remediation of barium contaminated groundwater a...

Remediation of Barium Contaminated Groundwater

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Final Year Project

School of Water Research


A Study of Barium Sulphate Mobility

Mala H. Batu

Supervisor: Dr Christoph Hinz

Co-Supervisor: Dr Chris Barber

3rd November 2003


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Abstract

Concentrations of barium exceeding Australian Drinking Water Guidelines have been found

in some aquifers near Perth, Western Australia. Barium has implications for human health,

because of the link between high barium in drinking water and cardiovascular disease. In

general high barium concentrations are found in deep confined aquifers, which are depleted in

sulphate. Therefore the possibility of injection of sulphate-amended water to precipitate

barium sulphate (BaSO4), in situ during aquifer storage and recovery (ASR) is a solution to

reduce concentrations of barium. This paper will investigate the mobility of precipitated

BaSO4, as a colloid through three types of sand with differing grain sizes. The stable colloid

kaolinite will be used as a reference for comparison with the BaSO4.

The mobility of the colloids was studied experimentally using a flow-through small diameter

sand column through which solutions containing colloidal suspensions BaSO4 and kaolinite

were eluted. The absorbance of column eluant was measured with a flow-through UV-VIS

spectrophotometer set at 420nm. The mixing interface between barium chloride and sodium

sulphate solutions was also investigated in the column for a study of in situ precipitation of

BaSO4. The effects of variable sand grain size and flow rate were studied to assess potential

colloid behaviour during ASR.

The results proved that kaolinite is a stable colloid while BaSO4 is unstable. The BaSO4

suspensions produced erratic results, indicating that the BaSO4 colloid was readily

flocculating and attaching to sand grains as well as column tubing. The positively charged

BaSO4 particles would attach and collect on charged mineral surfaces, but occasionally these

become displaced. Kaolinite (which has an overall negative charge) in contrast showed a

smooth elution pattern. The in situ precipitation test showed that BaSO4 was produced when

sodium sulphate was introduced into the barium chloride environment. Measured absorbance

of the eluant showed a characteristic pattern indicating that BaSO4 was accumulating in the

column with successive dispersive mixing events. This may be a reversible reaction as the

BaSO4 may be “bleeding” out from the column. Filtration coefficients were measured and

deduced that columns of smaller sand size filter out more kaolinite and barite colloids.


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Acknowledgements

The author expresses great appreciation for the excellent supervision of Dr Christoph Hinz

and Dr Chris Barber, for their advice, guidance and support throughout the year.

The author would also like to acknowledge Henning Prommer, Edgardo Alarcon Leon and

Michael Smirk.

Finally, the author especially wishes to thank family and friends for their love and constant

support throughout the year.


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Table of Contents

1 INTRODUCTION 1

2 LITERATURE REVIEW 3

2.1 DEFINITION OF A COLLOID 3

2.2 COLLOID FORMATION 4

2.2.1 AGGREGATION 4

2.3 MECHANISMS IN COLLOID TRANSPORT 4

2.3.1 RESTRICTIONS IN COLLOID TRANSPORT 4

2.4 COLLOID MOBILIZATION 6

2.4.1 IONIC STRENGTH 8

2.4.2 EFFECT OF PH 8

2.5 COLLOID COLUMN EXPERIMENTS 9

2.6 BARIUM SULPHATE 10

2.7 BARIUM AS A COLLOID 11

2.8 ROLE OF DISPERSANTS 11

2.9 TURBIDITY 12

2.10 SURFACE CHARGE 13

2.11 BARITE AND PERMEABILITY 14

2.12 BARIUM IN GROUNDWATER 14

3 METHODOLOGY 17

3.1 SAND COLUMN 17

3.1.1 SANDS 17

3.1.2 COLUMN 17

3.1.3 PACKING AND SATURATION 17

3.1.4 TRACER 18

3.2 KAOLINITE 20

3.2.1 CALIBRATION CURVE 20

3.2.2 BREAKTHROUGH CURVES FOR CONSTANT FLOW RATE 21

3.2.3 BREAKTHROUGH CURVES FOR VARIABLE FLOW RATE 22


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3.3.1 CALIBRATION CURVE 23

3.3.2 BREAKTHROUGH CURVES FOR BARIUM SULPHATE SUSPENSION 24

3.3.3 CREATING A MIXING INTERFACE BETWEEN BARIUM CHLORIDE AND SODIUM SULPHATE 24

3.4 CALCULATIONS 26

3.4.1 COLLOID FILTRATION RATE (KF) 26

3.4.2 DISPERSION COEFFICIENT (D) 26

4 RESULTS 27

4.1 SODIUM CHLORIDE TRACER 27

4.2 KAOLINITE 28

4.2.1 CALIBRATION CURVES 28

4.2.2 BREAKTHROUGH CURVES FOR DIFFERENT SAND SIZES 29

4.2.3 VARIABLE FLOW RATE 30


4.3.1 SUSPENSION 32

4.3.2 MIXING INTERFACE 33

5 DISCUSSION 37

5.1 SODIUM CHLORIDE TRACER 37

5.2 KAOLINITE 38

5.2.1 BREAKTHROUGH CURVES FOR DIFFERENT SAND SIZES 38

5.2.2 COLLOID FILTRATION RATE DEPENDENCE ON PORE WATER VELOCITY 39


5.3.1 BEHAVIOUR OF BASO4 40

5.3.2 MIXING INTERFACE 41

6 CONCLUSION 43

7 REFERENCES 45


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List of Figures

FIGURE 2-1 EXPERIMENTAL DETERMINATION OF THE COLLOID FILTRATION RATE (KF) USING A

LATEX-GLASS MODEL WITH A NACL TRACER. (A) STEP-INPUT BREAKTHROUGH CURVES,

(B) PULSE-INPUT BREAKTHROUGH CURVES, AND (C) COLLOID FILTRATION RATE

CALCULATED FROM BOTH METHODS. ......................................................................................7

FIGURE 2-2 CONCEPTUAL VIEW OF (A) INITIAL DEPOSITION, (B) BLOCKING, (C) AND RIPENING

(KRETZSCHMAR ET AL 1999). .................................................................................................7

FIGURE 2-3 SEM MICROGRAPHY OF A SYNTHETIC BASO4 CRYSTAL IN ITS RHOMBOHEDRAL

STRUCTURE (DUNN ET AL 1999) ...........................................................................................10

FIGURE 2-4 PARTICLE SIZE DISTRIBUTION FOR BASO4 (SUN AND SKOLD 2001).............................11

FIGURE 2-5 THE INFLUENCE OF SOLUTION PH ON THE SURFACE CHARGE OF BASO4 (COLLINS

1998)..................................................................................................................................13

FIGURE 2-6 EQUILIBRIUM BETWEEN CONCENTRATIONS OF BARIUM AND SULPHATE (BARBER

AND PROMMER 2002) ..........................................................................................................16

FIGURE 3-1 SCHEMATIC OF APPARATUS FOR NACL TRACER EXPERIMENTS.....................................19

FIGURE 3-2 DISPLACEMENT OF BACL2 BY NA2SO4 IN THE COLUMN TO INVESTIGATE THE MIXING

INTERFACE. .........................................................................................................................25

FIGURE 4-1 BREAKTHROUGH CURVE CONSTRUCTED WITH EC MEASUREMENTS FROM

EXPERIMENTS INVOLVING COLUMN OF 0.5-0.71MM AT TWO DIFFERENT FLOW RATES. .............27

FIGURE 4-2 RELATIONSHIP BETWEEN THE ABSORBANCE (NM) AND CONCENTRATION (PPT) FOR

KAOLINITE. .........................................................................................................................28

FIGURE 4-3 BREAKTHROUGH CURVES FOR 1PPT KAOLINITE IN DIFFERENT SAND GRAINS AT FLOW

RATE 1.25ML/MIN. ..............................................................................................................29

FIGURE 4-4 ABSORBANCE MEASURED WITH PORE VOLUME FOR 1PPT KAOLINITE IN COLUMN

WITH 1.4-2MM SAND SIZE WITH CHANGING FLOW RATE FROM 0.8ML/MIN TO 1.25ML/MIN

TO 1.8ML/MIN. ....................................................................................................................30

FIGURE 4-5 BREAKTHROUGH CURVE FOR SAND SIZE 1.4-2MM WITH 1PPT KAOLINITE FLOWING

THROUGH WHERE THE FLOW RATE CHANGES FROM 0.8ML/MIN- 1.25ML/MIN- 1.5ML/MIN. .....31

FIGURE 4-6 BREAKTHROUGH CURVE FOR 30PPM BASO4 SUSPENSION IN 1.4-2MM SAND COLUMN

AT FLOW RATE 0.52ML/MIN. ................................................................................................32

