Phosphate removal from aqueous solutions using kaolin
obtained from Linthipe, Malawi.
Kamiyango M.W.1, Masamba W.R.L., 1, 2, *and Sajidu S.M.I1.
1
2
Chemistry Department, Chancellor College, University of Malawi, P.O. Box 280,
Zomba, Malawi
Current address: Harry Oppenheimer Okavango Research Centre, P/Bag 285, Maun,
Botswana
Abstract
Earlier work on stream water and effluent from wastewater treatment plants in Blantyre,
Malawi, has revealed high phosphate levels ranging from 0.63 to 5.50 mg/L. These
phosphate levels would stimulate excessive growth of plants and toxic cyanobacteria in
stagnant receiving water bodies hence posing a threat to aquatic life and water quality.
Phosphate removal by kaolin obtained from Linthipe, Malawi (Linthipe clay), was
investigated as a function of pH, contact time, clay dosage, competing ions and initial
phosphate concentration by means of jar tests. Phosphate uptake was pH dependent with
adsorption mechanisms on clay and iron oxide surfaces dominant generally below pH 7
and precipitation by calcium ions dominant above pH 7. Maximum phosphate removal
occurred at high pH values of 11.22 (97.1 %) and 10.96 (100 %) for raw and treated clay
respectively. Acid treated clay indicated higher phosphate removal efficiency than raw
clay at the normal working pH of 9.0 ± 0.2. This is attributed to release of extra calcium
ions from CaO that is present in the clay during acid treatment. At the maximum clay
dosage of 80 g/L, percent phosphate removal was 69.7 ± 0.1% and 98.5 ± 0.0577% for
raw and treated clay respectively. This indicates requirement for high clay dosages to
effect phosphate removal from the aqueous solutions. Studies on the effects of competing
ions indicate that phosphate uptake decreased in the presence of CO32-/HCO3-, was little
affected by SO42-, and increased in the presence of Ca2+, F-, and Mg2+ ions. Furthermore,
phosphate uptake approached equilibrium slowly at 20 C and was faster at 40 C. Use of
*
Corresponding author: Email: wmasambamba@orc.ub.bw
1
Linthipe clay in wastewater treatment would require pretreatment with an acid and pH
adjustment of the water to higher values.
Key words: Adsorption; clay; hydroxyapatite; phosphate; pHpzc; precipitation
Introduction
Phosphorus is one of the primary nutrients that cause detrimental eutrophication in
aquatic environments. A continuing elevated level of phosphorus in stagnant water
systems stimulates growth of aquatic plants and toxic cyanobacteria. The common forms
of phosphorus in wastewater are orthophosphates, polyphosphates (polymers of
phosphoric acid), and organically bound phosphates with orthophosphate being the most
abundant phosphorus species (Hammer and Hammer Jr., 2005). Large quantities of
phosphates are used in many industrial applications with fertilizers being the most
important. Other applications include domestic use of detergents containing phosphate
formulations as well as excessive fertilizer application in the agricultural sector.
