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




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НазваниеThe application of calcium phosphate precipitation chemistry to phosphorus recovery: the influence of organic ligands
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THE APPLICATION OF CALCIUM PHOSPHATE PRECIPITATION CHEMISTRY TO PHOSPHORUS RECOVERY: THE INFLUENCE OF ORGANIC LIGANDS

Jacqueline A.M. van der Houwen* (jacv@nhm.ac.uk) and Eugenia Valsami-Jones (evj@nhm.ac.uk)

Department of Mineralogy, The Natural History Museum, Cromwell Road, SW7 5BD, London, UK

Abstract


Current knowledge of calcium phosphate precipitation chemistry is brought together and compared with the needs for understanding application in processes for recovering phosphates from wastewaters. The effects of the high concentrations of organic species present in such waters appears as a key factor requiring further research and laboratory experiments in this area are described. The supersaturation required for precipitation of calcium phosphate at 25ºC, pH 7, 0.1 M ionic strength and near-stoichiometric (for hydroxylapatite) calcium to phosphate molar ratio was determined under homogeneous precipitation conditions. The experiments were carried out in air. The phase precipitated at the critical concentration was allowed to grow using a constant composition method. 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. Good reproducibility of the experiments, which were carried out in triplicate, was observed.

The study assessed the supersaturation degree necessary for precipitation of hydroxylapatite to be 10.28, assuming a solubility constant of log K=-57.13. The required supersaturation was not affected by the presence of acetate. However, citrate was found to increase the degree of supersaturation to 11.07. It is proposed that this is due to binding of citrate on the active sites of newly formed nuclei, thereby inhibiting precipitation. All experiments showed formation of a poorly crystalline hydroxylapatite.


Keywords: citrate, homogeneous precipitation, hydroxylapatite, supersaturation
INTRODUCTION


The precipitation of calcium phosphates has been studied extensively in the past 30 years with the main focus on biological processes as, for example, bone and tooth mineralisation. In the last two decades further interest in calcium phosphate precipitation has arisen from its potential as a mechanism of phosphorus removal from wastewater (1,2), although few studies have considered the underpinning chemistry. 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) (3,4), octacalcium phosphate (OCP) (5) and the inhibition of hydroxylapatite (HAP) precipitation (6). 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 been studied (7). This study concluded that phospho-citrate complexes formed on active growth sites inhibited precipitation kinetics.

The study of calcium phosphates is further challenged by uncertainty over the first phase to precipitate. Previous research suggested the formation of an amorphous calcium phosphate preceding the formation (or transformation) of a more crystalline calcium phosphate (8). Which specific crystalline calcium phosphate forms will depend mostly on pH and kinetics (9). The phase predicted to be stable is DCPD (CaHPO4*2H2O) at acidic pH around 5, OCP (Ca4H(PO4)3*2.5H2O) at pH around 6 and HAP (Ca5(PO4)3OH) at pH of 7 and above (10).

Previous studies into homogeneous nucleation proceeded by mixing of calcium and phosphate solutions at a desired supersaturation degree and leaving the nuclei enough time to grow into crystals whilst keeping the pH constant (11). 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. Homogeneous nucleation has been criticised for giving non-reproducible results due to chance nucleation onto foreign particles in the solution (12). The 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 degree for homogeneous precipitation of calcium phosphate at pH of 7 was determined. The small molecular weight organic ligands studied were acetate and citrate, which possess one and three functional groups (carboxylates) respectively. Investigation into homogeneous nucleation is important in order to compare and contrast with previous work on heterogeneous nucleation, but also to help understand problems such as the formation of fines nucleating spontaneously in industrial processes.

MATERIALS AND METHODS

Homogeneous precipitation calculations


For homogeneous 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 (13).

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 (14)

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

Log Ksp = -28.9 (15)

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

Log Ksp = -46.9 (16)

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

Log Ksp = -57.4 (17)


For the calculation of the free ion concentrations of the lattice parameters and subsequently the supersaturation degree the computer program PHREEQC was used with the Minteq database (18). The activity was calculated using the Davies equation (vi).

Log f = -Az2I/(1+I-0.3I) (vi)

f = activity

z = valency

I = Ionic strength

A is 0.5 for water at 25C (19).