FIGURE 4-7 ALTERNATING SEQUENCE OF PUMPING IN NA2SO4 INTO A BACL2 COLUMN

ENVIRONMENT CONTAINING A SAND SIZE OF 1.4-2MM (A) 0.5-0.71MM (B) AT FLOW RATE

1ML/MIN.............................................................................................................................34


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FIGURE 4-8 ALTERNATING SEQUENCE OF PUMPING IN NA2SO4 INTO A BACL2 COLUMN

ENVIRONMENT CONTAINING A SAND SIZE OF 0.125-0.25MM AT FLOW RATE 1ML/MIN. ............35

FIGURE 5-1 RELATIONSHIP BETWEEN COLLOID FILTRATION RATE AND PORE WATER VELOCITY

FOR 1.4-2MM SAND..............................................................................................................40


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List of Tables

TABLE 2-1 DRINKING WATER STANDARDS (NH&MRC/ARMCANZ 1994, ROBERTSON 1991,

LANCIOTTI ET AL 1989)........................................................................................................15

TABLE 2-2 AQUIFERS IN PERTH THAT MAY BE DEPLETED IN SULPHATE ..........................................16

TABLE 3-1 COLUMN DIMENSIONS ................................................................................................17

TABLE 4-1 PEAK RELATIVE CONCENTRATIONS AND PORE VOLUMES TO REACH AVERAGE

RELATIVE CONCENTRATION FOR DIFFERENT SAND TYPES WITH 1PPT KAOLINITE FLOWING

AT 1.25ML/MIN. ..................................................................................................................30

TABLE 4-2 ABSORBANCE OF BASELINE LEVELS AFTER EACH INTERCHANGE FOR EACH SAND...........35

TABLE 5-1 THE PORE WATER VELOCITY (CM/MIN), DISPERSIVITY (CM) AND DISPERSION

COEFFICIENTS (CM2/MIN), FOR EACH SAND SIZE.....................................................................37

TABLE 5-2 COLLOID FILTRATION RATES CALCULATED FOR EACH SAND AT A FLOW RATE OF

1.25ML/MIN........................................................................................................................38

TABLE 5-3 COLLOID FILTRATION RATES MEASURED IN 1.4-2MM SAND FOR DIFFERENT PORE

WATER VELOCITIES..............................................................................................................39


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1 Introduction

Elevated barium concentrations have been found in aquifers near Perth, Western Australia

exceeding Australian Drinking Water Guidelines. The Australian non-exceedance guideline

for barium is 0.7mg/L (NH&MARC/ARMCANZ 1994), while concentrations have been

found in the Myalup region of the Yarragadee aquifer in Perth to be 0.8-4mg/L (Barber and

Prommer 2002). High barium concentrations have serious health implications for humans, as

it can increase the risk of constriction of blood vessels and contractions of ailmentary canal,

convulsions and paralysis. Exposure in low doses from drinking water can also increase the

risk of heart attack (NH&MARC/ARMCANZ 1994).

In general high concentrations of barium have been found in deeper, confined aquifers that

are sulphate depleted. This phenomenon occurs in Illinois (Gilkeson et al 1981) and Arizona

(Robertson 1991) in the US; Florence, Italy (Baldi et al 1996); and in China (Zhou and Li

1992). This is caused by the natural dissolution of barium from the aquifer minerals where

sulphate has been bacterial reduced. It can be explained by the solubility of barium sulphate

(BaSO4). BaSO4 precipitate is created by the interaction of Ba2+ ions and SO42- ions (Equation

1-1). The formation of the precipitate is dictated by the solubility of the precipitate, where the

solubility product Ksp is 1.1 x 10-10 (Shaw et al 1998). This value indicates that BaSO4 is very

insoluble. Ksp is derived from multiplying the activities (M) of dissolved Ba2+ and SO42- ions

together (Equation 1-2). So, if an aquifer is depleted in sulphate- the activity of sulphate will

decrease. Therefore the barium activity has to increase- leading to an increase in

concentration of barium. This is also true for the opposing case. If the barium activity

decreases, the activity of the sulphate has to increase. Therefore to decrease the barium

concentrations, the sulphate concentration has to increase.

Equation 1-1 Ba2+(aq) + SO4

2-(aq) →→→→ BaSO4(s)

Equation 1-2 Ksp = activityBa2+ x activitySO42-

A possible solution to lower the barium concentrations is to inject sulphate-amended water to

precipitate BaSO4, in situ during aquifer storage and recharge. In theory, injecting sulphate


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into the sulphate-depleted aquifer should precipitate out BaSO4, reducing the barium

concentrations to allowable levels. This technique has been employed in one simulation study

by Scrivner et al (1996), where sodium sulphate (Na2SO4) and sodium sulphide was added to

soils in a waste landfill to reduce metal concentrations of barium. However this technique has

not been used in the field. Therefore the research needs to be conducted to judge if this is a

feasible option.

This leads to the reason behind this study. The aim of this study is to investigate the

mobility of precipitated BaSO4 as a colloid through three types of sand with differing

grain size, and use a well studied colloid- kaolinite, as a reference for comparison with

BaSO4.

BaSO4 will be studied as a colloid to understand what happens when BaSO4 is formed in an

aquifer, and investigate how it is transported with groundwater flow. Three types of sand

were selected to understand the transport in different environments, but not to represent the

environments in the Perth aquifers, which is quite heterogeneous. A well-studied colloid

kaolinite will be used for comparison purposes, as little work has been done on the mobility

of BaSO4 as a colloid.

Colloid mobility is important in the aquifer especially if the colloid is a contaminant. The

transport of the colloid will indicate how far it may be advected and dispersed throughout the

porous media, showing the extent of contamination in an aquifer. This is especially important

in terms of the pumping of groundwater from the aquifer for drinking water consumption.

However in this case the formation of the colloid will reduce the amount of contamination.

The formation of the BaSO4 colloid during ASR is a relatively new concept to help reduce

contamination. Little is known about the behaviour of this colloid in terms of its formation

and transport through the aquifer. Many processes including adsorption and filtration will

impact heavily on the movement of this colloid. Therefore this study will provide preliminary

information about the behaviour of the colloid, to see if this option is feasible.


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2 Literature Review

This review will cover the general definition of a colloid and the mechanisms involved in the

transport of colloids through a porous media. The geochemistry of barium sulphate will also

be discussed and the previous investigations into the transport of this colloid.

2.1 Definition of a Colloid

Colloidal particles exist within the size range of 1nm to 1µm (Buddemeier & Hunt 1988).

These particles move according to Brownian motion, which is caused by the motion from the

bombardment by fluid molecules that move with a random thermal nature. This Brownian

motion helps keep the colloids in suspension.

Manahan (2000) classifies colloidal particles in terms of:

• Hydrophilic colloids – are large molecules or ions that strongly interact with

water, which results in the spontaneous formation of more colloids. Examples are

proteins and synthetic polymers.

• Hydrophobic colloids – interact less with water and are more stable, as they

have a charge. This surface charge and the opposing charge ions in the surrounding

solution form an electrical double layer. This enables the colloids to repel each other.

Clay particles are a good example of these colloids.

• Association colloids – are composed of miscelles. These are aggregates of

molecules and ions. A hydrocarbon chain may surround a spherical colloidal particle

where the ionic head is attracted to the positive ions in solution, and the tails entrain

the colloid.

The colloid BaSO4 investigated in this study is classified as a hydrophobic colloid.

Colloids have large specific areas usually greater than 10m2/g (Kretzschmar et al 1999). This

enables particles or other colloids to adsorb onto the surface of these colloids, or for the

colloids to adsorb onto grains in the aquifer. In general colloids are relatively stable and will

remain in suspension over large time periods. This will be the case unless the colloids

coagulate or deposit onto the soil. This study will test this theory, as BaSO4 is known to be

unstable in suspension. This will be discussed later in the review. The behaviour of the


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colloids will also depend on the high specific area, high interfacial energy and high surface

charge/ density ratio (Manahan 2000).

2.2 Colloid Formation

Colloids are mostly created by (Kretzschmar et al 1999)-

• in situ by changes in solution chemistry, which causes particles to be released.

• precipitation from saturated solutions

• introduction of “biocolloids” (viruses and bacteria) from anthropogenic inputs.

This study will focus on colloid formation by precipitation from ASR. ASR can produce

drastic changes in water chemistry by injecting sulphate-amended water into the aquifer to

induce the precipitation of BaSO4 colloids. A similar study was performed by Liang et al

(1993) who injected oxygenenated water into an anoxic aquifer to induce the precipitation of

iron (III) hydroxide colloids to reduce the iron (II) concentrations in the aquifer.

2.2.1 Aggregation

Aggregation occurs when colloids attract and form larger particles. This process is prevented

by the electrostatic repulsion of the electrical double layer. Aggregation can be classified into

two sections- coagulation and flocculation.

