PRECIPITATION OF CALCITE IN THE PRESENCE OF INORGANIC PHOSPHATE L.J. PLANT* AND W.A. HOUSE Centre for Ecology and Hydrology, Winfrith Technology Centre, Dorchester, Dorset DT2 8ZD, UK. Email: lop@ceh.ac.uk. ABSTRACT The precipitation of calcite from Ca(HCO3)2 solutions was studied over a wide range of dissolved inorganic phosphorus concentrations (0- 500 mol dm-3 KH2PO4) and a dissolved calcium concentration of 2.5 mmol dm-3 at 10 0C. At low phosphorus concentrations (< 20 mol dm-3), calcite growth was retarded and phosphate ions were coprecipitated according to theory and gave a phosphate surface density of coprecipitated phosphate of 0.150 mol m-2. As the phosphorus concentration was increased further, calcite growth ceased and a separate calcium phosphate phase was formed. In these experiments CO32ion concentrations were < 190 mol dm-3. The calcite seed played no part in the nucleation of the calcium phosphate phase as the growth rates were found to be independent of the amount of calcite seed material used in the reactions. The boundaries for the homogeneous nucleation of calcium phosphate were investigated by simple solution mixing of Ca(HCO3)2 and KH2PO4 over a wide range of supersaturations and quantified in terms of calcium hydroxyapatite, tricalcium phosphate and octacalcium phosphate saturation indices. Keywords: Phosphorus, calcite precipitation, calcium phosphate minerals INTRODUCTION There is a great deal of information available in the literature about the precipitation of calcite and the effects of inhibitors on the kinetics [1]. There have also been studies of the effects of inorganic phosphate on the crystal growth of calcite in both laboratory studies [2, 3, 4] and in field studies of hardwater lakes and rivers [5, 6, 7]. There is no doubt from these studies that phosphate in solution inhibits the formation of calcite by retarding the precipitation reaction and is coprecipitated as part of the calcite lattice [3, 8]. At low phosphate concentrations, typically < 20 mol dm-3, the coprecipitated phosphorus ions are incorporated in the calcite surface with no evidence for a distinct calcium phosphate phase being formed [8]. However, at higher phosphate concentrations there is the possibility that the calcite surface will favour the formation of a distinct calcium phosphate phase and therefore promote the over-growth of a calcium phosphate where calcite growth is impaired [9]. The present study extends previous seeded growth work on the coprecipitation of phosphate with calcite at low phosphate concentrations (< 20 mol dm-3) to examining the effects of higher phosphate concentrations (up to 500mol dm-3) more typical of the concentrations in treated sewage effluent. EXPERIMENTAL All chemicals were AR Grade unless otherwise specified. Solutions were prepared using ultra pure water (Purite Select, Analyst HP) of conductivity < 1 S cm-1. The pH glass electrode (Radiometer GK2401C) was calibrated before each experiment using freshly prepared NBS buffer solutions, KH 2PO4 / Na2HPO4 and Na2B4O7.10H2O [10]. The calcite seed material was AR Grade calcium carbonate, identified by XRay diffraction as calcite, with a specific surface area as determined by nitrogen gas adsorption and BET analysis of 0.23 m2 g-1. Free-drift experiments Treated sewage effluents are best mimicked using a Ca(HCO3)2 solution as a base that often contributes approximately 90 % of the ionic strength of real systems. A stock calcium bicarbonate solution (5 mmol dm-3) was prepared by dissolving calcium carbonate in water in the presence of carbon dioxide gas. The stock solution was diluted two-fold to give working solutions for the seeded growth experiments. Appropriate volumes of a 0.01 mol dm-3 KH2PO4 were added to the calcium bicarbonate solution to give concentrations of phosphorus between 3 and 18 mol dm-3. In later experiments the phosphorus concentration was increased to 50, 100, 250 and 500 mol dm-3 