Phosphorus content of effluents or streams that discharge directly into lakes and dams is
regulated by national and international water standard authorities with maximum limits
ranging from 0.1 to 2.0 mg/L as P, with many established at 1.0 mg/L (Hammer and
Hammer Jr., 2005). A study on water and wastewater quality in Blantyre, Malawi,
revealed phosphate concentrations ranging from 0.63 to 5.50 mg/L in some streams and
effluent from some wastewater treatment plants (Sajidu et al., 2007). This indicates need
for phosphate removal during wastewater treatment in an effort to avoid eutrophication in
stagnant water bodies onto which the effluent is discharged. Conventional methods for
phosphorus removal from wastewater include biological uptake and chemical
2
precipitation using iron and aluminum salts. Despite biological uptake and chemical
precipitation being established methods, problems are encountered owing to poor
operation stability or high economic costs. There has been growing interest in search for
locally available materials that can be used for phosphate removal during wastewater
treatment. A lot of studies have been done on phosphate removal using low cost materials
that include blast furnace slags (Johansson and Gustafsson, 2000; Kostura et al., 2005;),
alunite (Ozacar, 2006), calcite (Karageorgiou et al., 2007), iron oxide tailings (Zeng et
al., 2004), flyash (Can and Yildiz, 2006; Chen et al., 2007), bauxsol (Arkhurst et al.,
2006), red mud (Pradhan et al., 1998; Huang et al., 2008) and sludge from Fuller’s earth
production (Moon et al., 2007) just to mention a few. Phosphate is removed from
aqueous solutions via adsorption on iron and aluminum (hydr)oxides and precipitation by
Ca2+, Al3+ and Fe3+ ions by most of the studied materials. Goethite is the most studied of
the iron oxides owing to its role in nutrient adsorption and desorption by soils. Phosphate
adsorption on goethite is attributed to ligand exchange reaction on the adsorbent surface,
in which phosphate is exchanged with surface structural hydroxyl groups (Russell et al.,
1974). Phosphate adsorption on goethite proceed mainly via formation of innersphere
bidentate complexes, (FeO)2PO2(H) from pH 3.5 to 8.0 with abundance of a monodentate
complex, FeOPO3(H)1,2, increasing with increasing pH and decreasing phosphate surface
coverage (Zhong et al., 2007). In their study on phosphate adsorption on hematite,
Elzinga and Sparks (2007), reported surface complexes bound in a monodentate
(bridging) fashion as well as those coordinated as monodentate mononuclear complexes
to the hematite surface. Anion adsorption on aluminum oxide surfaces proceeds through
3
formation of ion-pairs with positively charged surface sites or by ligand exchange with
surface hydroxyls (Goldberg et al., 1996).
Reactions between phosphate species and iron and aluminum salts during chemical
precipitation are complex and not fully understood, however the primary reaction appears
to be the combining of orthophosphate with the metal cation (Hammer and Hammer Jr.,
2005). Investigations on phosphate removal using calcium containing materials such as
fly ash and calcite have shown that phosphate precipitates with calcium as apatites
especially hydroxyapatites
{Ca5(PO4)3.OH} (Johansson
and
Gustafsson,
2000;
Karageorgiou et al., 2007; Moon et al., 2007).
This paper reports studies on orthophosphate species removal from aqueous solutions by
clay obtained from Linthipe, Malawi, investigating key process parameters.
Materials and methods
Preparation and characterization of clay samples
Clay samples used in this study were obtained from kaolin deposits located at Linthipe in
Dedza district Malawi. The clay samples have chemical composition given in Table 1 as
determined by the Geological Survey Department of Malawi (GSoM). Prior to use, the
samples were thoroughly sun-dried and ground by a traditional pestle and mortar
followed by sieving through a 60-mesh size sieve. Clay samples prepared in this way are
referred to as ‘raw clay’. Acid treatment of the clay was achieved by addition of 100 mL
hydrochloric acid (0.3 mol.dm-3) to a 200g clay sample, ensuring thorough mixing in the
process. This was followed by drying of the wet clay in a sample-drying oven at 50 C
for 7 hours, crushing of the clay blocks, and sieving through a 60-mesh size sieve.
4
Table 1 Chemical composition of Linthipe clay
Parameter
Composition (wt.%)
SiO2
Al2O3
Fe2O3
CaO
MgO
K2O + Na2O
46.7
33.8
2.0
1.1
0.26
0.28
Potentiometric titration
Suspensions of 6.0 g sample/L were prepared for both raw and treated clay. For the
acidimetric titration, 25 mL of the clay suspension was titrated with 0.02359 mol.dm-3
nitric acid stepwise from 0 to 20 mL at increments of 0.1 mL using a Metrohm 775
Dosimat titrator equipped with a Metrohm 806 exchange unit. The pH reading after
equilibration of each step was recorded using Ross Sure Flow combination electrode
Orion 8172 suitable for pH determination in suspensions connected a Mettler-Toledo
SevenMulti meter. During the titration the flask containing the suspension was covered
with a parafilm to minimize formation of carbonates at pH greater than 6. An alkalimetric
titration was done following the same procedure using 0.020 mol.dm-3 NaOH. The
volume, Ve, corresponding to total proton monolayer coverage (which in turn was used to
calculate  SOH tot ), was calculated from the pH versus Vt (volume of acid added) plot.