Solution preparation

The solutions were made from stock solutions of ± 0.1 M organic acid. 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. The stock solutions were standardised with 0.1 M NaOH and were found not to deviate more than 7%. 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 in NaCl with a total ionic strength of 0.1 M. The initial calcium to phosphate concentrations were determined by calculation of concentrations of calcium and phosphate at a low supersaturation degree of 2. For these calculations the computer program PHREEQC was used (18). The solubility constant used in the calculations was determined in pH-drift experiments at 25C, ionic strength 0.1 M and final pH 7 ( 0.1) and was found to be log K= -57.4 (mol l-1)9 (17). The calcium and phosphate solutions added subsequently in order to slowly increase supersaturation, and the titrants were made at the molar ratio stoichiometric to hydroxylapatite (calcium to phosphate: 1.67  0.1) from the stock solutions, and adjusted to the required pH before making up the volume, in order to maintain exact concentrations. All solutions were made up in 0.1 M NaCl in order to maintain constant ionic strength during the experiment. The pH of the solutions was adjusted by addition of 0.1 M NaOH 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 (e.g. dust) 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 and phosphate 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: sample 1 was taken 1 minute after addition and sample 2 was taken immediately before the next addition to check stability of the calcium and phosphate concentrations during this induction period. A continuous drop in pH indicated precipitation. In order to maintain constant experimental conditions during precipitation, and to produce enough precipitate for identification, a constant composition method was used (20). The principle of the method is that at the critical supersaturation, where nucleation begins, the drop in pH triggers the addition 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.



Figure 1. Schematic experimental set-up for precipitation under constant composition conditions.

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 pH was only controllable when using a higher (than stoichiometric) concentration of base. The molar ratio of calcium: phosphate: hydroxyl-ions used for the titrants was 5:3:4. At the end-point of the experiment the solution was rapidly filtered through a 0.2 m syringe filter and then dried at 40°C for identification.


Analyses

During the experiments 1.5 ml samples were taken to determine calcium and phosphate concentration of the solution. 0.8 to 1 ml of sample was diluted 10 times and acidified with 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).

RESULTS AND DISCUSSION

Table 1 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 7.03 with an accuracy of ± 0.05 and experiments were found to be reproducible.


Table 1. Initial calcium and phosphate concentrations of the experiments.

Experiment


Calcium (10-4 mol l-1)

Phosphate (10-4 mol l-1)

Molar ratio (Ca:PO4)

Control (no organic ligand)

3.37

1.96

1.714

Acetate

3.48

1.99

1.746

Citrate

3.29

2.00

1.644


Experimental results are shown in figures 2a, b and c. The figures display one of each of the triplicate result graphs for the three different experimental conditions studied here (control, acetate, citrate). These graphs show the stepwise increase of calcium and phosphate over time until precipitation takes place.











Figures 2a, b and c. Determination of the critical point of precipitation in the different experiments.

In these figures the calcium and phosphate concentrations between additions is stable which indicates no precipitation occurs until the slow drop in pH (critical point of precipitation shown as the dashed line). The calcium to phosphate molar ratio slightly increases during the stepwise addition process but was found not to increase over 1.8 (on average). At the critical point of precipitation the automatic addition of the titrants commences.

Figures 2a, b and c also show that that the calcium and phosphate concentrations remain constant during precipitation. The addition of the titrants had to continue for up to 10 minutes, in order to recover enough precipitate. In table 2 the critical concentrations are given upon which precipitation took place. These concentrations are the average of three experiments and the range in concentration is also shown.

Table 2. Critical calcium and phosphate concentrations in the experiments.

Experiment

Calcium (10-3 mol l-1)

Phosphate (10-3 mol l-1)

Control (i.e. no organic ligand)

4.868 ± 0.306

2.644 ± 0.078

Acetate

4.877 ± 0.451

2.683 ± 0.341

Citrate

6.883 ± 0.184

3.874 ± 0.065


These concentrations are used in the calculation of the supersaturation degree using the computer program PHREEQC (18). The supersaturation degree calculated for the control and the system with acetate present was the same, within error, at 10.55 and 10.56 respectively. This is to be expected, from the concentrations listed in table 2. In the presence of citrate however, much higher concentrations of calcium and phosphate were necessary for precipitation to occur. The higher concentration is partly due to complex formation with citrate in the solution, which decreases free ion concentration of calcium and phosphate. However the supersaturation degree is calculated based on free ion concentration, and therefore excludes the effect of complexation. In other words, variations in the degree of supersaturation imply an effect at the mineral surface. The supersaturation degree in the presence of citrate was found to be 11.33, which is significantly higher than the control or the acetate experiment. As this is not an aqueous effect, it has to be the result of a surface inhibition. It is therefore likely that it is due to citrate binding onto active sites of the newly formed nuclei and inhibiting growth, until there is enough critical mass of phosphate and calcium in solution for precipitation to occur regardless.

The precipitate was studied by X-ray diffraction. In figure 3 the pattern of one of the precipitates is given in comparison with the pattern of natural hydroxylapatite; all precipitates showed a similar pattern.





Figure 3. XRD-pattern of the precipitate formed in comparison with natural hydroxylapatite.

