Environmental Technology, Vol. 22. pp 1325-1335
© Selper Ltd, 2001
THE APPLICATION OF CALCIUM PHOSPHATE
PRECIPITATION CHEMISTRY TO PHOSPHORUS
RECOVERY: THE INFLUENCE OF ORGANIC LIGANDS
J. A. M. VAN DER HOUWEN* AND E. VALSAMI-JONES
Department of Mineralogy, The Natural History Museum, Cromwell Road, London SW7 5BD, UK
(Received 22 April 2001; Accepted 20 August 2001)
ABSTRACT
This paper describes current knowledge of phosphate precipitation chemistry in the context of phosphorus recovery from
wastewaters, and presents experimental results on the effect of organic species, as key potential inhibitors, to the
precipitation of calcium phosphate. The supersaturation required for precipitation at 25ºC, pH 7, 0.1 M ionic strength and
near-stoichiometric (for hydroxylapatite) calcium to phosphate molar ratio was determined under spontaneous precipitation
conditions. The experiments were carried out in air. The phase precipitating at the critical concentration was allowed to
grow under constant supersaturation. The influence of organic ligands on the precipitation was investigated using two small
molecular weight organic ligands, acetate and citrate, present at a concentration of 10 -3 M. The precipitate was studied using
X-ray diffraction and scanning electron microscopy. Good reproducibility of the experiments, which were carried out in
triplicate, was observed. The study assessed the supersaturation necessary for spontaneous precipitation of hydroxylapatite
to be 10.93, calculated using a solubility constant of log K=-57.74. The required supersaturation was not affected by the
presence of acetate. However, citrate was found to increase the supersaturation required for precipitation to 11.73. It is likely
that this increase is due to binding of citrate on the active growth sites of newly formed nuclei, thereby inhibiting
precipitation. All experiments showed formation of a single phase: micro-crystalline hydroxylapatite.
Keywords:
Citrate, spontaneous precipitation, hydroxylapatite, supersaturation.
INTRODUCTION
The formation of phosphates is important for a number
of natural, biological and engineering processes. A major
focus for phosphate research in the last two decades has been
the study of their precipitation as a mechanism of phosphorus
removal from wastewater [1,2]. Phosphorus removal could be
made more sustainable if combined with recovery. However,
recovery to date has not developed beyond sporadic isolated
trials. Examples include wastewater treatment plants in South
Africa [3], Germany [4], Japan [1, 5] and Italy [6]; animal
manure treatment in Australia [7] and Holland [8]; and
industrial waste streams in Korea [9] and Turkey [10]. In the
majority of these cases struvite is the precipitating phosphate,
but pilot/full-scale trials producing calcium phosphate are
also known from Australia [11], Holland [12] and Germany
[13]. Most of these trials reported that the essential
requirements for the process to proceed were chemically
induced high pH, and CO2 stripping from the effluent, prior
to entering the precipitation reactor. They have also reported
problems with fines and precipitates of non-recyclable quality.
The current consensus is that the process needs further
development to produce a universal reactor design, to lower
process costs and to improve product quality/recyclability.
Further work is necessary, in particular to better
understand the complex chemistry of calcium phosphates. It
is currently believed, that although the most stable calcium
phosphate phase is hydroxylapatite (HAP), there is a range of
other calcium phosphate phases, that may precipitate as a
precursor to HAP, namely dicalcium phosphate dihydrate
(DCPD, [CaHPO4*2H2O]), octacalcium phosphate (OCP,
[Ca4H(PO4)3*2.5H2O]),
tricalcium
phosphate
(TCP,
[Ca3(PO4)2]) or amorphous calcium phosphate (ACP). The
formation of alternative calcium phosphates is a kinetic issue,
i.e. the result of a phase precipitating faster, despite being less
saturated [14, 15, 16].
A number of kinetic studies have attempted to establish
the specific conditions and order of formation of the precursor
phases, but the results often vary from study to study. For
example, Zawacki et al. [17] observed a nucleation of calcium
deficient HAP on HAP seeds, while Barone et al. [18] reported
a precipitation of DCPD onto the HAP seed crystals. Frèche
and Heughebaert [19] and Heughebaert et al. [20] observed
the precipitation of OCP and DCPD onto seed crystals of OCP
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and DCPD, but no precipitation of HAP was shown.