Coagulation occurs when the repulsion is reduced so much that the colloidal particles can join

and form larger particles. The repulsion is overcome by van der Waals forces. When the

particles come close to other particles through Brownian motion, van der Waals do the rest

and pull the particles together. The particles will then stick together and squeeze out the

solution between the particles. Hydrophyllic colloids tend to be more stable so they will not to

coagulate, while hydrophobic colloids are less stable and will want to coagulate. Flocculation

uses bridging compounds that connect the colloids through chemical bonds. Polyelectrolytes

cause flocculation to occur.

2.3 Mechanisms in Colloid Transport

2.3.1 Restrictions in colloid transport

The following processes restrict colloid movement (McGechan and Lewis 2002):


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• Straining/Filtration

• Adsorption

• Electrostatic charge mechanisms: which include the double diffuse layer (DDL).

Straining occurs where the colloid is greater than the size of the pore, and therefore cannot

pass through. According to Hunt et al (1987) and Ibaraki and Sudickey (1995), filtration is

achieved by particles that are captured within pores that are larger than the colloid. However

other authors refer to this as straining.

Adsorption occurs when reactive substances attach to surfaces of solids e.g. colloids attach

themselves to the soil matrix. The adsorption capacity of a solid depends on the surface area

to volume ratio. If it has high adsorption capacity, it means that the solid is small but has a

large specific area. Therefore there are more adsorption sites for the colloid to attach itself to.

Another important factor is the surface charge of the colloid. Colloids usually behave as

charged particles, which can interact strongly with the soil (McGechan and Lewis 2002).

The diffuse double layer is also known as the electrical double layer. This layer is created by

the arrangement of ions dissociated from a charged surface. These ions are subject to two

forces: adsorption generated by the electric field of the charged surface and diffusive force

formed from the concentration gradients trying to equilibrate the ion concentration throughout

the whole solution (Brady and Weil 1996). It will be formed when the colloid and the particle

it comes into contact with.

The thickness of the layer is usually large (10µm) for monovalent cations that have low ionic

strength; and thin (0.1µm) for multivalent ions that have a high ionic strength (McGechan and

Lewis 2001). Ionic strength will be discussed later. This study does not investigate ionic

strength, but it is recommended that it been done in the future.

To simplify analysis, this study will combine all of these terms and define the process as

filtration. Where filtration is the mechanism of removing colloid out of solution by trapping it

within the pores. This process is composed of transport of the colloid through the porous

media and attachment of colloid to the surface of the soil matrix.


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2.4 Colloid Mobilization

Colloid transport through granular porous media is modelled by the advection-dispersion

equation (Kretzschmar et al 1999):

Equation 2-1Ck

z

Cv

z

CD

t

Cf−

∂

∂−

∂

∂=

∂

∂2

2

Where C(t) is the colloidal concentration over time; D the dispersion coefficient that describes

how the colloids spread through the media; v the pore water velocity; kf the colloid filtration

rate; and z the distance travelled through the porous media.

Based on colloid filtration theory, this study assumes that colloid filtration is an irreversible

process. That indicates that once the colloids have been filtered out by the porous media they

will remain there indefinitely.

Colloid filtration occurs in two steps- first the colloids are transported through the porous

media by Brownian diffusion (transport step); and then the colloids may attach to the soil

matrix in the aquifer (attachment step). The transport step is affected by the size and density

of the colloids; accessibility to the surface area of the soil matrix; pore structure; and the pore

water velocity. Surface and solution chemistry between the colloid and the soil surface will

affect the attachment step. These include the van der Waals forces, electric double layer,

steric repulsion and hydration forces (Kretzschmar et al 1999).

The kf can be calculated from two methods. These are the step-input or short-pulse column

experiments (Figure 2-1). This study will employ the step-input method to determine the kf, as

both methods produce very similar results. Akbour et al (2002) determined the colloid

filtration rates (kf) for kaolinite using the pulse-input method. The rates varied between 1-12

hr-1. Even though this study will use the step-input method, vales will be able to be compared.


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Figure 2-1 Experimental determination of the colloid filtration rate (kf) using a latex-

glass model with a NaCl tracer. (a) Step-input breakthrough curves, (b) pulse-input

breakthrough curves, and (c) colloid filtration rate calculated from both methods.

After the initial attachment to the soil matrix, blocking and ripening processes can occur

(Figure 2-2). Blocking occurs when the colloid-colloid attraction is unfavourable due to

repulsion. Therefore the filtration rate decreases because the mobile colloids will repel against

the colloids that are already attached to the soil matrix. Conversely when colloid-colloid

attraction is favourable, the mobile colloids will attach to the colloids already attached to the

soil matrix. These phenomenon will not be studied exclusively, only the determination of if

the colloids are filtered or not.

Figure 2-2 Conceptual view of (a) initial deposition, (b) blocking, (c) and ripening

(Kretzschmar et al 1999).


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Colloid mobilization is caused by-

• Decrease in ionic strength

• Increasing in pH

• Increase in flow velocity

• High pumping rates

• Rapid infiltration

• Fractured flow

This review will focus on the first three factors, as they are important experimental

parameters.

2.4.1 Ionic Strength

Ionic strength can be reduced by infiltration of precipitation, irrigation or aquifer recharge.

Dilute solutions containing monovalent ions are the most effective at mobilizing colloids. The

dilute solutions will also reduce the hydraulic conductivity. And solutions with high ionic

strength and contain bivalent ions will not cause colloid mobilization. Low ionic strength

solutions will produce large (“thick”) double layers, because the surface charge of the colloid

needs to be balanced by a large layer because the ion concentration is lower. Decreasing the

ionic strength will increase the repulsion, as when two colloids approach each other, more of

the double layers will overlap increasing repulsion and therefore promote colloid mobilization

(Ryan and Elimelech 1996). The ionic strength is important in terms of coagulation occurring,

because high ionic strength enhances coagulation because the attractive van der Waals forces

dominate and the double layers compressed. At low ionic strength a larger repulsive barrier is

formed because there is a large electrostatic repulsion. Therefore there will be less attachment

of colloids to the soil matrix (Kretzschmar et al 1999).

2.4.2 Effect of pH

Kolakowski and Matijevic (1979) showed that chromium hydroxide colloids that were

attached to glass beads during low pH. However as the pH increased from 9.6 to 11.5 the

colloids were released. The increase in pH increases the level of repulsion between the

colloids and the medium. But this is with glass beads. The scope of this project will use

natural sediments.


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An increase in pH promotes colloid mobilization. Laboratory experiments were conducted

with columns containing disturbed and undisturbed sediments. Bunn et al (2002- as cited in

Ryan and Elimelech 1996) experiments were carried out on two columns- one with disturbed

sediments and the other undisturbed sediments. Disturbed meaning the sediment was not

changed; and pH was increased sequentially to the same sediment. Disturbed sediments will

be used in the present study. Undisturbed meaning fresh sediment was used each time the pH

was changed.

Twice as many colloids were released from the disturbed sediment as the undisturbed

sediments. This was attributed to the colloids being released during sediment drying and

coatings on the sand grains were disturbed, releasing colloids attached to it. Roy and

Dzombak (2001) and Grolimund et al (1996) have also produced work that has indicated that

the disturbed sediment releases more colloids that undisturbed.

Increasing the pH increases the electrostatic repulsion between the colloids and the sediment.

Therefore increasing colloid release. However a limit does occur when the pH is 12.5 in

disturbed and 13.1 in undisturbed (Bunn et. al. 2002). At these levels the increasing ionic

strength reduces the electrostatic repulsion and colloids release is diminished.

Bunn et al (1996) hypothesised that for colloid release to occur, the pH of the influent would

have to exceed the pH of the point of zero charge (pHpzc) on the ferric oxyhydroxides. If the

influent pH was below the pHpzc, then the positive ferric hydroxides would bind to the

negative clay colloids. However if the influent pH is above then this bond would become a

repulsion, as the ferric oxyhydoxides become negatively charged. BaSO4 point of zero charge

will be discussed later in the study.

2.5 Colloid Column Experiments

Most experiments conducted with colloids have suspensions containing colloids passed

through soil samples. Material from the bottom of the column as well as the soil column is

then analysed. Lahav and Tropp (1980) used this method with suspensions containing

synthetic latex microspheres of diameter 0.12 and 0.21µm. The column sample contained a


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high concentration of trapped spheres in the surface layers after analysis. Water entering and

leaving the column has also been analysed. From this high macroporosity aided colloid

transport; and the pH in the colloid suspension affected the ionic charge (Seta & Karathanasis

1997). Artificially created fractures have also been constructed in the soil samples. Toran and

Palumbo (1992) compared transport with and without fractures, however they found that

colloids smaller than the pores were being captured. This was attributed to chemical effects.

Experiments with metals involved the migration of copper and zinc in soil columns

containing clay and silty clay (Karanthanasis 1999). The transport of the metals was enhanced

with suspensions containing colloids by 5-50 fold, compared to suspensions with no colloids.

