probing and understanding chemical surfaces, obtaining thus information not readily accessible by other methods.
For this purpose, simple and inexpensive flow calorimeters were used to measure heats of reaction. The calorimeters used a pair of thermistors to monitor changes in the temperature of a solution as it flowed through columns containing about 15 mg of solid.
A change in solution temperature generated a signal that was amplified and fed into a computer for processing. The calorimeters exhibited high sensitivity, low thermal drift, and good signal-to-noise ratio. Calorimetric methodologies for probing surfaces and their properties were applied to the adsorption of a classic oxyanion contaminant, arsenate, onto amorphous aluminum hydroxides, a highly adsorptive, synthetic mineral oftentimes used in research in lieu of their naturally occurring counterparts.
Four different batches of amorphous aluminum hydroxides (AHO) used throughout this study were synthesized in our laboratory. They were found to be amorphous to X- rays except one, hydrated and porous. The AHO possessed high surface areas (64 to 443
2 -1 -1 m g ) and elevated anion exchange capacities (94 to 131 cmol(+) kg ).
Flow adsorption calorimetry was used to characterize the surfaces of these AHOs.
Direct measurements of the calorimetric effects of ion adsorption were first measured.
-1 Heats of exchange for Cl/NO3 were 3.5 to 5.8 kJ mol and those for K/Na at 1.1 to 1.6 kJ mol -1. In addition to heats of exchange, calorimetry provided insights into the kinetics of
88 89 the exchange reaction. Novel experiments were devised to determine the Zero Point of
Charge (ZPC) calorimetrically. This new method revealed valuable information about these surfaces: Results obtained supported the 2 pK charging mechanism and demonstrated unusual characteristics around the ZPC such as a change from weak field to strong field exchanger and decrease in energies of anion exchange. Applying this technique on the same AHO sample was also key in demonstrating the reversibility of surface charge, a difficult task if traditional methods are employed. Calorimetric data helped define a spatial geometry for the AHO surface that would account for observed effects, whereby no distinction is made between internal and external surfaces per se but rather is visualized as continuum of pores in an open structure.
The reaction of arsenate with AHO was exothermic; the net heat varied on average from 40 to 60 kJ mole -1 of arsenate sorbed at low fractional coverage and decreased with increasing surface coverage, indicating a shift in the mechanism of sorption. Based on rate of heat evolution, kinetically different reactions were identified and explained based on chemical and physical properties of the arsenate/AHO system. Despite a wide variation in molar Al:As ratios following arsenate sorption, they remained much lower than the Al:Clex on clean AHO, indicating that the AHO maximum sorption capacity was not attained. Arsenate sorption caused an OH- release followed by an H+ release.
Changes that occurred in the AHO surface properties as a result of arsenate sorption were investigated. Sorption of arsenate resulted in loss in AEC. The loss in AEC was proportional to the drop of heat of exchange indicating that arsenate sorption had no effect on the heat of Cl/NO3 exchange. No change in CEC was observed for the majority of the cases studied. These observations indicated that most of the arsenate was sorbed by
90 a process that involved no increase in surface negative charge. These reactions were proposed to be ligand exchange with aquo and hydroxo groups. A slower secondary reaction was identified as arsenate gaining access to the less accessible reactive sites that are dispersed throughout the structure. With increasing concentration, sorbed arsenate reacted with the bridging ol groups, opening the AHO structure, incorporating arsenate into it and creating negative sites for cation exchange. This bond breaking mechanism occurred also at a slower rate and was likely energy consuming. Sorbed arsenate caused the ZPC to shift by as much as 1 pH unit. In the absence of negative charge conferred to the surface by arsenate sorption, a shift in the pK of reactive surface was found to explain the ZPC changes. It was also postulated that the pK shift could account for the higher
CEC obtained on arsenated samples at pH 10.5, and the stoichiometry > 1:1 between
AEC lost and arsenate sorbed.