by addition of 0.1 mol dm-3 KH2PO4. All experiments were carried out in a 1 litre polypropylene reaction vessel immersed in a water thermostat (Heto BirkerØd). The vessel was continuously stirred at ca 200 rpm and left to equilibrate at 10 +/- 0.1 ºC (unless otherwise specified). The pH of the solution was adjusted by bubbling with nitrogen gas to remove dissolved carbon dioxide to achieve a desired calcite supersaturation. For experiments at lower phosphorus concentrations of 3 – 18 mol dm-3, an equilibrium period of approximately 20 minutes was allowed. However at the higher phosphorus concentrations of 50, 100, 250 and 500 mol dm-3, nitrogen gas was bubbled slowly through the solution for the entire experiment to maintain a low carbon dioxide concentration and high supersaturation with respect to calcite growth. The crystallisation of calcite was initiated by the addition of a known amount of calcite seed (ca 800 mg) as a suspension of concentration of 400 mg g-1. The pH and conductivity were monitored in the solution prior to seed addition, and for the duration of the experiment using a Radiometer PHM64 research pH meter and a Wayne Kerr automatic precision bridge B905 respectively. At time intervals samples (40 ml) were removed through a glass frit, filtered through a 0.45 m cellulose nitrate membrane, and analysed for total dissolved calcium and phosphorus. The established ‘Murphy and Riley’ colorimetric method [11], using a Beckman DU520 spectrophotometer, was employed for phosphorus determination (limits of detection, 0.3 mol dm-3). Calcium was determined using bench-top EDTA titration, where the limit of detection was controlled by one drop and was ca. 0.4 mol dm-3 [12]. A series of similar seeded growth experiments were also performed at an initial dissolved phosphorus concentration of 100 mol dm-3 using a range of seed amounts to promote crystal growth (0 to 4000 mg calcite seed). Experiments without calcite seed A series of experiments were carried out in test tubes at room temperature to investigate the conditions for homogenous precipitation of calcium phosphate minerals. Small volumes of KH2PO4, (0.5 mmol dm-3 to 0.5 mol dm-3), were added to a calcium bicarbonate solution (2.5 mmol dm-3). The solutions were mixed manually in the test tube and viewed using a Rank Brothers microelectrophoresis apparatus immediately after mixing. The field of view was divided into 50 m grid enabling sub-micron particles to be detected. The presence or absence of a precipitate was noted. THEORY Solution composition The kinetics of precipitation of phosphate with calcium, either as a coprecipitate with calcite or the formation of a distinct calcium phosphate mineral, is controlled by the solution composition at a fixed temperature. The chemical speciation of solutions used in the present experiments were determined using a speciation program, PHREEQ [13]. The program was modified to include other mineral phases likely to form as a precursor to the thermodynamically stable calcium hydroxyapatite (HAP), such as dicalcium phosphate dihydrate (DCPD), octacalcium phosphate (OCP) and crystalline tricalcium phosphate (TCP). The temperature dependence for the solubility data of each mineral phase was obtained from studies reported in the literature and entered using the format shown in Table 1. The solubility data for amorphous phases of tricalcium phosphate was not suitable for a temperature of 10 ºC [18]. The saturation index (SI) and Ion Activity Product (IAP) were output from the speciation program, e.g. for calcite: SI log 10 IAP K calcite IAP aCa 2 aCO 2 (i) (ii) 3 where ai is the activity of the ith ion, and Kcalcite is the solubility product of calcite at the temperature of the reaction. Coprecipitation kinetics. A simple two-component adsorption model relating the simultaneous precipitation of calcite and adsorption of inorganic phosphorus on the calcite surface can be used to describe