2.6 Barium Sulphate

Barium sulphate (BaSO4) is formed according to the following reaction-

Equation 2-2 BaSO4 ↔↔↔↔ Ba2+ + SO42-

where the equilibrium constant and solubility product, Ksp = 1.23 x 10-10. The solubility, S =

1.1 x 10-5 mol/L (Manahan 2000).

BaSO4 is typically an orthorhombic crystal, but it can also form several other different shapes

(Figure 2-3). Collins (1998) measured an average size of 16µm.

Figure 2-3 SEM micrography of a synthetic BaSO4 crystal in its rhombohedral structure

(Dunn et al 1999)


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Figure 2-4 Particle size distribution for BaSO4 (Sun and Skold 2001)

A particle size distribution was formed from a 5mM BaSO4 suspension in Sun and Skold

(2001) work. The average effective diameter of the particles was 156nm with a range of 125-

225nm (Figure 2-4).

SEM micrographs and scintillation counters were used to measure the barium particle sizes

(Aliaga et. al. 1989). However they produced significantly different results. The SEM

micrographs measured particles with size 20µm, while the counter measured particles in the

range of 0.5-5µm. This inconsistency was attributed to the irregular sizes of the particles.

2.7 Barium as a Colloid

Synthetic BaSO4 is the one of the most insoluble sulphate minerals, as it has a pKsp of 9.96

(Bishop 1988). It is commonly used in powder coatings, paints, inks, rubber, pigments,

storage batteries, plastics, paper etc. BaSO4 is formed by three steps- nucleation, crystal

growth and aggregation. The size of the BaSO4 particles formed is dependent on the

conditions during these phases.

2.8 Role of Dispersants

Sun and Sköld (2001) created barium sulphate (BaSO4) particles by precipitation and used

light scattering measurements to determine its turbidity. Turbidity was measured for a wide

range of temperatures and concentrations. Titrating BaCl2 and NaSO4 created the solution of

BaSO4, while light scattering, pH and conductivity were measured throughout the process.


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However BaSO4 is a very unstable mineral, so two different dispersants were used in separate

experiments. The dispersants were a- 5% aqueous solution of non-ionic surfactant, and a 2%

aqueous solution of polyacryl amide. Polyacryl amide (PAM) is used as a flocculent and a

stabilizer. This depends on the concentration- at low concentrations it acts as a flocculent, and

at high concentrations a stabilizer.

The BaSO4 suspension had a pH of approximately 7.4 throughout all the experiments. Results

show that BaSO4 precipitate is stable throughout the pH range of 0-14. The main findings

include-

1 . BaSO4 particles aggregated forming larger particles in the absence of a dispersant.

Therefore the BaSO4 suspension becomes unstable.

2. In the presence of the non-ionic surfactant, aggregation does not occur beyond 50°C,

therefore forming a stable suspension. The surfactant disperses the particles and stops

them from aggregating.

3. For BaSO4 to form a concentration of 1.5mM is needed in the presence of a PAM.

Therefore the critical nucleation is 1.5mM i.e. turbidity at this concentration leads to

nucleation and crystal growth. It is inferred that this is due to the Ba2+ sequestering action

of the PAM.

From this work it can be deduced that BaSO4 is unstable, but can be stabilised with use of a

non-ionic surfactant or PAM. However these will not be used in this study, as this does not

represent what is happening in the field. In is not practical to add a one these agents to the

aquifer to stabilise the BaSO4 colloid.

2.9 Turbidity

Turbidity is associated with light scattering, because when a beam of light passes through a

solution, the light will scatter when it hits the BaSO4 particles. The scattering reduces the

intensity of the light. Therefore the turbidity depends on the concentration and size of the

BaSO4 particles.

In this present study will use turbidity measurements to show firstly that BaSO4 colloid

particles are forming in the sand columns. The analytical method to form BaSO4 will also be

used.


13

2.10 Surface Charge

The surface charge in soils is important in colloid transport (Brady & Weil 1996). The

colloids can be seen as charged particles that interact with the soils. The stability of a colloid

depends on hydration and the surface charge. Hydration prevents contact with the colloidal

particles, which results in the formation of larger particles. The surface charge determines

whether aggregation occurs. If the particles have the same charge they will repel, and if they

have opposite charges they will attract. In terms of surface charge, when there are excess

sulphate ions, they adsorb onto the BaSO4 particles, causing an overall negative charge. If

there are excess barium ions, then a positive charge will result.

The pH will affect the surface charge of a colloid. The point of zero charge occurs when at a

given pH, when the overall net charge of a colloid will be zero (Manahan 2000). This will

favour the onset of aggregation and precipitation. DLVO theory says that highly charged

particles will form stable suspensions, while low surface charge will allow particles to come

together and form larger particles. Collins (1998) measured the surface electrical properties of

barite, and investigated the effect polyasparate has on the surface charge. This polyasparate

has functions to help inhibit barium sulphate precipitate. Collins (1998) shows that barite is

positively charged below pH of 5 (and negatively charged above). That is the point of zero

charge for BaSO4 is 5 (Figure 2-5).

Figure 2-5 The influence of solution pH on the surface charge of BaSO4 (Collins 1998).


14

In theory below a pH of 5 the colloid will be positively charged and will attract to a

negatively charged soil matrix. While above a pH of 5 it will repel from the soil matrix.

Affect of changing the pH around the point of zero charge will not be analysed in this study.

2.11 Barite and Permeability

Aliaga et al (1989) investigated how permeability was reduced in sandpacks with the aid of

barium and calcium sulphate precipitation. Usually one solution of barium ions and another

containing sulphate ions was pumped at 0.2cm3/min simultaneously into the sandpacks. This

resulted in a precipitate forming inside the sandpack, when the two solutions encountered

each other. Some of this precipitate would then collect within the sand, causing the

permeability to decrease. A 60% reduction was recorded in the permeability trend. However

this trend did oscillate due to the instability of the precipitate. Once again another study has

found that BaSO4 colloid was unstable. This shows that particle bridges were forming and

breaking throughout the experiment. However these oscillations decrease as the pore size is

decreased.

The group also experimented with calcium sulphate or potassium sulphate already mixed into

the sandpacks. Then barium chloride was injected into the sandpacks, causing precipitation of

barium sulphate. A permeability decrease of approximately 60% was also found in this

experiment.

Experiments performed using a CT scan enabled the detection of the wave-like progress of

barium sulphate through the sandpack. The leading edge of the barium sulphate was spread

out- more than what would be accounted for by dispersion or the linear function deduced by

the group. The CT scan was also used to display the path that the barium sulphate takes in the

sandpack. It does not progress uniformly, but a tortuous path.

2.12 Barium in Groundwater

High barium concentrations in aquifers is not only localised to Perth, as occurrences of

elevated concentrations of barium have been found all over the world. In the US, (where the

guideline level is 1mg/L (Table 2-1), high concentrations have been found in New Jersey

aquifers (Czarnik and Kozinski 1994); public supply wells in Illinois (Voelker 1989); and


15

alluvial basins in Arizona, Nevada, New Mexico and California (Robertson 1991). Barium

was found to exceed the European Union standards (Table 2-1) in Birmingham, UK (Ford and

Telham 1994) and the Netherlands (Frapporti et al 1996); 30% of St Petersberg aquifers

exceeded 2mg/L (Barvish and Shvarts undated); and Tuscany, Italy contained barium

concentrations ranging from 7µg/L-1160µg/L.

Table 2-1 Drinking Water Standards (NH&MRC/ARMCANZ 1994, Robertson 1991,

Lanciotti et al 1989)

Country/Organisation Drinking Water Guidelines (mg/L)

Australia 0.7

USA 1

European Union (EU) 0.1

World Health Organisation (WHO) 0.7

Work conducted by Barber and Prommer (2002) established the equilibrium between barium

and sulphate concentrations (Figure 2-6). Therefore concentrations below Australian

guidelines have an equilibrium of 20mg/L of sulphate. So aquifers below 20mg/L of sulphate

were considered depleted of sulphate. Therefore concentration needs to be raised to 100mg/L

to reduce concentrations of barium to below the Australian guidelines if equilibrium

established. Therefore suggest in situ treatment of barium by ASR where sulphate containing

up to 50-100mg/L would induce precipitate of barium sulphate in a short storage period.


16

0

1

2

3

4

5

6

7

0 200 400 600

Sulphate mg/L

log

Bar

ium

ug

/L

Figure 2-6 Equilibrium between concentrations of barium and sulphate (Barber and

Prommer 2002)

Barber and Prommer (2002) also compiled some of the aquifers in Perth that may be depleted

in sulphate (Table 2-1). There are three formations in Perth that have recorded depleted

sulphate levels with depleted sulphate, which would indicate elevated concentration of

barium. Superficial aquifers have less than 20mg/L but due to an unconfined aquifer and the

existence of silicate minerals such as feldspars, this is not likely. The only barium data for the

formations was for the Myalup region. This region recorded elevated concentrations of

barium 0.8-4mg/L, well above 0.7mg/L.