By allowing multiple sorption cycles to be applied to the same sample and changes in surface properties to be easily measured on the same sample, flow adsorption calorimetry yielded unique information about a classical sorption system. Used in conjunction with other techniques, it has greatly improved our understanding of the surface chemical properties of AHO. The results of this investigation were fit together in a comprehensive reaction mechanism model that was found to complement the existing state of knowledge on inorganic oxyanions sorption in general. Flow adsorption calorimetry was proven to be a uniquely informative yet rapid experimental tool that can be adapted to numerous applications in surface chemistry studies.
APPENDIX RAW DATA
Table A-1. Ensemble of data collected for columns in arsenate sorption experiments Column 1B1 3B1 4B1 Description ------arsenate sorption experiments ------Number of As cycles 4 1 2 Weight in mg 12.4 14 15.3 Al content in % 13.36 9.91 7.33 Al content in µmol 61.357 51.385 41.537 a Cl/NO3 peak in V ml 0th 32.34 28.5 34.68 1st 21.17 (34.54)b 17.42 (38.88) 22.96 (33.79) 2nd 15.83 (25.22) -- c 16.11 (29.83) 3rd 5.5 (65.25) -- -- 4th ------AEC in cmol kg-1 n.a d 80.61 58.91 Lost AEC in µmol n.a 1.875 5.369 As sorbed in µg g-1 24,200 6,920 9,657 As sorbed in µmol 4.001 1.292 1.97 Al:As mole ratio 15.33 44.98 21.08 As ∆H in kJ mol-1 63.99 63.5 66.29 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
91 92
Table A-2. Ensemble of data collected for columns in arsenate sorption experiments Column 6B1 8B1 9B1 Description ------arsenate sorption experiments ------Number of As cycles 3 + 1 1 2 Weight in mg 6.6 12.6 13.3 Al content in % 8.1 15.95 15.49 Al content in µmol 19.8 74.43 76.30 a Cl/NO3 peak in V ml 0th 17.2 36.56 37.42 1st -- 27.22 (25.54) 22.11 (40.91) 2nd -- -- 16.12 (27.09) 3rd 4.66 (72.91) -- -- 4th 2.48 (46.78) -- -- AEC in cmol kg-1 n.a. 82.1 35.2 Lost AEC in µmol n.a. 1.50 7.82 As sorbed in µg g-1 31,060 15,536 19,548 As sorbed in µmol 2.733 2.61 3.467 Al:As mole ratio 7.24 28.51 22 As ∆H in kJ mol-1 38.7 37.4 43.43 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
93
Table A-3. Ensemble of data collected for columns in arsenate sorption experiments Column 11B1 12B1 13B1 Description ------arsenate sorption experiments ------Number of As cycles 1 + 0.5 0.5 1.25 Weight in mg 15.2 15 15.7 Al content in % 13.59 15.1 14.39 Al content in µmol 76.51 83.89 83.68 a Cl/NO3 peak in V ml 0th 40.51 30.73 55.4 1st 25 (38.29) 24.56 (20.08) 41.85 (24.46) 2nd 19.13 (23.48) -- -- 3rd ------4th ------AEC in cmol kg-1 56.41 79.87 69.27 Lost AEC in µmol 5.71 2.12 3.88 As sorbed in µg g-1 21,053 6,000 8,519 As sorbed in µmol 4.267 1.2 1.78 Al:As mole ratio 17.93 69.9 46.92 As ∆H in kJ mol-1 18 24.5 35.6 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
94