The total concentration of surface charge, Q, at every titration point and the mean
number of protons per surface site (Z) were also calculated. Plots of apparent pKa
(pKaapp) values against Z for both the acidimetric and alkalimetric titrations were made
for determination of the intrinsic pKas, pKsa1intr and pKsa2intr. The points of zero charge for
the clay were calculated as; pHpzc = 0.5(pKsa1intr + pKsa2intr)
5
Standard solutions
A standard phosphate stock solution (1000 mg/L) was prepared by dissolving appropriate
amount of analytical grade anhydrous KH2PO4 (Glassworld, SA) in distilled water.
Intermediate standard solutions (100 mg/L) were prepared by diluting the stock solution
with distilled water. The intermediate standard solutions were used in obtaining desired
phosphate concentrations. Similarly, carbonate and bicarbonate standard solutions were
prepared from Na2CO3 (Glassworld, SA) and NaHCO3 (BDH) salts respectively.
Fluoride, sulphate, calcium (Ca2+), and magnesium (Mg2+) standard solutions were
prepared by dissolving NaF (SAARCHEM, SA), anhydrous Na2SO4 (ACE, SA), CaCl2
(SAARCHEM, SA), and MgCl2 (Glassworld, SA) salts, respectively, in distilled water.
Phosphate removal experiments
Defined masses of clay (either raw or treated) were weighed into centrifuge tubes to
which was added a defined volume of either 0.1 mol.dm-3 NaNO3 or distilled water
followed by a defined volume of 100 mg/L standard phosphate solution to make a total of
10 mL. Depending on parameters being studied, defined volumes of 100 mg/L standard
solutions of species of interest were added before introduction of the phosphate solution
with adjustments made to added volumes in order to obtain a total volume of 10 mL. An
appropriate amount of nitric acid (0.02 mol.dm-3) or sodium hydroxide (0.02 mol.dm-3)
was added to the clay suspension to adjust pH to a certain value. The mixture was shaken
for a defined period of time using a flask shaker then centrifuged at 2500 rpm using a
Hettich Roxita/AP centrifuge. The supernatant (2 to 3 mL) was passed through a Titan,
0.45 m, PES, membrane filter to remove solid particles. A component suppression ion
6
chromatography system composed of a Dionex CDM-1 conductivity detector, an Ionpac
AS14 anion exchange column and Data Apex Clarity chromatography software, was used
to determine phosphate concentration of the filtered supernatants.
Ca2+ and Fe3+ ion concentration in supernatants was determined using a Buck scientific
atomic absorption spectrophotometer (AAS) in accordance with standard methods
(APHA, 1985).
Results and discussion
Potentiometric titration of raw and treated clay
Linthipe clay is highly alkaline as indicated by initial pH values higher than 9 shown in
Figures 1 and 3 as well as 2 and 4 for both raw and treated clay. Equilibrium volumes
obtained from Figures. 1 and 3, were used to calculate total concentration of acid base
active sites,  SOH tot , as 4.188 x 10-3 mol.dm-3 (6.980 x 10-4 mol/g) and 3.517 x 103
mol.dm-3 (5.862 x 10-4 mol/g) for raw and treated clay respectively. Raw Linthipe clay
indicated buffering properties in the pH ranges of 5 to 6.5 and 7 to 8. Buffering properties
in the pH range of 5 to 6 could be attributed to deprotonation of water-soluble carboxylic
groups of humic and fulvic acids with pKa values in the ranges of 4 to 6. Between pH 6
and 6.5, buffering properties could be due to presence maleic and carbonic acids with
pKa values of 6.22 and 6.35 respectively. Buffering properties indicated between pH 7
and 8 could be attributed to presence of soluble inorganic phosphates considering that the
pK2 value for phosphoric acid is 7.2. Acid treatment did not alter buffering properties of
the clay as indicated by the same buffering regions (Figures 1 and 2). The point of zero
7
charge (pHpzc) for the raw clay was calculated to be 9.105 (Figure 5), and 9.041 for
treated clay (Figure 6).