Table 2-2 Aquifers in Perth that may be depleted in sulphate

Formation Region Range of SO4 2-(mg/L)

Mirrabooka - 1-18

Gwelup 5-26

Mirrabooka 3-90

Leederville

Wanneroo 0-19

Wanneroo 0-39Yarragadee

Myalup 5-19*

10 mg/L

1 mg/L

0.1 mg/L

0.01 mg/L


17

3 Methodology

3.1 Sand Column

3.1.1 Sands

Three sand size grains were investigated in this study-

• 1.4-2mm

• 0.5-0.71mm

• 0.125-0.25mm

These three were chosen to understand colloid mobility in a large sand grain, a medium sand

and a small sand grain.

3.1.2 Column

Three glass columns with the same dimensions were constructed to house the three sizes of

sands. The column properties are listed below (Table 3-1). Each column has tubing and

valves attached that will control flow to and from the column. A valve is placed before and

after the column, so that flow can bypass the column if necessary.

Table 3-1 Column dimensions

Inner Diameter (cm) 1

Length (cm) 7.5

Volume (cm3) 5.89

3.1.3 Packing and Saturation

Each column was packed with a different sand size. Care was taken when packing the column

with sand, as the sand needs to be packed tightly and uniformly so that the flow travels

uniformly throughout the column. The column was then saturated with deionised water (DI)

ready for experiments.

3.1.3.1 Materials

The following materials are needed to pack and saturate the columns-

• 3 x columns


18

• 3 x sand sizes (1.4-2mm, 0.5-0.71mm, 0.125-0.25mm)

• Retort Stand

• Plunger

• CO2

• DI water

3.1.3.2 Procedure

1. Wash and dry column thoroughly.

2. Weigh column and tubing associated with it.

3. Place column upright in a retort stand.

4. Start to fill column with some sand. Use a plunger to pack sand and level it out in the

column.

5. Add some more sand and use the plunger to pack it.

6. Repeat step 4 until the entire column is packed.

7. Seal column and weigh.

8. Now fill the column with CO2. This is used to push out all of the air from the column.

9. Saturate the column with DI water by either distillation or pumping water through at a

very slow flow. The DI water will absorb the CO2 producing a saturated column. Care

must be taken not to let air back into the column.

3.1.4 Tracer

A sodium chloride (NaCl) tracer was used to determine the dispersivity (cm) of each sand

size. Dispersivity is a material property, and from this the Dispersion Coefficient (cm2/min)

can be derived. This coefficient is used to describe the extent that colloids will spread

throughout the porous media. 0.2mM NaCl was used as a background solution, and then a

pulse of 0.5mM NaCl entered the column. This tracer was used to ensure that it would be

unreactive in the column, and that the background and pulse had similar ionic strength. Two

flow rates were chosen (one high, and one low), so that the appropriate range of dispersion

coefficients values could be calculated.

3.1.4.1 Materials

• 3 columns with differing sand size

• 0.2mM NaCl


19

• 0.5mM NaCl

• Flow-through electrical conductivity detector

• Peristaltic pump

• Timer

3.1.4.2 Procedure

Figure 3-1 Schematic of apparatus for NaCl tracer experiments

The following should be conducted for each of the columns.

1. Set up apparatus as shown (Figure 3-1).

2. Flush column with 0.2mM NaCl solution. Ensure that no air enters the column.

3. Set flow rate to 1.4mL/min.

4. Simultaneously input pulse into column and start EC measurements at an interval of

8secs. Again ensure that no air enters the column.

5. Input pulse for exactly 5mins. Therefore 7mL of 0.5mM flows through the column.

6. After 5mins replace the pulse with the background solution.

7. Stop recording EC measurements once a stable reading has been reached.

8. Repeat steps 2-8 for flow rate 0.6mL/min.

UV-VIS PUMP

Column

Beaker 1 Beaker 2

Valve


20

3.2 Kaolinite

3.2.1 Calibration Curve

Calibration curve was devised to convert absorbance measured by the UV-VIS

spectrophotometer to concentration. Producing kaolinite solutions of differing concentrations

and measuring their corresponding absorbance with the UV-VIS spectrophotometer achieved

this. This relationship was then plotted to find a linear relationship, which should have a

correlation coefficient greater than 0.98 to ensure that the relationship is strong.

Concentrations of kaolinite were chosen so that there was a broad range, but also that a

relationship could be formed from the values. Therefore the concentrations were chosen at

regular intervals between 0.1-5ppt. The solutions were not too turbid, so that the absorbance

detected would be too high for the UV-VIS spectrophotometer to measure. Also solutions

with too high a concentration deviate from a linear relationship with absorbance.

3.2.1.1 Materials

• 10mL of each Kaolinite concentration- 5, 2.5, 1, 0.75, 0.5, 0.1ppt.

• 5 x 4.5 mL cubettes

• Magnetic Stirrer

• Stirrer Piece

• UV-VIS Spectrophotometer

3.2.1.2 Procedure

1. Fill a cubette with the 5ppt kaolinite solution and measure the absorbance in the UV-

VIS spectrophotometer.

2. Place the solution back into the 5ppt kaolinite container and place the container on the

magnetic stirrer with stirrer piece inside.

3. Turn on stirrer and mix for 1min.

4. Take container off stirrer and fill the same cubette with the solution.

5. Place cubette in the UV-VIS spectrophotometer and measure the absorbance.

6. Repeat steps 1-9 for the following concentrations of kaolinite- 2.5ppt, 1ppt, 0.75ppt,

0.5 ppt and 0.1ppt.


21

3.2.2 Breakthrough Curves for Constant Flow Rate

Breakthrough curves were constructed for-

• each sand size

• different concentrations

The following method and procedure will outline how to produce a breakthrough curve for

each column at one flow rate (1.25mL/min) and at one concentration (1ppt kaolinite).

Breakthrough curves are used to determine the colloid filtration rate (kf). kf is the rate at which

colloids are filtered out of solution by the porous media. This is very important in

understanding colloid mobility in an aquifer.

Having different concentrations, as this is to represent what is happening in the aquifer. It also

to see the affect different concentrations and flow rates on colloid transport. Concentrations

were chosen so that they were within the sensitivity range of the UV-VIS. Not enough of a

turbid solution would produce no absorbance readings and a too turbid solution would be

beyond the limits of the UV-VIS detectable range. Also the concentration would have to have

a linear relationship with absorbance (Section 3.2.1).

Flow rates were chosen depending on how many pore volumes would flow through the

column in a minute. If too many pore volumes flowed through in a minute, the UV-VIS

would not be able to represent what truly was happening the in column. The interval that the

UV-VIS would measure absorbance may “miss” what is coming out of the column. Therefore

a flow rate 1.25mL/min ensured that a pore volume would exit the column every 1.1-2.2

minutes for the different sand sizes.

3.2.2.1 Materials

• 1ppt kaolinite solution

• DI water

• Sand column

• Peristaltic pump

• Flow-through UV-VIS spectrophotometer

• Timer


22

3.2.2.2 Procedure

First the initial concentration of kaolinite (CO) must be measured before it enters the column.

Then the following is conducted-

1. Set-up the experimental apparatus as shown (Figure 2-1).

2. Flush column thoroughly with DI water in beaker 1.

3. Set flow rate to 1.25mL/min.

4 . Simultaneously start UV-VIS measurements and input the kaolinite solution by

switching to beaker 2.

5. Once a stable reading has been reached, switch back to beaker 1.

6. Again once a stable reading has been met, stop recording absorbance.

The maximum absorbance is the initial absorbance of the kaolinite (AO), which can then be

converted to CO using the linear relationship in the calibration curve (Section 3.2.1).

Now to see the affects with the column, change the experimental set-up so that solution now

flows through the column. Repeat steps 1-5 to record the absorbance from the column. These

are the A(t) readings which can then be converted to C(t). This procedure can be conducted

for each column. To construct the breakthrough curve pore volume is plotted against C(t)/CO.

3.2.3 Breakthrough Curves for Variable Flow Rate

This experiment was devised to understand if kf was affected by a changing flow rate. A slow

flow rate (0.8mL/min) was used in the beginning of the experiment and then it was gradually

increased to 1.25mL/min and then 1.8mL/min during a pulse of kaolinite to the column.