Table A-4. Ensemble of data collected for columns in arsenate sorption experiments Column 14B1 15B1 16B2 Description ------arsenate sorption experiments ------Number of As cycles 1 1.5 2 Weight in mg 14.9 15.4 17.7 Al content in % 14.3 14.32 17.52 Al content in µmol 78.92 81.68 114.85 a Cl/NO3 peak in V ml 0th 40.01 50.01 102.28 1st 27.12 (32.21) 29.59 (40.94) 79.29 (22.47) 2nd -- -- 62.54 (21.12) 3rd ------4th ------AEC in cmol kg-1 56.07 55.75 52.64 Lost AEC in µmol 5.65 5.89 9.13 As sorbed in µg g-1 10,721 19,586 n.a As sorbed in µmol 2.13 4.022 n.a Al:As mole ratio 40.69 20.3 n.a As ∆H in kJ mol-1 37.05 34.28 n.a a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
95
Table A-6. Ensemble of data collected for columns in arsenate sorption experiments. Column 26B2 1 B4 Description ------arsenate sorption experiments ------Number of As cycles 1 1 Weight in mg 17.7 6.9 Al content in % 19.33 19.37 Al content in µmol 108.1 49.5 a Cl/NO3 peak in V ml 0th 87.64 41.09 1st 58.32 (33.45) 20.62 (49.82) 2nd -- -- 3rd -- -- 4th -- -- AEC in cmol kg-1 63.31 13.12 Lost AEC in µmol 6.17 8.13 As sorbed in µg g-1 39,142 22,246 As sorbed in µmol 7.8 2.05 Al:As mole ratio 13.86 37.73 As ∆H in kJ mol-1 6.77 24.19 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
96
Table A-7. Ensemble of data collected for columns of clean AHO in ZPC experiments. Column 3B3 16B3 11B2 Description ZPC ZPC ZPC Weight in mg 14.4 15.9 12.6 Al content in % 14.65 15.74 12.1 Al content in µmol 78.13 92.74 56.47 a Cl/NO3 peak in V ml pH 5.75 56.35 48.44 66.95 pH 8.0 18.98 (66.32)b 15.82 11.77 (81.88) pH 10.5 0 0 0 K/Na peak a in V ml-1 pH 5.75 --c -- 0 pH 8.0 0 0 0 pH 10.5 7.62 2.74 1.65 AEC in cmol kg-1 Initial 111 111 104 Final 0 0 1.98 CEC in cmol kg-1 Initial 0 0 0 Final 2.47 n.a 7.198 AEC in kJ mol-1 3.835 2.47 3.15 CEC in kJ mol-1 2.331 n.a 1.147 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
97
Table A-8. Ensemble of data collected for columns of clean AHO in ZPC experiments. Column 2B4 3B4 Description ZPC ZPC Weight in mg 7 7.4 Al content in % 14.68 17.65 Al content in µmol 38.06 48.37 a Cl/NO3 peak in V ml pH 5.75 21.29 9.98 pH 8.0 5.678 (73.33) 5.16 (48.27) pH 10.5 0 0 K/Na peak a in V ml pH 5.75 0 0 pH 8.0 1.88 1.90 pH 10.5 2.7 5.29 AEC in cmol kg-1 Initial 131 131 Final 0 0 CEC in cmol kg-1 Initial 0 0 Final n.a. 38.35 AEC in kJ mol-1 3.76 0.887 CEC in kJ mol-1 n.a. 1.606 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
98
Table A-9. Ensemble of data collected for columns in ZPC experiments after arsenate adsorption. Column 10B3 11B3 1B3 Description 1 As + ZPC 1 As + ZPC 2 As + pH 8.0 Weight in mg 15.6 15 14.3 Al content in % 15.37 15.27 n.a Al content in µmol 88.8 84.83 n.a As sorbed in µg g-1 n.a 11,667 13,589 As sorbed in µmol n.a 2.33 Al:As mole ratio n.a 36.4 n.a As ∆H in kJ mol-1 n.a 33 a Cl/NO3 peak in V ml 0th 62.25 64.23 43.48 1st As 42.71 (31.39) 38.21 (40.5) 30.31 (30.29) 2nd As -- -- 19.33 (36.23) pH 8.0 4.86 (88.62) 8.8 (76.97) 8.06 (58.3) pH 10.5 0 0 0 K/Na peak a in V ml 1st As 0 0 0 2nd As 0 0 0 pH 8.0 0.23 0.65 3.838 pH 10.5 3.05 8.65 -- AEC in cmol kg-1 Initial 111 111 111 Final 2.756 0 0 CEC in cmol kg-1 Initial 0 0 0 Final 8.43 3.67 0 AEC in kJ mol-1 3.511 4.117 n.a CEC in kJ mol-1 6.312 n.a n.a a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
99