As can be seen in this figure, the pattern shows a poorly crystalline hydroxylapatite but no other phase is present. These results imply that the first and only crystalline phase to precipitate in all experiments presented here is hydroxylapatite.

The observation of HAP precipitation in this study is in agreement with the results of Boskey and Posner (11), who investigated HAP homogeneous precipitation at low supersaturation (saturation degree between 5 and 9) at pH=7.4 and ionic strength 0.15 M. Their experiments were carried out by rapidly mixing a calcium and phosphate solution to the desired final supersaturation degree. After nuclei formation the crystals were allowed to grow keeping the pH constant during this time. The precipitate was identified using electron microscopy and was found to be identical, even though poorly crystalline, to hydroxylapatite. They concluded the precipitation of HAP without the formation of a precursor phase.

Nancollas and Tomazic (21) investigated heterogeneous precipitation of calcium phosphate onto HAP seeds at different levels of supersaturation. Precipitation in seeded experiments can take place at lower supersaturations than homogeneous nucleation. The range of supersaturation degree for calcium phosphate to precipitate in the presence of seeding material in their study was between 7 and 11. The saturated solutions were prepared by mixing calcium and phosphate solutions in which precipitation was initiated by the addition of a seed. The pH during the experiments was kept constant. The precipitate was studied using X-ray diffraction and infrared spectroscopy. They found that at high supersaturation (SI=11) an amorphous calcium phosphate (ACP) was formed as a precursor phase. At a low supersaturation degree (SI=7) the study found formation of hydroxylapatite as the first phase to precipitate. These previous studies (11, 21) worked with a solubility constant for HAP of 1.8*10-58 (log K = -57.74) which is similar to the solubility constant used here. Also Koutsoukos et al (22) reported the direct precipitation of a highly crystalline hydroxylapatite onto seeding material using a constant composition method and low supersaturation (SI=7). The precipitate was studied using X-ray diffraction.

A comparison of these studies with the results presented here indicates that the formation of ACP may depend on the supersaturation degree. In seeded experiments at high supersaturation, similar to the supersaturation degree necessary for homogeneous nucleation to take place (SI = 11), it is possible for ACP to form as precursor. Another possibility might be that both ACP and HAP will form simultaneously at a degree of supersaturation high enough for homogeneous precipitation to occur.

It should be noted that even though no precursor phase to HAP was observed here, and although every care was taken to minimise the chance of any transformations (e.g. the filtration of the precipitate was carried out as soon as the precipitate was visible), it is still possible that the first phase formed was another calcium phosphate, which rapidly transformed into HAP. Appropriate experiments are currently being designed to address this possibility (precipitation of calcium phosphate from supersaturated solution in a specially designed chamber within an X-ray diffractometer with a position sensitive detector). The possible formation of a mixture of HAP with ACP cannot be disregarded as a non-crystalline phase cannot be identified using X-ray diffraction. Chemical analyses of the precipitate can only provide information on the calcium to phosphate molar ratio of the precipitate, but will not necessarily prove which calcium phosphate phase has formed, as it is very likely a non-stoichiometric hydroxylapatite may precipitate (23, 24). Finally, it should be noted that the supersaturated solutions used had the stoichiometric HAP molar ratio, although the precipitated apatites had a molar ratio ranging from 1.49 to 1.65.

CONCLUSIONS

The results of this study show that reproducible homogeneous precipitation experiments can be performed under thoroughly controlled conditions. It was found that precipitation in a control experiment at pH 7 took place at a supersaturation degree of 10.55. The influence of acetate was found to be negligible as supersaturation was similar, at 10.56. Citrate, however, had a pronounced effect, as it required a higher degree of supersaturation (11.33) for precipitation to take place. This is interpreted as the result of citrate binding to the active sites of the nuclei inhibiting precipitation. All experiments at pH 7 showed the precipitation of poorly crystalline hydroxylapatite with no apparent precursor phase.

ACKNOWLEDGMENTS

The authors wish to thank CEEP (Centre Européen d’ Etudes des Polyphosphates, a sector group of CEFIC, the European Chemical Industry Council) for funding of this project. Gordon Cressey assisted with XRD identification. Vic Din and Gary Jones are thanked for support and advice with chemical analyses and experimental set-up. Alan Pethybridge’s academic supervision of Jacqueline van der Houwen is acknowledged.

REFERENCES

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2. Seckler, M.M., Bruinsma, O.S.L. and van Rosmalen, G.M. Calcium phosphate precipitation in a fluidized bed in relation to process conditions: a black box approach. Water research, 30: 1677-1685 (1996).

3. Frèche M., Rouquet N., Koutsoukos P. and Lacout J.L. Effect of humic compounds on the crystal growth of dicalcium phosphate dihydrate. Agrochimica, vol. 36, no.6, pp. 500-510 (1992).

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