The same experimental apparatus was used in Section 3.2.3. The initial concentration was

again measured before it enters the column, but this time the following procedure was

implemented-

1. Set flow rate to 0.8mL/min and set path to beaker 1 to flush column.

2. Simultaneously start measuring absorbance and switch path to beaker 2.

3. Once a stable reading has been reached, change flow rate to 1.25mL/min.

4. Again once a stable reading has been reached, increase the flow rate to 1.8mL/min.

5. Once a stable reading is met, switch to beaker 1.

6. Stop measuring absorbance once the reading is stable.


23

3.3 Barium Sulphate

3.3.1 Calibration Curve

Calibration curve was devised to convert absorbance measured by the UV-VIS

spectrophotometer to concentration. The materials and procedure used to create the

calibration curve was adapted from the analytical Turbidimetric method (Clesceri et al 1998).

Producing BaSO4 solutions of differing concentrations and measuring their corresponding

absorbance with the UV-VIS spectrophotometer achieved this. This relationship was then

plotted to find a linear relationship, which should have a correlation coefficient greater than

0.98.

3.3.1.1 Materials

The following materials were used in the creation of the calibration curve-

• 10mL of each Na2SO4 concentrations- 10, 7.5, 5, 2.5, 1ppm.

• 10mL of Buffer B

• BaCl2 (Dihydrate) Analytical Reagent

• 5 x 30mL containers with cap

• 5 x 4.5 mL cubettes

• Weighing Balance

• Magnetic Stirrer

• Stirrer Piece

• UV-VIS Spectrophotometer

3.3.1.2 Procedure

7. Take 10mL of 10ppm Na2SO4 and place in a small 30mL container.

8. Add 2mL of Buffer B solution. Screw cap on container and shake.

9 . Fill a cubette with the solution and measure the absorbance in the UV-VIS

spectrophotometer.

10. Place the solution back into the 10ppm Na2SO4 container and place the container on

the magnetic stirrer with stirrer piece inside.

11. Turn on stirrer.


24

12. Measure 1.25g of BaCl2 and add to 10ppm Na2SO4 container and mix for 1 minute

exactly. This is to ensure all the BaCl2 is dissolved inside the solution

13. Take container off stirrer and fill the same cubette with the solution.

14. Leave for exactly 5mins.

15. Place cubette in the UV-VIS spectrophotometer and measure the absorbance.

16. Repeat steps 1-9 for the following concentrations of Na2SO4- 7.5ppm, 5ppm, 2.5ppm

and 1ppm.

Timing is important in this experiment, as the BaSO4 is forming and precipitating out of

solution. Mixing allows for the Na2SO4 and BaCl2 to fully mix, and allowing it to settle for 5

mins ensures that the same amount of precipitate is formed for each concentration of solution.

3.3.2 Breakthrough Curves for Barium Sulphate Suspension

The same materials and procedure were used as in Section 3.2.2 except that-

• the background solution was 10ppm Na2SO4 (beaker 1)

• the pulse was 10ppm BaSO4 suspension (beaker 2), which was constantly mixed by a

magnetic stirrer.

However results from this experiment proved erratic, so the concentration of BaSO4 was

increased to 30ppm to see if the sensitivity range was the problem. Other concentrations were

experimented, which also produced the same results. Explanations of the erratic results are

discussed in Section 5.3.1. Therefore another method of experimenting with BaSO4 colloids

was found.

3.3.3 Creating a Mixing Interface between Barium Chloride and Sodium Sulphate

Since creating BaSO4 outside the column and then letting it infiltrating the column produced

erratic results, another method was chosen. The BaSO4 would be produced within the column

by flushing the column with BaCl2, and then displacing it with a Na2SO4 solution (

Figure 3-2). BaSO4 should be formed at the interface of the two solutions. This is

representative of the situation of ASR. Sulphate-amended water is used to displace the water

in aquifer, which contains barium. It will also give an idea of how much colloids are produced

in the interface and their mobility with the interface. The affects of dispersion should be

easily investigated too. Once Na2SO4 has displaced the BaCl2, the reverse was done. BaCl2

then displaced Na2SO4 to see what affect this process had.


25

Figure 3-2 Displacement of BaCl2 by Na2SO4 in the column to investigate the mixing

interface.

3.3.3.1 Materials

The same experimental apparatus is used as in Section 3.2.2 except-

• beaker 1 contained 9mM BaCl2 constantly mixed by a magnetic stirrer.

• beaker 2 consisted of 10mM Na2SO4.

The concentrations were chosen so that SO42- was in excess of Ba2+. This is to maintain that

all the Ba2+ would react and form BaSO4.

3.3.3.2 Procedure

To create a mixing interface the following was conducted in each column-

1. Flush the column with BaCl2.

2. Set the flow rate to 1.25mL/min.

3. Simultaneously start recording absorbance and switch flow to beaker 2.

4. Once a stable reading has been established, switch back to beaker 1.

5. Again when the reading is stable switch back to beaker 2.

6. Repeat steps 4-5 twice.

7. Once the reading is stable stop measurements.

Throughout this experiment the pH and EC was measured every 3mins.

BaCl2

Na2SO4

Flow

Mixing Zone


26

3.4 Calculations

3.4.1 Colloid Filtration Rate (kf)

Colloid deposition rates are determined from step-input or short-pulse column experiments

(Kretzschmar et al 1999). This study has assumed that this is a filtration process not a

deposition process, so it will refer to the colloid deposition rate as the colloid filtration rate.

Again this study has focused on step-input experiments to derive the breakthrough curves.

The following method to determine the kf is adapted from Kretzschmar et al (1999) work.

Columns that have a high Peclet Number can calculate the colloid filtration rate from:

Equation 3-1 FL

vk p

f ln−=

where vp is the pore water velocity, L the length of the column and F the ratio between C/CO

once the plateau has reached a stable value.

3.4.2 Dispersion Coefficient (D)

The Dispersion Coefficient was derived using a spreadsheet created by Bromly (2003) based

on Yu et al (1999) moment method. It can be calculated by measuring the EC during the

NaCl tracer experiments (Section 3.1.4).


27

4 Results

4.1 Sodium Chloride Tracer

Figure 4-1 Breakthrough curve constructed with EC measurements from experiments

involving column of 0.5-0.71mm at two different flow rates.

Electrical conductivity (EC) measurements were made by adding a pulse of 0.5mM NaCl

solution to a background solution of 0.2mM NaCl in each of the columns (Figure 4-1). These

experiments were performed at two different flow rates- 0.6 and 1.4mL/min. Results show

that the EC measurements are consistent with a breakthrough pattern. Plotting relative EC, the

maximum critical value is 1 in all cases.

Changing the flow rate in each sand size did not make an impact on the results, as the same

pulse volume was added to each column. However the gradient at which the EC increases is

steeper for the larger sand size columns.

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 1 2 3 4 5 6 7 8

Pore Volume

Re

lati

ve

EC

1.4mL/min 0.6mL/min


28

4.2 Kaolinite

4.2.1 Calibration Curves

y = 1.6916x - 0.0897

R2 = 0.9964

-0.5

0

0.5

1

1.5

2

2.5

3

0 0.5 1 1.5 2

Absorbance (nm)

Co

nce

ntr

atio

n (

pp

t)

Figure 4-2 Relationship between the absorbance (nm) and concentration (ppt) for

kaolinite.

A relationship was established between the concentration of the differing kaolinite solutions

and the absorbance, which was measured with a UV-VIS spectrophotometer (Figure 4-2).

Fitting a linear relationship to this data produced a correlation coefficient of 0.9964 indicating

a strong linear relationship. Therefore the relationship exists such that-

Equation 4-1 Concentration = 1.6916 x Absorbance – 0.0897

This relationship is strong but has a negative intercept of 0.0897, which is unusual. This was

ignored, as this value is small in comparison to the concentration used.


29

4.2.2 Breakthrough Curves for Different Sand Sizes

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 2 4 6 8 10 12 14 16 18 20

Pore Volumes

C/C

o

1.4-2mm 0.5-0.71mm 0.125-0.25mm

Figure 4-3 Breakthrough curves for 1ppt kaolinite in different sand grains at flow rate

1.25mL/min.

The breakthrough curves for 1ppt kaolinite solution in different sand with a constant flow rate

of 1.25mL/min are shown above (Figure 4-3). The largest sand size (1.4-2.0mm) produced the

highest peak of relative concentration with a value of 0.9695. The middle sand grain size (0.5-

0.71mm) produced a peak of 0.948, and the smallest grain size (0.125-0.250mm) had the

smallest relative concentration peak- 0.942 (Table 4.1).


30

Table 4-1 Peak relative concentrations and pore volumes to reach average relative

concentration for different sand types with 1ppt kaolinite flowing at 1.25mL/min.

Sand Size (mm) Peak Relative Concentration

1.4-2 0.9695

0.5-0.71 0.9480

0.125-0.25 0.9420

4.2.3 Variable Flow Rate

Figure 4-4 Absorbance measured with pore volume for 1ppt Kaolinite in column with

1.4-2mm sand size with changing flow rate from 0.8mL/min to 1.25mL/min to

1.8mL/min.