It is often suggested that at pH>7 and high
supersaturation, the precursor phase is an amorphous calcium
phosphate (ACP). The ACP may dissolve again and form
HAP nuclei [21]. A three-stage formation of hydroxylapatite
starting with the formation of ACP, followed by OCP has also
been suggested [15]. Kibalczyc [14], under pH driftconditions, observed a transformation of one type of ACP into
another ACP, which then transformed into HAP. Similar work
[22] showed the same ACP(1) and ACP(2) formation but also
OCP as precursor phases to HAP; these phases were
determined using TEM and electron diffraction analyses on
the precipitate. Precipitation of non-stoichiometric apatites
has also been observed [17, 23].
A number of ions have been shown to act as inhibitors
to precipitating phases, by forming a surface complex on the
newly forming surfaces, and blocking further precipitation or
transformation to a more stable phase. The influence of
inhibitor ions, such as carbonate and magnesium, to calcium
phosphate precipitation has been discussed by many authors
[e.g. 24]. Co-precipitation of magnesium with the calcium
phosphates for example may induce firstly the formation of
ACP, which will transform into the more stable HAP. It is
possible for HAP to incorporate a small percentage of
magnesium into its structure, but this causes structural
changes and has an inhibitory effect to further HAP formation
[25]. Salimi et al. [26] studied the crystal growth of DCPD and
OCP in the presence of magnesium and carbonate using
constant composition and pH stating methods. They found
that the crystal growth of DCPD was not affected while the
OCP crystal growth rate decreased considerably. The presence
of CO2 may also affect precipitation of HAP, either by
blocking phosphate nucleation sites, or by inducing calcium
carbonate precipitation instead [e.g. 27]. Other dissolved
cations in solution may also block calcium phosphate
formation, or even cause HAP dissolution. Examples include
the dissolution of HAP to form the less soluble fluorapatite in
the presence of fluoride ions (28) or to form metal phosphates
in the presence of lead or cadmium [29]. There is currently a
particular need for research on calcium phosphate
precipitation in organic rich systems that simulate wastewater
environments. A small number of studies exist, which have
assessed the effect of organic ligands on the precipitation of
dicalcium phosphate dihydrate (DCPD) [30, 31], octacalcium
phosphate (OCP) [32] and the inhibition of hydroxylapatite
(HAP) precipitation [33]. All these studies were carried out as
seeded growth experiments (heterogeneous nucleation) and
all reported inhibition of precipitation kinetics by the presence
of the organic ligands. Adsorption of the ligands on active
growth sites of the seeding material was given as the cause for
this inhibition. The influence of citrate on calcium phosphate
precipitation onto OCP seeding material has also been studied
[34]. This study concluded that phospho-citrate complexes,
formed onto active growth sites, inhibited precipitation
kinetics. An earlier study into the influence of citrate on
calcium phosphate formation [35] suggested that calcium
citrate complexes inhibited transformation of ACP into a
more stable crystalline calcium phosphate.
Previous studies into spontaneous precipitation
proceeded by the mixing of calcium and phosphate solutions
at a desired supersaturation and leaving the nuclei enough
time to grow into crystals whilst keeping the pH constant [36].
A disadvantage of this method is that the solution
composition will change continuously (i.e. calcium and
phosphate concentrations fall), as soon as precipitation starts.
This, in turn, may result in variation in the nature
(mineralogy and/or crystallinity) of the precipitate.
Spontaneous precipitation has been criticised for giving nonreproducible results due to chance nucleation onto foreign
particles in the solution [37], however our study suggests that
by using carefully controlled systems (e.g. fresh solution
preparation, minimisation of potential nuclei by filtration) it is
possible to reproduce results; this was demonstrated by
triplicate experiments.
Previous research has shown that adsorption of organic
ligands to active sites of seeds, in heterogeneous nucleation
experiments, inhibits growth; it is now important to compare
and contrast this previous work with precipitation in the
absence of seeds, i.e. during well-controlled spontaneous
precipitation experiments. Investigation into un-seeded
precipitation is also important for the understanding of
problems such as the formation of fines, nucleating
spontaneously in industrial processes.
The specific aim of the present study was to elucidate
the chemical principles of calcium phosphate precipitation at
neutral pH, particularly in the presence of organic ligands.
The critical supersaturation for spontaneous precipitation of
calcium phosphate at pH 7 was determined, and the results
were compared with those of experiments carried out in the
presence of organic ligands. The small molecular weight
organic ligands studied were acetate and citrate, which
possess one and three functional groups (carboxylates)
respectively.
MATERIALS AND METHODS
Spontaneous Precipitation Calculations
For spontaneous precipitation to occur, the solution
needs to be supersaturated with respect to a mineral phase.