the reaction coprecipitation kinetics [19]. The literature proposes the combination of the relevant isotherm equation and an equation describing the growth of calcite. Assumptions relating to precipitation are made in order to derive a general coprecipitation equation [19]: t dnCaT N A hsol dt dt 0 dnPT where nPT (iii) = Total number of moles of inorganic phosphorus coprecipitated. = Surface density of coprecipitated phosphorus (mol m-2). NA = Avagadros number (6.022 x 1023 mol-1) = Molecular area of CaCO3 on the surface (20.101 x 10-20 m2) nCa T = Number of moles of calcium containing species (total dissolved calcium) h(sol) = Function describing how the form of the isotherm depends on the solution (sol) under consideration. t = time / min This can be further simplified to give: n PT N Ag (t ) (iv) t where g (t ) h( sol )dnCaT 0 The function h(sol) illustrates how the form of the isotherm depends on the solution under consideration. This can be determined by measuring the four independent variables, temperature, pH, total dissolved phosphorus and total dissolved calcium, and adsorption constants K1 and K2 for PO43- and HPO42respectively as given in a two-component Langmuir adsorption isotherm: h(sol) K1[PO 34 ] K 2 [HPO 24 ] 1 K1[PO 34 ] K 2 [HPO 24 ] (v) The surface density parameter, , (equation iii) is an important indicator of the amount of phosphorus coprecipitated on the surface of calcite. The surface density parameter, , was calculated from equation (iv), i.e. a plot of the amount of phosphorus coprecipitated Δ n PT against g(t) to yield a slope of NA. RESULTS AND DISCUSSION Seeded growth experiments at low phosphorus concentrations The results from six experiments obtained using the seeded growth method in free-drift conditions are summarised in Table 2. All experiments were carried out at 10 +/- 0.1 C, except for experiments 1 and 3 which were performed at 20 +/- 0.2 C. It is shown that the higher temperature results in a greater saturation index with respect to HAP and a lower surface density of coprecipitated phosphorus on the calcite surface in comparison to experiments at 10 +/- 0.1 C. The experiment without phosphate addition gave the most rapid kinetics with a fast decrease in the solution pH at the beginning of the reaction as calcite formed from the Ca(HCO3)2 solution, and then a much slower increase in pH as the reaction slowed and carbon dioxide gas was lost from the reactor (Figure 1). In the other experiments as the initial dissolved phosphorus concentration increased, the precipitation reaction became slower. To achieve a reaction at the higher phosphorus concentrations, the initial pH (and saturation index) was increased as shown in Table 2. In experiment 6, even with an initial saturation index of 1.63 with respect to calcite, the reaction was relatively slow as shown by the gradual decrease in pH throughout the reaction (Figure 1). These results are consistent with earlier findings that dissolved phosphorus inhibited calcite growth but that the effects of the inhibition could be partly overcome by increasing the initial supersaturation prior to calcite seed addition [2]. The results from the experiments were analysed using equation (iii) to obtain values of the surface density of coprecipitated phosphorus, . Pearson’s correlation coefficients of between ca 0.8 and 0.98 were calculated with values of the surface density given in Table 2. The mean value for all the experiments was 0.150 mol m-2 at 10 0C, which corresponds with 0.150 mol m-2 obtained by House and Donaldson [3] and is similar to values calculated by other researchers at a similar temperature [4]. Seeded growth experiments at high phosphorus concentrations In contrast to the previous experiments, the addition of the calcite seed did not produce a decrease in pH during experiments 7-10. Instead the pH gradually increased during all the reactions as a result of the continual nitrogen purging with