The application of calcium phosphate precipitation chemistry to phosphorus recovery: the influence of organic ligands

НазваниеThe application of calcium phosphate precipitation chemistry to phosphorus recovery: the influence of organic ligands
Дата конвертации25.10.2012
Размер76.6 Kb.
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Environmental Technology, Vol. 22. pp 1325-1335

© Selper Ltd, 2001


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)


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.


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 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 drift-conditions, 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 non-reproducible 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.


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-) (ii)

Log Ksp = -6.68 [39]

Log IAPTCP = 3 log (Ca2+) + 2 log (PO43-) (iii)

Log Ksp = -28.9 [40]

Log IAPOCP = 4 log (Ca2+) + 3 log (PO43-) + log (H+) (iv)

Log Ksp = -46.9 [41]

Log IAPHAP = 5 log (Ca2+) + 3 log (PO43-) + log (OH-) (v)

Log Ksp = -57.74 [42]

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] (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-Ca-CO2-H2O-system.

(Minteq database in PHREEQC [43]).

Thermodynamic dissociation products

Log k

Ca2+ + PO43- + H+ = CaHPO4


Ca2+ + PO43- = CaPO4-


Ca2+ + PO43- + 2H+ = CaH2PO4+


PO43- + H+ = HPO42-


PO43- + 2H+ = H2PO4-


PO43- + 3H+ = H3PO4


Ca2+ + CO32- + H+ = CaHCO3+


Ca2+ + CO32- = CaCO3


Organic acid equilibrium reaction

Acetate- + H+ = Hacetate


H+ + Citrate3- = CitrateH2-


2H+ + Citrate3- = CitrateH2-


3H+ + Citrate3- = CitrateH3


Ca2+ + Citrate3- = CaCitrate-


Ca2+ + Citrate3- + H+ = CaCitrateH


Ca2+ + Citrate3- + 2H+ = CaCitrateH2+


Ca2+ + Acetate- = CaAcetate+


Solution Preparation

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 ICP-AES 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 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 on-line 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.


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).

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