Equilibrium pH against volume of acid added on
raw clay
9
9x
28
-1.
5
8
.09
+9
pH
y=
7
3
y = -0.0543x + 3.6168
1
0
5
10
0.02358M HNO3 (mL)
15
20
Figure 1 Acidimetric titration of raw clay
Equilibrium pH against volume of acid added
on treated clay
9
4x
24
-1.
5
66
.85
+8
pH
y=
7
3
y = -0.053x + 3.633
1
0
5
10
0.02358M HNO3 (mL)
15
Figure 2 Acidimetric titration of treated clay
8
20
Equilibrium pH against volume of base added on
raw clay
13
11
pH
9
7
5
3
1
0
5
10
15
20
25
0.020M NaOH (mL)
Figure 3 Alkalimetric titration of raw clay
Equilibrium pH against volume of base added
on treated clay
13
11
pH
9
7
5
3
1
0
5
10
15
20
0.020M NaOH (mL)
Figure 4 Alkalimetric titration of treated clay
9
25
pKa against mean number of protons bound
per surface site (for raw clay)
14
y = 4.2967x + 10.626
12
PKa
10
8
y = 0.09x + 7.5837
6
4
2
0
0
0.1
0.2
0.3
0.4
0.5
0.6
Z
Figure 5 Apparent acidity constants (pKa app.) calculated from figures 1 and 3 plotted against Z for
raw clay
pKa against mean number of protons bound per
surface site ( for treated clay)
14
y = 1.7197x + 10.654
12
PKa
10
8
y = 0.5449x + 7.4274
6
4
2
0
0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
Z
Figure 6 Apparent acidity constants (pKa app.) calculated from figures 2 and 4 plotted against Z for
treated clay
10
Effect of suspension pH
Variation of percent phosphate removal with initial suspension pH is shown in Figure 7.
Phosphate removal increased sharply with pH reaching a maximum at around pH 3 for
raw clay and pH 5 for treated clay followed by a sharp decrease to around pH 7 in both
cases. A valley was formed from pH 7 to 8 for raw clay followed by a sharp increase in
removal leveling off at pH 11. For the treated clay, phosphate removal increased sharply
from around pH 7 and leveled off at around pH 9. Treated clay showed higher phosphate
removal efficiency than raw clay above around pH 5.
% Phosphate removed against pH
% Phosphate removed
100
80
60
40
20
0
1
3
5
7
9
11
13
pH
raw clay
treated clay
Figure 7 Effect of clay suspension pH on phosphate removal. Reaction conditions were, 0.1M NaNO 3
background electrolyte, 40 g/L clay dosage, 10 mg/L initial phosphate concentration and 90 minutes
run time.
Higher phosphate removal at around pH 3 and 5 for raw and treated clay respectively,
suggest adsorptive mechanisms involving active sites on the clay surface as well as those
on metal oxides present in the clay mainly iron oxide. The proposed sites of anion
adsorption on layer silicates are positively charged aluminol (≡AlOH2+) functional groups
11
at exposed crystal edges (Manning and Goldberg, 1996). Since the aluminol group,
≡AlOH, is considered to be the main active site for complexation, phosphate adsorption
mechanisms on aluminum hydroxides can be inferred on layered silicates. The Visual
Minteq program was used to calculate fractional composition of phosphate species as a
function of pH given in Figure 8.
Fraction of species
1
0.8
0.6
0.4
0.2
0
0
2
4
6
8
10
12
14
pH
H2PO4-
H3PO4
HPO4-2
PO4-3
Figure 8 Dominant phosphate species in solution as a function of pH.
Between pH 2 and 7, H2PO4-, is the dominant phosphate species implying its use in
describing adsorption mechanisms in this pH region. Higher phosphate removal between
pH 3 and 5 should be guided by inner sphere complexation of phosphate on the clay
surface via ligand exchange reaction mechanisms. Equations 1 and 2 show the ligand
exchange reaction mechanism involving the aluminol functional group.