The degree of supersaturation is given by the following
equation (i).
SCa/P = log(IAP/Ksp)
(i)
Where IAP is the ionic activity product and Ksp is the
solubility constant for the calcium phosphate mineral [38].
To determine the supersaturation degree required for
precipitation it is necessary to state the ionic products for each
of the expected calcium phosphate minerals (equations ii-v):
Log IAPDCPD = log (Ca2+) + log (HPO42-)
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(ii)
Log Ksp
= -6.68 [39]
Log IAPTCP = 3 log (Ca2+) + 2 log (PO43-)
Log Ksp = -28.9 [40]
Log IAPOCP = 4 log (Ca2+) + 3 log (PO43-) + log (H+)
Log Ksp = -46.9 [41]
Log IAPHAP = 5 log (Ca2+) + 3 log (PO43-) + log (OH-)
Log Ksp = -57.74 [42]
(iii)
(iv)
(v)
(vi)
f=
activity
z = valency
I=
Ionic strength
A is 0.509 for water at 25C [38].
The experiments were carried out in air, therefore the
presence of CO2 was included in the modelling with
PHREEQC. Thermodynamic data used in PHREEQC are
given in Table 1.
Table 1.
Solubility product constants in the Na-Cl-PO4-CaCO2-H2O-system.
(Minteq database in PHREEQC [43]).
Thermodynamic dissociation products
Log k
Ca2+ + PO43- + H+ = CaHPO4
Ca2+ + PO43- = CaPO4Ca2+ + PO43- + 2H+ = CaH2PO4+
PO43- + H+ = HPO42PO43- + 2H+ = H2PO4PO43- + 3H+ = H3PO4
Ca2+ + CO32- + H+ = CaHCO3+
Ca2+ + CO32- = CaCO3
15.085
6.459
20.96
12.346
19.553
21.7
11.33
3.15
Organic acid equilibrium reaction
Acetate- + H+ = Hacetate
H+ + Citrate3- = CitrateH22H+ + Citrate3- = CitrateH23H+ + Citrate3- = CitrateH3
Ca2+ + Citrate3- = CaCitrate-
4.76
6.33
11.05
14.18
4.73
3.02
1.29
1.18
Solution Preparation
For the calculation of the free ion concentrations of the
lattice parameters and subsequently the supersaturation
degree the computer program PHREEQC (a standard
program, which allows calculation of saturation by taking all
ionic interactions in solution into consideration) was used
with the Minteq database [43]. The activity was calculated
using the extended Debye-Hückel equation proposed by
Davies [44] (vi):
Log f = -Az2[(I½/(+I½))-0.3I]
Ca2+ + Citrate3- + H+ = CaCitrateH
Ca2+ + Citrate3- + 2H+ = CaCitrateH2+
Ca2+ + Acetate- = CaAcetate+
All chemicals used were of Analar® purity or better.
Acetic acid was made up from glacial acetic acid solution and
citric acid was made up using citric acid salt. Stock solutions
of 0.1 M organic acid were prepared which were standardised
with 0.1 M NaOH. The NaOH solution was standardised with
potassium hydrogen phthalate (COOHC6H4COOK) solutions.
These titrations were carried out with a Mettler DL55
automatic titrator. Stock solutions of 0.1 M calcium and 0.06
M phosphate were made from calcium chloride and sodium
phosphate salts; their concentrations were confirmed by ICPAES analyses. Working solutions of 10-3 M organic acid were
made freshly for each experiment using the stock solutions
and adding NaCl as background electrolyte to ensure a final
ionic strength of 0.1 M.
Since the purpose of the experiments was to build up
supersaturation in solution to the point that spontaneous
precipitation occurs, the experiments started at a low
supersaturation of SCa/P= 3. Calculations with PHREEQC [43]
were again carried out, as described above, to predict the
calcium and phosphate concentrations in the working
solution necessary for this supersaturation. Two individual
titrant solutions, one of calcium and one of phosphate, were
made up, to be used subsequently for the incremental
increase of supersaturation. These titrants were made at
concentrations corresponding to the molar stoichiometric
ratio for hydroxylapatite (calcium to phosphate: 1.67±0.1, in
concentrations of 0.5 M or 0.05 M calcium and 0.3 M or 0.03 M
phosphate) from the stock solutions and added in equal
quantities, to maintain the stoichiometric ratio in solution.