no evidence of carbon dioxide production as a result of calcite precipitation. As shown in Table 3, all the solutions were greatly supersaturated with respect to calcite and also calcium phosphate minerals prior to any reaction. However it was noticed that the soluble phosphorus and calcium concentrations measured in samples taken immediately before the calcite seed was added to the reactor, were lower than the expected values. This is illustrated in Table 4 where the differences are particularly marked for experiments 9 and 10, e.g. for the highest phosphorus concentration over half of the dissolved phosphorus was missing from solution after equilibrating for approximately 30 min. This indicated that precipitation of some phase was occurring prior to the calcite seed addition. After the seed addition, the concentrations of calcium and phosphorus decreased slowly (Figure 2). Changes in dissolved calcium were relatively small and difficult to measure at small time intervals by the titration method. However, changes in the dissolved phosphorus and calcium over the entire experiment, i.e. P and Ca respectively, gave much lower molar ratios of Ca/P than obtained for calcite precipitation at phosphorus concentrations < 20 mol dm-3, i.e. predicted Ca/P ratios are greater than 50 [20] and compare with an estimate from experiment 6 (Table 2) of Ca/P=90. The ratios are much lower than this for experiments 7 and 8 (at initial phosphorus of 50 and 100 mol dm-3) but not low enough to be consistent with the precipitation of a calcium phosphate phase alone. Also the change in solution composition in these experiments was relatively small compared with experiments 9 and 10, causing larger uncertainties in the Ca/P ratios. The ratios for experiments 9 and 10 (Table 4) are in close agreement to the ratio expected for calcium hydroxyapatite (HAP), viz 1.67 for HAP compared with 1.50 for tricalcium phosphate (TCP). The precipitate was mixed with the calcite seed material and therefore could not be identified separately by X-Ray diffraction. The application of the results to the coprecipitation model (equation iii) failed to produce linear relationships that were observed at the lower dissolved phosphorus concentrations. As precipitation occurred prior to calcite addition (see Table 4), even in experiments at initial phosphorus concentration of 100 mol dm-3, further experiments were performed to assess whether precipitation was influenced by the presence of calcite seed. These experiments are listed in Table 3 and include a reaction without seed addition (experiment 13) and with 800, 1600 and 4000 mg of seed for experiments 8, 11 and 12 respectively. The reaction kinetics are summarised in Figure 3 with the initial solution conditions given in Table 3. As shown, the reaction without seed addition gave a similar decrease in dissolved phosphorus concentration in the reactor compared with experiment 8 with 800 mg of seed, both starting at a similar saturation index (Table 3). Experiments 11 and 12 with 1600 and 4000 mg of seed respectively, produced a similar decrease in dissolved phosphorus over a period of ca 2000 min (Figure 3) and both experiments started with SI values for calcite, HAP and TCP in good agreement. The initial growth rates for these experiments were calculated for t< 220 min using the soluble calcium concentrations. Values varied between 0.35 to 0.78 mol min-1 but with no systematic trend with the amount of seed used in the experiments, e.g. the initial calcium loss rate in the experiment without seed addition was 0.36 mol min-1 which compares with 0.59, 0.78 and 0.35 mol min-1 for seed additions of 800, 1600 and 4000 mg respectively. There is evidence from studies in solutions dominated by the Ca(HCO3)2 system that the precipitation of a calcium phosphate (identified as an amorphous tricalcium phosphate) is possible at a SIHAP of ca. 14.4 corresponding to a solution of pH 8.6, calcium