AlOH + H2PO4-
AlH2PO4 + OH-
12
(1)
Al
2
HPO4 + H2O + OH-
AlOH + H2PO4Al
(2)
The ion exchange mechanisms involves the release of hydroxyl ions into solution as such
phosphate adsorption increased with decreasing pH starting from around pH 7. The
decrease in phosphate adsorption with decreasing pH starting from around pH 3 for raw
clay and pH 5 for treated clay can be attributed to dissolution of the kaolinite surface
which releases soluble Al3+ below pH 4.5 (Weiland and Stumm, 1992), as well as
preferential adsorption of the phosphate species H2PO4- that dominates in solution
between pH 2 and 7 (Figure 8). The role of iron oxides in phosphate removal between pH
3 and 5 cannot be ignored considering that they form 2.0 % (wt.) of the clay (Table 1)
and that a higher clay dosage was used. Phosphate adsorption on the iron oxide surfaces
could be through the formation of protonated or unprotonated monodentate and bidentate
complexes (Kubicki et al., 2007).
The increase in phosphate uptake above pH 7 for both raw and treated clay, suggest
phosphate precipitation by metal ions as the main predominant uptake mechanism. The
clay suspensions contained significant amounts of calcium ions at different suspension
pH (Table 2), raising a possibility of phosphate precipitation as calcium phosphates.
13
Table 2 Ca2+ concentration in solution at different pH and saturation indices of
hydroxyapatite (HAP) and tricalcium phosphate (TCP).
Saturation index (SI)
SAMPLE
Equilibrium pH
Ca2+ (mg/L)
HAP
TCP
Raw clay
6.241
440.9
4.634
-0.9730
7.239
130.2
7.723
0.598
7.928
53.48
8.958
1.092
8.952
12.18
9.641
1.101
9.819
6.283
7.513
-0.519
10.758
4.316
1.386
-4.575
8.403
157.7
12.88
3.237
Treated clay
The Visual Minteq program was used to calculate saturation index (SI) values of possible
calcium phosphate minerals. Hydroxyapatite (HAP) and tricalcium phosphate (TCP)
minerals had more positive saturation index (SI) values as compared to the other possible
calcium phosphate minerals across the pH range studied as such their SI values are given
in Table 2. It is evident from the calculated SI values that all the solutions were over
saturated with respect to HAP whereas some were under saturated with respect to TCP. It
follows from the calculated SI values that the increase in phosphate uptake with
increasing pH starting from around pH 7 was due to phosphate precipitation as calcium
phosphates mainly HAP as given by equation 3.
5Ca2+ + 3PO43- + OH-
Ca5(PO4)3OH
(3)
Acid treatment of raw clay resulted in release of more Ca2+ ions (Table 2) hence higher
phosphate removal efficiency of treated clay than raw clay with increasing pH.
14
Effect of clay dosage
Using an initial phosphate concentration of 10 mg/L and a run time of 90 minutes, the
effect of clay dosage on phosphate uptake was studied and the results are shown in Figure
9.
% phosphate removed against clay dosage
% Phosphate removed
100
80
60
40
20
0
0
20
40
60
80
Dosage (g/L)
Raw clay
Treated Clay
Figure 9 Effect of raw and treated clay dosage on phosphate removal.
Phosphate removal increased with clay dosage reaching 98.5% and 69.7% removal for
treated and raw clay respectively at a dosage of 80 g/L. It is evident from Figure 9 that
the clay’s ability to remove phosphate increased with acid treatment as a result of an
increase in free calcium ions.
Effect of contact time
Effect of contact time on phosphate uptake was studied for raw clay at a fixed dosage of
40 g/L and 10 mg/L initial phosphate concentration at temperatures of 20 and 40 C.
Phosphate uptake approached equilibrium faster at 40 C and slowly at 20 C but
eventually leveled off after 270 minutes at both temperatures (Figure 10).
15
Phosphate uptake (mg/g)
Phosphate uptake against time
0.25
0.2
0.15
0.1
0.05
0
0
50
100
150
200
250
300
350
Time (minutes)
293K
313K
Figure 10 Effect of contact time on phosphate removal at 20 and 40 C.