Both titrants were made up in 0.1 M NaCl and were adjusted
to the same pH as the working solution. A pH of 7 in the
working solutions was controlled by addition of 0.1 M NaOH
by automatic titration and equilibrated for 24 hours in air. All
solutions used in these experiments were filtered through a
0.2m Millipore® filter to minimise the possibility of
introducing particles which could interfere with nucleation.
Precipitation Experiments
In these experiments an overhead propeller stirrer was
used at a speed of 180 rpm. The experiments were carried out
in a water bath at 25C. Experiments were carried out in
triplicate to ensure reproducibility.
To determine the saturation degree necessary for
precipitation, calcium (0.5 M) and phosphate (0.3 M) were
introduced stepwise. After each addition of calcium and
phosphate solutions, 30 minutes was given as an induction
period before the next addition. Two samples were taken at
each increment: the first sample was taken 1 minute after
addition, and the second sample was taken immediately
1327
before the next addition to confirm stability of the calcium
and phosphate concentrations during this induction period;
these were analysed after the experiment. During the
experiment, the onset of precipitation was recognised by online monitoring of pH. A continuous drop in pH indicated
precipitation. Once precipitation started, and in order to
maintain constant saturation conditions during precipitation,
for as long as necessary to produce enough precipitate for
identification, a method of maintaining constant composition,
as described previously in the literature [45], was used. The
principle of the constant composition method is that at the
critical supersaturation, where precipitation begins, the drop
in pH triggers the automatic addition (via a pH controlled
titrator) of calcium, phosphate and a pH adjusting solution
(base). This maintains calcium, phosphate and pH at a
constant level and allows the precipitate to grow. The
experimental set-up is given schematically in Figure 1.
The titrant concentrations were provided at the
stoichiometric ratio of hydroxylapatite (calcium to phosphate
1.67±0.1) for calcium and phosphate. Preliminary experiments
showed that the supersaturation conditions and pH remained
constant when the molar ratio of calcium: phosphate: base in
the titrants was 5:3:5; the concentrations were 0.05 M calcium,
0.03 M phosphate and 0.05 M NaOH. At the end-point of the
experiment the solution was rapidly filtered through a 0.2 m
syringe filter to retrieve the precipitate, which was
subsequently dried at 40C for identification.
Analyses
During the experiments 1.5 ml samples were taken and
filtered through a 0.2 m syringe filter to determine calcium
and phosphate concentration of the solution. An aliquot of 1
ml of the filtrate was diluted 10 times and acidified at 2%
HNO3. Calcium and phosphorus were measured using
Inductively Coupled Plasma Atomic Emission Spectrometry
(ICP-AES) with an ARL 3410. The precipitate was identified
using X-ray diffraction (Enraf Nonius). The retrieved solid
was further characterised using scanning electron microscopy
(SEM; Philips XL30).
RESULTS AND DISCUSSION
Table 2 lists the initial concentrations, which show that
the experiments started at the stoichiometric ratio for HAP of
1.67±0.08. The experiments were controlled at pH 7.01 with an
accuracy of ±0.05 and were found to be reproducible.
Two representative graphs, showing experimental
progress, are presented in Figures 2a and b, in order to
facilitate understanding of how the experiments were carried
out. The two graphs display one of the (triplicate) sets of data,
for the control experiment, and one for the experiment carried
out in the presence of citrate. The graphs show the stepwise
addition of calcium and phosphate until precipitation takes
place, at which point concentration remains constant via
continuous automatic titrant addition. As demonstrated in the
graphs, the interval between stepwise addition of calcium and
phosphate concentrations is fixed, but the titrant volume
added could vary from one experiment to the other. This had
no effect on the point at which precipitation occurred,
which is the critical point (vertical line) shown in the graphs.
1328
Figure 1.
Schematic experimental set-up for precipitation under constant composition conditions.
Table 2. Starting experimental concentrations for calcium and phosphate.
Experiment
Calcium
(10-4 mol l-1)
Phosphate
(10-4 mol l-1)
Molar ratio
(Ca:PO4)
Control
(no organic ligand)
Acetate
Citrate
3.37
1.96
1.714
3.48
3.29
1.99
2.00
1.746
1.644
(a)
(b)
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Figure 2. Determination of the critical point of precipitation (vertical line) on a concentration against time plot, for two different
experiments: a. control (no organic ligand); b. citrate.  represents calcium and  represents phosphate.