0.000

0.100

0.200

0.300

0.400

0.500

0.600

0.700

0 5 10 15 20 25 30 35 40 45

Pore Volumes

Ab

so

rban

ce (

nm

)

Flow rate change

(0.8-1.25mL/min)

Flow rate change

(1.25-1.8mL/min)


31

Increasing the flow rate in a column will increase the concentration of kaolinite that is

produced in the eluant (Figure 4-4). As soon as the flow rate is increased from 0.8mL/min to

1.25mL/min or 1.25mL/min to 1.5mL/min, a rise in concentration is recorded by the UV-VIS

spectrophotometer. The change in flow rate from 0.8mL/min to 1.25mL/min occurs at 18.75

pore volumes, and the increase from 1.25mL/min to 1.5mL/min takes place at 31.875 pore

volumes.

Figure 4-5 Breakthrough curve for sand size 1.4-2mm with 1ppt kaolinite flowing

through where the flow rate changes from 0.8mL/min- 1.25mL/min- 1.5mL/min.

From the measured concentration a breakthrough curve can be created for changing the flow

rate within the 1.4-2mm sand column (Figure 4-5). When the flow rate changes during the

experiment, the peak in relative concentration increases. The stable plateau reaches 0.867 at

flow rate 0.8mL/min; a reading of 0.966 for 1.25mL/min; and 0.989 at 1.8mL/min.

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 5 10 15 20 25 30 35 40 45

Pore Volume

C/C

o


32

4.3 Barium Sulphate

4.3.1 Suspension

This graph represents one of many results that were obtained from pumping a BaSO 4

suspension through a sand column. Results were not reproducible, however this was one of

the “better” data that shows the erratic behaviour of BaSO4. This breakthrough curve displays

a 30ppm BaSO4 suspension flowing through the 1.4-2mm sand column at a flow rate of

0.52mL/min (Figure 4-6). There is no constant peak or plateau produced in the data, only

“jagged” peaks. Initially there is a marked increase in concentration near 2 pore volumes, but

then the concentration “zig zags” until reaching a constant relative concentration of around

0.365. However this does not remain constant, as the relative concentration fluctuates. If any

quantitative analysis could be made it would be the fact that the relative concentration hovers

between 0.3 and 0.5.

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 1 2 3 4 5 6 7 8 9 10 11

Pore Volume

C/C

o

Figure 4-6 Breakthrough Curve for 30ppm BaSO4 suspension in 1.4-2mm sand column

at flow rate 0.52mL/min.


33

4.3.2 Mixing Interface

The mixing interface between BaCl2 and Na2SO4 was observed by pumping Na2SO4 into a

BaCl2 environment, and then once a Na2SO4 environment had been established, BaCl2 was

then pumped in. This was repeated a number of times. These experiments were conducted in

each sand for a constant flow rate of 1mL/min. Experiments performed with 1.4-2mm and

0.5-0.71mm produced marked peaks each time the environment changed from either BaCl2 to

Na2SO4, or Na2SO4 to BaCl2 (Figure 4-7). The red lines indicate when the environment was

changed. When comparing these peak heights in the two sand sizes, no relationship could be

established. This is because all the peaks in the larger sand reached absorbance values near

0.23-0.27 nm, yet peak heights in the 0.5-0.71mm sand did not reach a common range.

For the medium column the first peak from Na2SO4 entering BaCl2 column was 0.25nm but

the next peak (when BaCl2 entered Na2SO4 environment) reached 0.43nm. The peak then

dropped for then next interchange to 0.32nm but then increased to 0.40nm. An alternating

peak sequence was reproduced in the 1.4-2mm column, as the peaks increased when Na2SO4

entered the BaCl2 environment not the other way around as seen in the 0.5-0.71mm column.

In both columns there is am increasing trend of the baseline levels of absorbance recorded

after each environment change (Table 4-2). Both baselines levels of absorbance begin at 0nm

but then they increase once the first environment has been changed. Column with 1.4-2mm

has it’s baseline level increased to 0.02nm, and column 0.5-0.71mm to around 0.04nm after

the first interchange. The baseline levels seem to increase by approximately 0.25-0.3nm in the

1.4-2mm sand, and increases by about 0.45nm in the 0.5-0.71mm sand after each interchange.


34

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4

0.45

0.5

0 5 10 15 20 25 30 35 40

Pore Volume

Ab

sorb

ance

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4

0.45

0.5

0 5 10 15 20 25 30 35 40

Pore Volume

Ab

sorb

ance

Figure 4-7 Alternating sequence of pumping in Na2SO4 into a BaCl2 column

environment containing a sand size of 1.4-2mm (a) 0.5-0.71mm (b) at flow rate

1mL/min.


35

Table 4-2 Absorbance of baseline levels after each interchange for each sand

Absorbance (nm) Width of peaks (pore volumes)

1.4-2mm sand 0.5-0.71mm sand 1.4-2mm sand 0.5-0.71mm

sand

1st interchange 0.02 0.04 2.16 1.35

2nd interchange 0.05 0.08 2.5 2.4

3rd interchange 0.07 0.12 2 1.8

4th interchange 0.10 0.17 2.4 1.5

5th interchange 0.12 - 2.1 -

The width of the peaks can also be compared with these sands. The larger sand produced

peaks over a larger pore volume than the medium sand. The 1.4-2mm sand recorded an

average of 2.2 pore volumes while the 0.5-0.71mm sand produced peaks over 1.7 pore

volumes.

The results from the smallest column (0.125-0.25mm) produced no peaks in absorbance, but

did share this increasing trend in baseline values with the other sands (Figure 4-8).

0

0.025

0.05

0.075

0.1

0.125

0.15

0 2 4 6 8 10 12 14 16 18 20

Pore Volume

Ab

sorb

ance

(n

m)

Figure 4-8 Alternating sequence of pumping in Na2SO4 into a BaCl2 column

environment containing a sand size of 0.125-0.25mm at flow rate 1mL/min.


36

Baseline levels increased initially by approximately 0.03nm but then the rate started to

decrease. The difference at the last interchange was 0.02nm.


37

5 Discussion

5.1 Sodium Chloride Tracer

EC measurements were used to calculate the dispersivity (λ) and the dispersion coefficient

(D) for each size of sand. For the results to make sense the relative EC should have a critical

value of 1. This indicates that none of the pulse of 0.5mM NaCl is being filtered by the

column- i.e. whatever enters the column leaves the column.

A spreadsheet devised by Bromly (2003) uses the Yu et al (1999) moment method to

calculate the λ and D from the EC measurements. This spreadsheet was used in this study for

this same purpose by entering values for EC, properties of the sand and the flow rate. The λ

for each sand is listed below (Table 5-1).

Table 5-1 The pore water velocity (cm/min), dispersivity (cm) and dispersion coefficients

(cm2/min), for each sand size.

Sand Size

(mm)

Pore Water

Velocity

(cm/min)

Dispersivity

(cm)

Comparison of

dispersivity

(cm)

Dispersion

Coefficient

(cm2/min)

1.4-2 1.778 0.148 0.278 0.268

0.5-0.71 2.225 0.136 0.249 0.302

0.125-0.25 2.281 0.134 0.221 0.306

The dispersivity values for comparison were obtained from Bromly (2003). The dispersivity

values from this study were lower in all cases, but still in the same magnitude. They also

decreased with decreasing sand size indicating that colloids will spread or disperse more in a

column that has larger sand grains. Therefore it is easier for the colloids to disperse in a pore

volume that is larger and has fewer obstructions in the way to restrict colloid movement.

The dispersion coefficient was calculated by multiplying the pore water velocity with the

dispersivity. Obviously it too will increase for larger sand grains. The pore water velocity

increases with sand size, as the pore volume decreases. The flow rate remains the same, but


38

since the volume the solution has to travel is smaller, the velocity has to increase to keep the

flow rate constant.

5.2 Kaolinite

5.2.1 Breakthrough Curves for Different Sand Sizes

When 1ppt kaolinite was placed through the differing sand sizes at a constant flow rate of

1.25mL/min, one important observations was made (Section 4.2.2)- the larger the sand grain

size, the larger relative concentration peaks. The largest sand grain size produced the largest

relative concentrations peak indicating that there is less colloid filtration occurring inside the

column. If a value of 1 was reached for the relative concentration, this means that there is no

colloid filtered inside the column. Therefore the closer the relative concentration peak is to 1,

the less colloid mass is being retained in the column.

From the relative concentration peak the colloid filtration rate (kf) can be calculated (Section

3.4.1). The kf for each sand is listed below (Table 5-2). These values were consistent with kf

calculated by Akbour et al (2002) for kaolinite colloids. The kf decreases with decreasing sand

size. Therefore there are more colloids being filtered in the columns with a smaller sand size.