First order plot
0
log (Qe-Qt)
-0.5 0
5000
10000
15000
20000
y = -9E-05x - 0.7061
R2 = 0.7837
-1
-1.5
-2
-2.5
y = -1E-04x - 1.2093
R2 = 0.9099
-3
-3.5
293K
Time (s)
Linear (293K)
313K
Linear (313K)
Figure 11 Kinetics data fitted to Largergren pseudo first order model.
Second order plot
t/Qt (g.s/mg)
120000
y = 4.8416x + 15982
R2 = 0.9635
100000
80000
60000
y = 5.3704x + 2642.3
R2 = 0.999
40000
20000
0
0
5000
10000
15000
20000
Time (s)
293K
313K
Linear (293K)
Linear (313K)
Figure 12 Kinetics data fitted to Largergren pseudo second order model
16
In their study on effects of solution conditions on phosphate precipitation, Song et al.
(2002), noted that temperature affects not only precipitation of HAP, but also the
association/dissociation reactions of the related species and the thermodynamic solubility
of HAP. On overall SI values for HAP increase with temperature hence the faster
approach towards equilibrium. The kinetic mechanism of phosphate uptake was
evaluated with two different models: the Largergren pseudo first order and pseudo second
order models (Huang et al., 2008). The pseudo first order equation takes the form
log(qe – qt) = log qe – k1t/2.303
(4)
where qt is the amount of phosphate removed at time t (mg/g), qe is the amount of
phosphate bound at equilibrium (mg/g), and k1 is the equilibrium rate constant for pseudo
first-order kinetics (s-1). The pseudo second-order equation is expressed as
t/qt = 1/k2qe2 + t/qt
(5)
where k2 is the equilibrium rate constant for the second-order kinetics (g mg-1 s-1). Data
obtained at both temperatures fitted better to the pseudo second-order model than to the
pseudo first-order model as indicated by R2 values shown in Figures 11 and 12.
Equilibrium rate constants at 20 and 40 C were calculated as 1.467 x 10-3 g/mg.s and
1.092 x 10-2 g/mg.s respectively. Activation energy for the reaction (Ea), was calculated
using the relation;
ln(k2(a)/k2(b)) = Ea/R (1/Tb – 1/Ta)
17
(6)
where k2(a) and k2(b) are second order equilibrium constants at 20 and 40 C respectively,
Ta and Tb are temperatures in Kelvin and R is the universal gas constant (8.314 J/K.mol)
and it was found to be 76.56 kJ/mol.
Effect of initial phosphate concentration
The effect of initial phosphate concentration on phosphate uptake (Figure 13) shows
increasing uptake with increasing initial concentration followed by leveling off at
maximum uptake.
Phosphate uptake against initial concentration
Phosphate uptake (mg/g)
0.8
0.6
0.4
0.2
0
0
10
20
30
40
50
60
70
Initial phosphate concentration (m g/L)
raw clay treated clay
Figure 13 Effect of initial phosphate concentration on phosphate removal. Reaction conditions were,
40g/L clay dosage, and 90 minutes run time.
The saturation index for precipitation of hydroxyapatite is given by the equation,
SI = 5log(CCa2+) + 3 log(CPO43-) + log(COH-) + log(f25f33f1) – logKsp
(7)
where Cx is the concentration of the x ion and fi is the activity coefficient of the i-valent
ion (Song et al., 2002). It is evident from the equation that holding other parameters
18
constant, the saturation index of the solution will increase with increasing initial
phosphate concentration. Phosphate precipitation is also guided by available calcium ions
hence attainment of maximum uptake with increasing initial phosphate concentration.
Faster attainment of maximum uptake by raw clay follows from lower amounts of
calcium ions as compared to treated clay.
Effect of fluoride, calcium, magnesium, sulphate and carbonate/bicarbonate ions on
phosphate uptake
Figures 14 and 15, show effect of fluoride, calcium, magnesium, carbonate/bicarbonate
and sulphate ions on phosphate uptake by raw and treated clay respectively.