concentration 2.9 mmol dm-3 and a dissolved phosphorus of 110 mol dm-3 (Temperature of 17 ºC) [21]. The saturation condition is similar to that described for experiment 8 (Table 3). Other research in the absence of CO32- ions using Ca(NO3)2 and KH2PO4 solutions at 25 ºC, has also indicated the precipitation of an amorphous calcium phosphate phase at SIHAP values of 14.5 and greater [22]. The amorphous phases subsequently transform to the more thermodynamically stable phase of calcium hydroxyapatite (HAP), possibly through other precursor phases [23]. Experiments without calcite seed The experiments at the higher dissolved phosphorus concentrations indicated that precipitation was occurring before calcite seed addition. There is evidence to therefore suggest that calcite is ineffective as a nucleating material for calcium phosphate growth and that in experiments 7-10 homogenous nucleation is occurring. This is contrary to experiments reported by Donnert [24], involving direct precipitation of calcium phosphate induced by calcite as the seeding material. Use of calcite seed material was shown to be effective as a nucleating media, particularly in solutions containing CO32- ions, and increasing the calcite seed concentration was seen to improve the removal of phosphorus from a “German” waste-water [25]. However it was also reported that according to recent results the calcite seed material contained only some 6 % phosphorus after the application of calcite crystallisation [24]. The initial saturation of initial solutions with respect to HAP was not reported. The conditions for homogeneous precipitation of calcium phosphate phases is uncertain particularly in solutions containing CO32- solutions. The results from mixing Ca(HCO3)2 and KH2PO4 solutions at room temperature (ca 20 0C) are given in Figure 4. Each solution composition used in the experiments was analysed using the speciation program to obtain values for the saturation indices with respect to calcite, HAP, OCP, TCP and DCPD (see Table 1). The solutions tested included compositions that were undersaturated with respect to calcite but supersaturated with respect to the calcium phosphate phases. A clear division between solutions that spontaneously precipitated and those that did not was observed, and is indicated by the open and filled symbols in Figure 4. The dashed “tramlines” shown in Figure 4 also mark the limits of supersaturation with respect to HAP that were found in the seeded growth experiments 7-13 (Table 3). It was found that solutions with a saturation index greater than 9.4 with respect to HAP (and 2.2 with respect to TCP) were unstable and precipitated a solid spontaneously on mixing the solutions. As the supersaturations for the initial solutions conditions obtained in the experiments at the higher phosphorus concentrations were above these values, homogeneous precipitation is expected. This was confirmed by the initial decrease in dissolved phosphorus and calcium in experiments 8 to 10 (Table 4) measured immediately before seed addition and after mixing the KH2PO4 solution and allowing thermal equilibration. The results are also consistent with the coprecipitation of phosphate with calcite at the lower phosphorus concentrations, where initial saturation indices with respect to HAP formation were < 8.05 at temperatures of 10 +/- 0.1 C (Table 2). Precipitation indicated by loss of dissolved phosphorus and calcium, before seed addition was not found to occur. CONCLUSION At dissolved phosphorus concentrations less than 20 mol dm-3 calcite precipitation is able to occur as long as the initial supersaturation with respect to calcite is sufficient to promote growth. During precipitation from Ca(HCO3)2 solutions, inorganic phosphate ions are removed from solution and coprecipitated with calcite with the kinetics corresponding to a theoretical interpretation proposed previously. However, as the concentration of dissolved