This makes sense, as there is a smaller pore volume in a column with finer sand grains. The

colloids are more likely to become trapped within this finer pore space, and will then fall out

of solution and collect with the column. Conversely for a larger sand grain size there are

larger pore spaces, so the likelihood of colloids becoming trapped decreases and there are less

colloid filtered.

Table 5-2 Colloid filtration rates calculated for each sand at a flow rate of 1.25mL/min.

Sand Size (mm) kf (min-1)

1.4-2 0.0148

0.5-0.71 0.0168

0.125-0.25 0.0175


39

5.2.2 Colloid Filtration Rate Dependence on Pore Water Velocity

The flow rate affects the concentration in the eluant from the column. As the flow rate is

increased, a higher concentration of kaolinite is recorded indirectly by the UV-VIS

spectrophotometer. This is because at higher flows more colloidal mass can be transported

through the column.

From the kf’s calculated from the experiment with variable flow rate (Section 4.2.3), it

indicates that kf is dependent on the pore water velocity. When the flow rate is increased (and

hence velocity), the relative concentration peak increases, and also the kf. The kf’s for each

velocity are listed below (Table 5-3). The data was then plotted to see what the relationship

was (Figure 5-1).

Table 5-3 Colloid filtration rates measured in 1.4-2mm sand for different pore water

velocities.

Pore Water Velocity (cm/min) Colloid Filtration Rate (min–1)

2.29 0.0435

3.57 0.0165

5.14 0.0076

This relationship needs to be studied more, as three points of data is not enough to establish a

proper relationship. But from equation 3.1 it shows that there is a relationship between the

two parameters. The negative sign shows that there is a negative relationship, which agrees

with the results.


40

0

0.01

0.02

0.03

0.04

0.05

0 2 4 6

Pore Water Velocity (cm/min)

Co

llo

id F

iltr

ati

on

Rate

(min

-1)

Figure 5-1 Relationship between colloid filtration rate and pore water velocity for 1.4-

2mm sand.

5.3 Barium Sulphate

5.3.1 Behaviour of BaSO4

Creating a BaSO4 suspension and pumping it into the column produced very erratic results

that were not reproducible (Section 4.3.1). Trends that did occur were that the peak in relative

concentration fluctuated in a range but did not reach some critical value. Initially the relative

concentration rose to some value then diminished to some constant range. This erratic

behaviour may be attributed to the BaSO4 colloids readily flocculating inside the column and

attaching to the sand grains. Colloids may be forming particle bridges between sand grains,

and these bridges may be broken by a mobile colloid. As these bridged colloids become

dislodged, large spikes in absorbance (and therefore concentration) will be recorded. Even

experiments conducted with no column produced these erratic results indicating that colloids

were attaching to the tubing in the apparatus as well. It can be deduced that colloids are being

deposited in the column from most results. The breakthrough curve for the 30ppm BaSO4

suspension in the 1.4-2mm sand column showed a peak range between 0.3-0.5 (Figure 4-6).

Since this peak does not reach anywhere near 1, this shows that colloids are being filtered out


41

of suspension and are being deposited inside the column. However due to the quality of the

data, it can only be analysed qualitatively. Therefore a kf can not be calculated.

This behaviour is consistent with past studies done by Sun and Skold (2001) and Aliaga et al

(1992). Both studies encountered the precipitate to be unstable in their experiments. Aliaga et

al (1992) even tried to stabilise the precipitate by adding a non-ionic surfactant and PAM.

5.3.2 Mixing Interface

The peaks in the experiments investigating the mixing interface between BaCl2 and Na2SO4

represent the formation of BaSO4 precipitate (Section 4.3.2). The peaks in absorbance show

that the UV-VIS spectrophotometer is detecting a more turbid solution than BaCl2 or Na2SO4

alone. This can only be explained by the fact that when the two solution meet inside the

column, they interact and form BaSO4, but only at the interface. This is called the zone of

mixing. This explains why the peak only occurs over a few pore volumes not over many pore

volumes. The BaSO4 is being formed at the interface of the two solutions and is then

transported out of the column by the flow. The sand grains promote some mixing of the two

solutions also, to help this reaction occur.

The baseline levels of absorbance are on the increase, which indicates that some proportion of

BaSO4 colloids that have accumulated inside the column. This is in agreement with the results

from creating BaSO4 outside the column and pumping it in, as it proves filtration is occurring

(Section 5.3.1). However the increasing levels suggest that the BaSO4 colloids may be

“bleeding” out of the column at some rate. Therefore this filtration process may be a

reversible reaction. Again this can only be analysed qualitatively. Therefore from the results a

kf can not be calculated.

The larger sand produced peaks over the largest pore volume in comparison with the medium

sand. This means that the BaSO4 colloids are being produced and dispersing throughout the

column more in the larger sand. The colloids are spreading out over a larger pore volume, and

are thus more mobile. This would indicate that its dispersion coefficient would be greater in

the larger sand. However this can not be computed due to the type the data. This observation

would be attributed to the greater pore volume within the 1.4-2mm sand column. A greater


42

pore volume allows for the colloids to disperse more easily than in a smaller pore volume.

The medium sand size column has more sand grains than the large column, so this allows for

more obstructions for a mobile colloid. The more sand grains there are, the harder the path is

to traverse through the column for the colloid. This is agreement with the results from the

NaCl tracer experiments (Section 4.1). It was found that the dispersion coefficient was larger

for sand grains of larger size.

No peaks in absorbance are observed in the finest sand column. This is because BaSO4 is

being produced in the column but a major proportion is not leaving the column. It can be

deduced that BaSO4 is being produced in the column by the increasing trend of baseline

absorbance after each interchange. Again we can see these “bleeding” phenomena, which is

evident by the increasing baseline levels.


43

6 Conclusion

From the experiments performed in this study the following was found-

• the colloid filtration rate (kf) is higher for smaller sand sizes. This means that more

colloids will be filtered out in the smaller sand.

• kf is dependent on pore water velocity (vp). The larger the vp, the lower the kf will be,

as a higher velocity will flush out more colloids.

• the dispersivity (λ) and the dispersion coefficient (D) are greater for larger sands. This

is because in larger sand there is a larger pore volume and less grain surface area for

the colloids to attach to. Therefore the colloids will disperse more throughout the sand.

• when a BaSO4 suspension is produced outside the column, the colloids readily

flocculate and create particle bridging between the sand grains. These bridges may be

broken by the displacement by mobile colloids. From these erratic results it is difficult

to determine a kf, but qualitatively it can be inferred that the sand filters out a

proportion of the BaSO4 colloids.

• injecting Na2SO4 into BaCl2 column environment was used to simulate ASR, where

BaSO4 was produced in the zone of mixing. The zone of mixing is greatest for larger

sands, as there is greater dispersion. However in the finest sand it appeared that no

BaSO4 was produced in the eluant, as it was accumulating inside the column. In all the

sands there appeared to be some “bleeding” of BaSO4, which indicates that colloids

were being released at some rate from the sands. Therefore the filtration of BaSO4

may be a reversible reaction.

This study gave introductory knowledge of the mobility of BaSO4 as a colloid. This is a

complex issue though as the behaviour of BaSO4 was very unstable. Most of the

conclusions about BaSO4 could only be done qualitatively not quantitatively. Therefore it

is recommended that further work be conducted to describe the behaviour quantitatively,

which was done for kaolinite. Work should be done to understand why BaSO4 is so

unstable. Possible factors responsible could be the effect of ionic strength or point of zero

charge. Experiments involving BaSO4 suspensions with differing ionic strength and pH

above and below the point of zero charge should be conducted. This should give an idea

why the precipitate is unstable.


44

Creating the mixing interface is the best experiment to represent what’s happening in the

aquifer. Na2SO4 represents sulphate-amended water being pumped into the aquifer.

BaSO4 is formed in the zone of mixing and is then filtered out by the porous media. It

then continues to accumulate in the column. Therefore with the formation of BaSO4 this

reduces the barium concentration in the water. Results show that the porous media

“bleeds” out BaSO4. This indicates that the filtration process is reversible as BaSO4 is

adsorbing to the sand grains but is then released after some time. More research should be

conducted to understand if this is indeed a reversible reaction. The mixing interface could

be investigated more by injecting Na2SO4 into a BaCl2 column environment and reversing

the flow to promote mixing. This would be representative of what occurs in the aquifer,

because more dispersion and mixing is expected in the field.

Kaolinite colloids proved a stable colloid for comparison. Therefore further research to

determines the kf for BaSO4 can be compared with kaolinite.

kf is dependant on pore water velocity, so this has implications for ASR. During ASR

velocities are greater near the well and decrease proportionately to the radius squared

from the well. Therefore near the well the filtration of colloids will be smaller, so the

colloids should form and disperse away from the well.

Another recommendation is to conduct experiments in soils characteristic to Perth’s

geology because the sands that were chosen were based on size only.


45

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