Phosphate uptake against concentration of other ions
(raw clay)
Phosphate uptake (mg/g)
0.25
0.2
0.15
0.1
0.05
0
0
100
Fluoride
Calcium
Magnesium
200
300
400
500
Concentration (m g/L)
Carbonate/Bicarbonate
Sulphate
Figure 14 Effect of fluoride, calcium, magnesium, sulphate, and carbonate/bicarbonate ion
concentration on phosphate removal. Raw clay dosage of 40 g/L and 90 minutes run time conditions
were applied.
19
Phosphate uptake (mg/g)
Phosphate uptake against concentration of other
ions (treated clay)
0.25
0.2
0.15
0.1
0.05
0
0
100
Fluoride
Magnesium
Sulphate
200
300
400
500
Concentration (m g/L)
Calcium
Carbonate/Bicarnonate
Figure 15 Effect of fluoride, calcium, magnesium, sulphate, and carbonate/bicarbonate ion
concentration on phosphate removal. Treated clay dosage of 40 g/L and 90 minutes run time
conditions were applied.
Phosphate uptake increased with increasing fluoride, calcium and magnesium ion
concentration, it was little affected by sulphate ions and decreased with increasing
carbonate/bicarbonate concentration for both raw and treated clay. The increase in
phosphate uptake with increasing calcium concentration was more pronounced for raw
clay than treated clay probably due to release of more calcium ions during acid treatment
such that additional calcium ions had little effect on saturation of the solution. The
increase in phosphate uptake with increasing fluoride concentration can be explained by
over saturation of the solution with respect to fluorapatite, (Ca5(PO4)F), and hence
phosphate precipitation as given by equation 8.
5Ca2+ + 3PO43- + F-
Ca5(PO4)F
(8)
The increase in phosphate uptake with increasing magnesium ion concentration could be
as a result of phosphate precipitation as magnesium phosphate, Mg3(PO4)2, though the
20
reactions may be complex considering a wide range of possible mineral phases with
different thermodynamic stabilities that could precipitate. Calculation of SI values for
possible mineral phases in the presence of sulphate ions showed that all the solutions
were under saturated with respect to possible sulphate minerals. This explains the
observed little effect of sulphate ions on phosphate uptake. The decrease in phosphate
uptake with increasing carbonate concentration could be as a result of formation of
carbonate hydroxyapatite from the precipitated HAP. At lower carbonate concentrations,
both phosphate and hydroxyl groups in the HAP lattice are substituted by carbonate
resulting in release of phosphate and hydroxyl ions into solution (Barralet et al., 1997).
At higher carbonate concentrations, substitution of phosphate for carbonate dominates.
This implies that the observed decrease in phosphate uptake with increasing carbonate
concentration could have been guided by formation of carbonate hydroxyapatite from the
precipitated HAP.
Conclusion
Phosphate removal by Linthipe clay was pH dependent proceeding via adsorptive and
precipitation mechanisms at lower and higher pH values respectively. Equilibrium was
approached faster at 40 C than at 20 C but was attained after 270 minutes of the
reaction at both temperatures. The study has also shown that phosphate removal would be
enhanced by presence of higher amounts of calcium, fluoride and magnesium ions but
would not be affected by sulphate ions whereas carbonate/bicarbonate ions would
decrease phosphate removal. In an overall assessment, Linthipe clay has low phosphate
21
removal efficiency as indicated by requirement for high dosages to effect removal hence
the need to study other locally found materials.
Acknowledgement
The authors acknowledge support from the International Programme in Chemical
Sciences (IPICS) through the University of Uppsala, Sweden (project code MAW:02) for
a grant that provides Msc scholarship for M. Kamiyango to study at the University of
Malawi.
References
Akhurst, D.J., Jones, G.B., Clark, M., McConchie, D., 2006. Phosphate removal from
aqueous solutions using neutralized bauxite refinery residues (Bauxsol). Environmental
chemistry 3: 65-74
American Public Health Association (APHA), 1985. Standard methods of the
examination of water and wastewater, 16th Edition. APHA, AWWA, and WPCF, New
York.