Table A-10. Ensemble of data collected for columns in ZPC experiments after arsenate adsorption. Column 17B2 25B2 9B3 Description 2 As + ZPC 2 As + ZPC 1 As + ZPC Weight in mg 14.7 14.9 15.5 Al content in % 12.9 14.9 16.25 Al content in µmol 70.23 82.23 93.29 As sorbed in µg g-1 10,178 11,030 n.ad As sorbed in µmol 1.99 2.19 n.a Al:As mole ratio 35.29 37.54 n.a As ∆H in kJ mol-1 48.91 15.86 n.a a Cl/NO3 peak in V ml 0th 94.21 57.26 65.32 1st As 66.88 (29) -- 32.18 (50.7) 2nd As 57.06 (39.43) 29.76 (48.03) -- pH 8.0 5.86 (89.73) 2.29 (92.31) 7.38 (77.07) pH 10.5 0 0 0 K/Na peak a in V ml 1st As 0 0 0 2nd As 0 0.05 -- pH 8.0 0 0 0.93 pH 10.5 2.85 4.32 6.34 AEC in cmol kg-1 Initial 104 104 111 Final 0 0 0 CEC in cmol kg-1 Initial 0 0 0 Final 11.23 6.74 6.69 AEC in kJ mol-1 6.426 3.56 3.33 CEC in kJ mol-1 1.8 4.14 5.36 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
100
Table A-11. Ensemble of data collected for columns in ZPC experiments after arsenate adsorption. Column 2B3 12B3 Description 2 As + pH 8.0 1 As + ZPC Weight in mg 16.3 16.1 Al content in % 17.1 16.24 Al content in µmol 103.23 96.8 As sorbed in µg g-1 22,776 20,219 As sorbed in µmol 4.95 4.31 Al:As mole ratio 20.8 22.4 As ∆H in kJ mol-1 31.9 14.77 a Cl/NO3 peak in V ml 0th 69.38 70.43 1st As 47.08 (32.14) 39.09 (44.5) 2nd As 32.4 (31.18) -- pH 8.0 15.49 (52.21) 12.98 (66.8) pH 10.5 -- 0 K/Na peak a in V ml 1st As 0 0 2nd As 0 -- pH 8.0 0 0 pH 10.5 -- 0 AEC in cmol kg-1 Initial 111 111 Final 15.51 0 CEC in cmol kg-1 Initial 0 0 Final 0 11.27 AEC in kJ mol-1 6.629 3.485 CEC in kJ mol-1 0 0 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
101
Table A-12. Ensemble of data collected for columns in ZPC experiments after arsenate adsorption. Column 13 B3 14 B3 15 B3 Description 5 min As 10 min As 15 min As Weight in mg 15 16 15.5 Al content in % 20.63 21.39 20 Al content in µmol 114.6 126.79 115.31 As sorbed in µg g-1 29,066 39,656 67,290 As sorbed in µmol 5.81 8.46 13.91 Al:As mole ratio 19.73 14.99 8.29 As ∆H in kJ mol-1 3.042 6.194 4.69 a Cl/NO3 peak in V ml 0th 55.65 60.38 36.52 1st As 29.93 (46.22) 40.21 (33.4) 20.44 (44) 2nd As ------pH 8.0 ------pH 10.5 ------K/Na peak a in V ml 1st As 0 0.84 2.7 2nd As ------pH 8.0 ------pH 10.5 ------AEC in cmol kg-1 Initial 111 111 111 Final 0 0 0 CEC in cmol kg-1 Initial 0 0 0 Final 0 1.97 3.98 AEC in kJ mol-1 2.995 3.005 1.877 CEC in kJ mol-1 0 2.357 3.869 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
102
Table A-13. Ensemble of data collected for columns in ZPC experiments after arsenate adsorption. Column A0B3 A1B3 A2B3 Description As’ted in batch Batch As + ZPC Batch As + ZPC Weight in mg 11.7 16 14.7 Al content in % 28.07 18.19 27.98 Al content in µmol 121.64 107.61 152.34 As sorbed in µg g-1 22,476 n.a 25,842 As sorbed in µmol 3.51 n.a 5.065 Al:As mole ratio 34.65 n.a 30.07 As ∆H in kJ mol-1 a Cl/NO3 peak in V ml pH 5.75 -- 66.45 0 pH 8.0 -- 4.99 (92.5) 0 pH 10.5 -- 0 0 K/Na peak a in V ml pH 5.75 -- 0 1.804 pH 8.0 -- 4.3 3.056 pH 10.5 -- 11.50 14.34 AEC in cmol kg-1 Initial 104 104 104 Final -- 0 0 CEC in cmol kg-1 Initial -- 0 0 Final -- 15.5 12.93 AEC in kJ mol-1 -- n.a. CEC in kJ mol-1 - 2.338 7.234 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply d not available