In other words, precipitation only occurred when the critical
calcium concentration in the presence of citrate is also
concentration was reached, regardless of the rate of titrant
significantly higher than in the other two experiments. In
addition. The supersaturation conditions were maintained
quantitative terms, the increase in free calcium concentration
until a visual inspection suggested there was enough
in the presence of citrate compared to the control is 18%. The
precipitate for identification; the duration of this second stage
calculated free phosphate concentration given in Figure 3b
was therefore not time-dependent. The results are plotted here
also shows a higher free phosphate concentration in the
as a function of time, only as a means of demonstrating
presence of citrate. The percentage increase in comparison
experimental progress.
with the control is 43%. This is higher than the increase in free
The calcium to phosphate molar ratio in solution
calcium, but it is to be expected because the phosphate
slightly increased during the stepwise addition process but
concentration is not influenced by the presence of citrate (no
was found not to increase over 1.8 (on average). Overall, the
complex formation between citrate and phosphate).
calcium and phosphate concentrations remain constant
By excluding aqueous complexation from the potential
during precipitation. A summary of experimental results is
mechanisms causing the observed increase in saturation index
shown in Table 3, where the average critical concentration,
necessary for precipitation to occur in the presence of citrate,
upon which precipitation occurs in each experiment, is given.
the only other alternative mechanism we are left with is
These concentrations are the average of three experiments;
inhibition of nucleation and/or growth. It is likely that, due to
standard deviations are also shown and demonstrate the good
citrate binding onto active growth sites of the newly formed
reproducibility of the experiments.
nuclei, growth of the nuclei is inhibited until there is enough
These concentrations are used in the calculation of
critical concentration of phosphate and calcium in solution for
supersaturation using the computer program PHREEQC [43].
precipitation to occur regardless. An interaction between
The degree of supersaturation calculated for the control and
citrate and calcium phosphate precipitate has been suggested
the experiment containing acetate was the same, within error,
previously [35]. This previous study found that the presence
at 10.93 and 10.94 respectively. This is to be expected, from the
of citrate delays the occurrence of crystals and the
concentrations listed in Table 3. In the presence of citrate
crystallisation rate. It was suggested in that paper that citrate
however, much higher concentrations of calcium and
adsorbs onto colloidal calcium phosphate precipitates and
phosphate were necessary for precipitation to occur. The
slows down the transformation of these particles into a more
higher concentration is partly due to complex formation with
stable calcium phosphate.
citrate in solution, which results in lower free ion
The precipitates from the present study were studied
concentration of calcium according to Table 1. However, the
by SEM and X-ray diffraction. In figure 4 the X-ray diffraction
supersaturation degree is calculated based on free ion
pattern of one of the precipitates is given, as well as the
concentrations, and therefore excludes the effect of
pattern of a well-crystalline natural hydroxylapatite, for
complexation. This implies that any increase in the required
comparison. All precipitates from this study showed very
supersaturation for precipitation to occur is the result of a
similar patterns, of poorly crystalline hydroxylapatites, with
specific to the organic ligand mechanism of inhibition. The
no evidence for any other phase being present. The
supersaturation in the presence of citrate was found to be
precipitates were studied using electron diffraction, but these
11.73, which is significantly higher than that in the control or
results will be discussed in detail in a future publication. In
the acetate experiment, and well outside any experimental
summary, electron diffraction also showed hydroxylapatite as
variation. To emphasise the significance of the results, the
the only phase present, and confirmed that all precipitates
calculated free calcium and phosphate concentrations are
were fully crystalline, although the size of some crystallites is
shown in Figure 3a and b respectively. In Figure 3a the total
so small, that they appear as X-ray amorphous. This is further
calcium concentration is compared with the calculated free
demonstrated in Figure 5, where an SEM image shows a
Ca2+ concentration in the absence and presence of the studied
precipitate consisting of extremely small crystallites. Given
organic ligands. The figure shows the effect of aqueous citrate
the size of the crystallites, SEM was not able to decipher
complexation on total calcium, and emphasises that the free
differences between different precipitate particle sizes.
Table 3.
Critical experimental concentrations for calcium and phosphate.
Experiment
Calcium (10-3 mol l-1)
Phosphate (10-3 mol l-1)
Control
(i.e. no organic ligand)
Acetate
4.868 ± 0.266
2.644 ± 0.071
4.877 ± 0.393
2.683 ± 0.319
1330
Citrate
6.883 ± 0.159
3.874 ± 0.057
(a)
(b)
Figure 3. The average concentrations of calcium and phosphate in the three experiments (control, acetate, citrate): a. calcium in
which  represents the total calcium and  represents the calculated free calcium Ca2+; b. phosphate in which 
1331
represents the calculated free phosphate PO43-.