phosphorus increases, calcite precipitation ceases and there is evidence from the changes in the solution composition, for the formation of a calcium phosphate phase. This has been confirmed by simple solution mixing experiments with Ca(HCO3)2 and KH2PO4 solutions covering a wide range of supersaturations in the absence of calcite seeds, but in the presence of CO32- ions in alkaline solution. These results indicate that homogeneous nucleation of a calcium phosphate phase occurs at SIHAP>9.4. Further research is needed to identify the phase formed and how the presence of CO32- ions controls the reaction kinetics. A surprising result from the study is that calcite was ineffective in promoting the nucleation of a calcium phosphate phase. Precipitation occurred prior to the addition of the calcite seed material and the results were consistent with homogeneous nucleation in solution. ACKNOWLEDGMENTS We thank Dr A.D. Pethybridge for his comments and help and CEEP for funding this research. REFERENCES 1. Brady P.V. and. House W.A., Surface-controlled dissolution and growth of minerals. In: Physics and Chemistry of Mineral Surfaces, Brady P.V. (ed.), CRC Press, New York, pp. 225-305 (1996). 2. House W.A., Inhibition of calcite crystal growth by inorganic phosphate. J. Coll. Interface Sci., 119, 505-511 (1987). 3. House W.A. and Donaldson L., Adsorption and coprecipitation of phosphate on calcite. J. Coll. Interface Sci., 112, 309-324 (1986). 4. Hartley A.M., House W.A., Callow M.E. and Leadbeater B.S.C., Coprecipitation of phosphate with calcite and the presence of photosynthesising green algae. Water Res., 31, 2261-2268 (1997). 5. Koschel R., Benndorf J., Proft G. and RecKnagel F., Calcite precipitation as a natural control mechanism of eutrophication. Arch. Hydrobiol., 98, 380-408 (1983) 6. Reddy M.M., Carbonate precipitation in Pyramid Lake, Nevada. In: Mineral Scale Formation and Inhibition, Amjad Z. (ed.), Plenum Press, New York, pp. 21-32 (1995). 7. Woodruff S.L., House W.A., Callow M.E. and Leadbeater B.S.C., The effects of a developing biofilm on chemical changes across the sediment-water interface in a freshwater environment. Internat. Rev. Hydrobiol., 84, 509-532 (1999). 8. Hinedi Z.R., Goldberg S., Chang A.C. and Yesinowski J., A 31P and 1H MSA NMR study of phosphate sorption onto calcium carbonate. J. Coll. Interface Sci., 152, 141-160 (1992). 9. Stumm W. and Morgan J.J., Aquatic chemistry. An Introduction emphasizing chemical equilibria in natural waters. Wiley-Interscience, New York, pp. 558-563 (1970). 10. Eisenman G., In: Glass electrodes for hydrogen and other cations, Eisenman G. (ed.), Marcel Dekker Inc., New York, pp. 35 (1967). 11. Murphy J. and Riley J.P., A modified single solution method for the Determination of Phosphate in Natural Waters. Anal. Chim. Acta., 27, 31-36 (1962). 12. Vogel A.I., A Textbook of Quantitative Inorganic Analysis Including Elementary Instrumental Analysis, Longmans, London, pp. 426-427 (1961). 13. Parkhurst D.L., PHREEQC 2.2 – A computer program for speciation, reaction-path, advective transport, and inverse geochemical calculations. U.S. Geological Survey, Colorado (2000). 14. Mcdowell H., Gregory T.M. and Brown W.E., Solubility of Ca5(PO4)3OH in the system Ca(OH)2H3PO4-H20 at 5, 15, 25 and 37 ºC. J. Res. National Bureau of Standards, 81A, 273-281 (1977). 15. Gregory T.M., Moreno E.C. and Brown W.E., Solubility of CaHPO4.2H2O in the system Ca(OH)2H3PO4-H20 at 5, 15, 25 and 37.5 ºC. J. Res. National Bureau of Standards, 74A, (1970). 16. Gregory T.M., Moreno E.C., Patel J.M. and Brown W.E., Solubility of -Ca3(PO4)2 in the system Ca(OH)2-H3PO4-H2O at 5, 15, 25 and 37 C. J. Res. National Bureau of Standards, 78A, 667-674 (1974). 17. Tung M.S., Eidelman N., Sieck B. and Brown W.E., Octacalcium phosphate solubility product from 4 to 37 ºC. J. Res. National Bureau of Standards, 93, (1988). 