103
Table A.14. Ensemble of data collected from columns used in reversibility experiments. Column 19B2 6B3 Weight in mg 14.6 16.1 Al content in % 8.71 14.27 Al content in µmol 47.1 85.09 a Cl/NO3 peak in V ml pH 5.75 62.23 27.02 pH 8.0 7.18 (88.46)b 9.6 (64.47) pH 10.5 0 0 Cl/NO3 peak in V ml pH 10.5 0 0 pH 8.0 2.26 3.03 pH 7.5 --c 16.54 AEC in cmol kg-1 Initial 104 111 Final 12.88 24.41 AEC in kJ mol-1 Initial 3.53 1.66 Final 1.04 4.60 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area c does not apply
Table A.15. Ensemble of data collected for columns used in energetics experiments. Column 20B2 22B2 5B3 7B3 Weight in mg 14.1 15.6 14.6 16.6 Al content in % 19.55 19.37 21 19.9 Al content in µmol 102.1 111.92 113.56 22.35 a Cl/NO3 peak in V ml pH 5.75 101.55 101.8 59.63 51.94 pH 7.25 12 (88.18)b 15.72 (84.55) 22.56 (62.17) 16.2 (68.81) AEC in cmol kg-1 Initial 104 104 111 111 Final 46.91 33 20.14 30.66 AEC in kJ mol-1 Initial 7.15 6.85 3.18 2.50 Final 1.87 3.33 6.62 2.82 a averages of 3 to 6 values after correction for flow rate differences b numbers in parenthesis are percent reductions from previous peak area
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BIOGRAPHICAL SKETCH
Nadine Jack Kabengi was born on a sunny July 3rd day, 1975, in Beirut, Lebanon.
When she did finally speak (around 3 years old after giving a scare to her parents),
Nadine had problems pronouncing many letters and was never expected to become a good student. She did however grew out of her speech impediments, and finished every year first of her class. She graduated high school in 1993 earning a double, French and
Lebanese, baccalaureate. Nadine joined the American University of Beirut in that year as a premedical student ready to fulfill the traditional destiny as a bright student and become a medical doctor. To the chagrin of her parents, she was attracted more to the political life of college than to medical classes. To pursue her extracurricular interests, she shifted from the highly demanding major and joined the Faculty of Agricultural and Food
Sciences, and graduated four years later with a bachelor’s in agriculture and a diploma of agricultural engineer. In her second year, she sat for a soil science class and was so amazed by what she learned, she pursued an M.S in soil science which was conferred in
1999. Armed with more unanswered questions and an insatiable curiosity about her field, she started a Ph.D. program in soil chemistry at the University of Florida. Having worked for three and a half years with two remarkable advisors, Dr. R.D. Rhue and Dr. S.H.
Daroub, she is done now and feels very fortunate to have been given the opportunities she had. Nadine hopes to pursue a career in research in the field about which she is passionate as well as earning a Nobel peace prize for making the world a better place.
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