Figure 4. XRD-pattern of an experimental precipitate, shown in comparison with natural hydroxylapatite. All precipitates (with
and without organic ligands) produced very similar patterns. Experimental acquisition conditions: 120 degrees
position-sensitive detector (PSD), flat (spinning) powder sample on silicate substrate, fixed beam-sample detector
geometry, beam to sample angle = 7.5 degrees and pattern acquisition time = 600 minutes.
Figure 5. Scanning electron micrograph of a precipitate formed in one of the control experiments, at pH 7. All experiments (with
1332
and without organic ligands) produced precipitates of similar (under SEM) appearance.
The results from the present study will be discussed
explains the discrepancy.
next, in the context of previous work. The observation of HAP
To conclude, a comparison of previous studies with the
precipitation in this study is in agreement with the results of
results presented here indicates that the formation of ACP
Boskey and Posner [36], who investigated spontaneous
may require extremely high supersaturation. If the
precipitation of calcium phosphate at low supersaturation
experiments are seeded, then a supersaturation similar to
(saturation degree between 5 and 9) at pH=7.4 and ionic
what we have identified as required for spontaneous
strength 0.15 M. Their experiments were carried out by
precipitation to take place (SI = 11), may result in an ACP
rapidly mixing a calcium and phosphate solution to the
precursor.
desired final supersaturation degree. After nuclei formation
It should be noted that even though no precursor phase
the crystals were allowed to grow keeping the pH constant.
to HAP was observed here, and although every care was
The precipitate was identified using electron microscopy and
taken to minimise the chance of any transformations (e.g. the
was found to be identical, even though poorly crystalline, to
filtration of the precipitate was carried out as soon as enough
hydroxylapatite. They concluded the precipitation of HAP
precipitate was visible), it is still possible that the very first
without the formation of a precursor phase.
phase that was formed was rapidly transformed into HAP.
Nancollas and Tomazic [42] investigated heterogeneous
New experiments are currently being designed to investigate
precipitation of calcium phosphate onto HAP seeds at
this possibility.
different levels of supersaturation [7 and 11]. The saturated
solutions were prepared by mixing calcium and phosphate
CONCLUSIONS
solutions in which precipitation was initiated by the addition
of a seed. The pH during the experiments was kept constant.
The results of this study show that reproducible
The precipitate was studied using X-ray diffraction and
spontaneous precipitation experiments can be performed
infrared spectroscopy. They found that at high
under thoroughly controlled conditions. It was found that
supersaturation (SI=11) an amorphous calcium phosphate
precipitation in a control experiment at pH 7 took place at a
(ACP) was formed as a precursor phase. At a low
supersaturation degree of 10.93. The influence of acetate was
supersaturation degree (SI=7) the study found formation of
found to be negligible as supersaturation was similar, at 10.94.
hydroxylapatite as the first phase to precipitate. These
Citrate, however, had a pronounced effect, as it required a
studies [36, 42] worked with solubility constant for HAP of
higher degree of supersaturation (11.73) for precipitation to
1.8*10-58 (log K = -57.74), which is the same as the one used in
take place. This is interpreted as the result of citrate inhibiting
the present study. Also Koutsoukos et al. [46] reported the
precipitation by binding to active growth sites of the nuclei.
direct precipitation of a highly crystalline hydroxylapatite
All experiments at pH 7 showed the precipitation of microonto seeding material using a constant composition method
crystalline hydroxylapatite with no apparent precursor phase.
and low supersaturation (SI=7). The precipitate was studied
using X-ray diffraction.
ACKNOWLEDGMENTS
Another previous study however [22], carried out
under spontaneous precipitation conditions, identified a
The authors wish to thank CEEP (Centre Européen d’
mixture of different amorphous precipitates and OCP; the
Etudes des Polyphosphates, a sector group of CEFIC, the
phases were identified with TEM and electron diffraction.
European Chemical Industry Council) for funding of this
However, this work was not carried out at constant pH.
project. Gordon Cressey assisted with XRD identification. Vic
Furthermore, our calculation of the concentrations at the
Din and Gary Jones are thanked for support and advice with
beginning of that study, indicates a very high (15 times)
chemical analyses and experimental set-up. Alan
supersaturation with respect to HAP. This is significantly
Pethybridge’s academic supervision of Jacqueline van der
higher than what was used in the present work, and perhaps
Houwen is acknowledged.
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