Barralet, J., Best, S., Bonfield, W., 1997. Carbonate substitution in precipitated
hydroxyapatite: An investigation into the effects of reaction temperature and bicarbonate
ion concentration. Journal of biomedical materials research part A, 41(1): 79-86
Can, M.Y., Yildiz, E., 2006. Phosphate removal from water by fly ash: Factorial
experimental design. Journal of hazardous materials B, 35: 165-170
Chen, J., Kong, H., Wu, D., Chen, X., Zhang, D., Sun, Z., 2007. Phosphate
immobilization from aqueous solution by fly ashes in relation to their composition.
Journal of hazardous materials B 139: 293-300
22
Elzinga, J.E., Sparks, D.L., 2007. Phosphate adsorption onto hematite: An in situ ATRFTIR investigation of the effects of pH and loading level on the mode phosphate surface
complexation. Journal of Colloid and Interface Science, 308: 53-70
Goldberg, S., Davis, J.A., Hem, J.D., 1996. The surface chemistry of aluminum oxides
and hydroxides. In: The environmental chemistry of aluminum, G. Sposito (Ed). Lewis
publishers, Florida, USA, pp 271-318
Hammer, M.J., Hammer, Jr., 2005. Water and wastewater technology, 4th Edition.
Prentice Hall of India, New Delhi.
Huang, W., Wang, S., Zhu, Z., Li, L., Yao, X., Rudolph, V., Haghseresht, F., in press.
Phosphate removal from wastewater using red mud. Journal of hazardous materials, doi:
10.1016/j.jhazmat.2008.01.061
Johansson, L., Gustafsson, J.P., 2000. Phosphate removal using blast furnace slags and
opoka-mechanisms. Water research 34(1): 259-265
Karageorgiou, K., Paschalis, M., Anastassakis, G.N., 2007. Removal of phosphate
species from solution by adsorption onto calcite used as natural adsorbent. Journal of
hazardous materials A, 139: 447-452
Kostura, B., Kulveitova, H., Lesko, J., 2005. Blast furnace slags as sorbents of phosphate
from water solutions. Water research 39: 1795-1802
Kubicki, J.D., Kwon, K.D., Paul, K.W., Sparks, D.L., 2007. Surface complex structures
modelled with quantum chemical calculations: carbonate, phosphate, sulphate, arsenate
and arsenite. European journal of soil science 58: 932-944
Manning, B.A., Goldberg, S., 1996. Modelling arsenate competitive adsorption on
kaolinite, montmorillonite and illite. Clays and clay minerals, 44(5): 609-623
23
Moon, Y.H., Kim, J.G., Ahn, J.S., Lee, G.G., Moon, H., 2007. Phosphate removal using
sludge from fuller’s earth production. Journal of hazardous materials, 143: 41-48
Ozacar, M., 2006. Contact time optimization of two-stage batch adsorber design using
second-order kinetic model for the adsorption of phosphate onto alunite. Journal of
hazardous materials B 137: 218-225
Pradhan, J., Das, J., Das, S., Thakur, R.S., 1998. Adsorption of phosphate from aqueous
solution using activated red mud. Journal of colloid and interface science 204: 169-172
Russell, J.D., Parfitl, R.L., Fraser, A.R., Farmer, V.C., 1974. Surface structures of
gibbsite goethite and phosphated goethite. Nature 248: 220-221
Sajidu, S.M.I., Masamba. W.R.L., Henry, E.M.T., Kuyeli, S.M., 2007. Water quality
assessment in streams and wastewater treatment plants of Blantyre, Malawi. Physics and
chemistry of the earth, 32:1391-1398
Song, Y., Hahn, H.H., Hoffmann, E., 2002. Effects of solution conditions on the
precipitation of phosphate for recovery: A thermodynamic evaluation. Chemosphere, 48:
1029-1034
Wieland, E., Stumm, W., 1992. Dissolution kinetics of kaolinite in acidic aqueous
solutions at 25 C. Geochimica et Cosmochimica Acta, 56: 3339-3355
Zheng, L., Xiaomei, L., Liu, J., 2004. Adsorptive removal of phosphate from aqueous
solutions using iron oxide tailings. Water research, 38: 1381-1326
Zhong, B., Stanforth, R., Wu, S., Chen, J.P., 2007. Proton interaction in phosphate
adsorption onto goethite. Journal of Colloid and Interface Science 308: 40-48
24