18. Christoffesen M.R., Christoffersen J. and Kibalczyc W., Apparent solubilities of two amorphous calcium phosphates and of octacalcium phosphate in the temperature range 30-42 ºC. J. Crystal Growth., 106, 349-354 (1990). 19. House W.A., The prediction of phosphate coprecipitation with calcite in freshwaters. Wat. Res., 24, 1017-1023 (1990). 20. House W.A., The physico-chemical conditions for the precipitation of phosphate with calcium. Env. Tech. 20, 727-733 (1999). 21. Moutin T., Gal J.Y., Halouani H.EL., Picot B. and Boutoux J., Decrease of phosphate concentration in a high rate pond by precipitation of calcium phosphate: theoretical and experimental results. Wat. Res., 26, 1445-1450 (1992). 22. Meyer J.L. and Eanes E.D., A thermodynamic analysis of the amorphous to crystalline calcium phosphate transformation. Calcif. Tiss. Res., 25, 59-68 (1978). 23. Heughebaert J.C, De Rooij J.F. and Nancollas G.H., The growth of dicalcium phosphate dihydrate on octacalcium phosphate at 25 C. J. Crystal Growth, 77, 192-198 (1986). 24. Donnert D. and Salecker M., Elimination of phosphorus from waste water by crystallization. Env. Tech. 20, 735-742 (1999). 25. Donnert D., Investigations on phosphorus removal from waste water, progress report. Kernforschungszentrum Karlsruhe GmbH, Karlsruhe, (1988). Table 1. Temperature dependence for the solubility data for calcite and calcium phosphate mineral phases. Solubility dependence expressed in following form: log K A1 A2T A3/T A4logT A5/T 2 where K is the thermodynamic solubility product, T is the temperature (K) and Ai, are coefficients. Calcium phosphate phase Temperature dependence of solubility. Ref Hydroxyapatite (HAP) Log K = -8219.41/T-1.6657–0.098215T 14 Dicalcium dihydrogen Log K = -3649.5701/T+18.180752–0.0420307T 15 Log K = -45723.26/T+287.4536–0.546763T 16 Octacalcium phosphate Log K = 1039082.71+196.004417T-(48333918.7/T) - 17 (OCP) (388521.579logT)+(2303600498/T2) Calcite (cal) Log K = 13.543 - (0.0401T)-(3000/T) phosphate (DCPD) Crystalline Tricalcium phosphate (TCP) 13 Table 2. Summary of results from the seeded growth experiments with phosphorus concentrations < 20 mol dm-3. Mass of seed in each experiment was ca 1000 mg. Initial data is representative of samples taken immediately prior to seed addition. Expt. Temp ºC Initial P Initial pH Initial SIcalcite /mol dm-3 No. Final P Surface Initial /mol dm-3 density SIHAP / mol m-2 1 19.8 0 8.34 1.26 0 - - 2 9.9 3.3 8.54 1.27 0.17 0.159 4.51 3 19.8 4.7 8.46 1.45 0.03 0.27 9.97 4 10.1 8.3 8.45 1.24 4.0 0.139 5.50 5 10.0 11.4 8.70 1.47 6.1 0.123 6.77 6 9.9 18.4 8.95 1.63 11.5 0.221 8.05 Table 3. Results of the seeded growth experiments at higher dissolved phosphorus concentration at 10 +/0.1 C . Initial P concentrations are calculated from known standard additions. Saturation indices are given for calcite, calcium hydroxapatite and crystalline tricalcium phosphate. Expt. No. Length of Mass of seed/ mg Initial pH Initial P SIcalcite SIHAP SITCP /mol dm-3 expt./ min 7 2673 ca 800 9.00 50 1.70 13.9 3.76 8 7200 ca 800 8.83 100 1.52 14.3 4.10 9 1486 ca 800 8.67 250 1.38 14.9 4.59 10 5923 ca 800 8.80 500 1.44 16.2 5.41 11 3266 ca 1600 8.67 100 1.42 13.79 3.87 12 2755 ca 4000 8.72 100 1.44 13.92 3.93 13 1649 0 8.93 100 1.66 14.60 4.25 Table 4. Comparison of the seeded growth experiments with the same mass of calcite seed but high dissolved phosphorus concentrations (50- 500 mol dm-3) at 10 +/- 0.1 C . Ca / P are the molar amounts of calcium and phosphorus removed from the solution at the end of the experiment. Expt No. Nominal P measured immediately Initial Ca measured P prior to seed addition measured Ca immediately prior to seed / mol / mol dm-3 / mmol dm-3 addition Ca / P / mol dm-3 dm-3 7 50 46.1 2.50 2.50 6.0 8 100 95.8 2.44 2.42 10.6 9 250 192 2.46 2.26 1.67 10 500 206 2.47 2.00 1.68 Figure legends Figure 1. Changes in pH and dissolved phosphorus concentrations in experiments 1 and 6 (conditions shown in Table 2). Figure 2. Calcium and phosphorus concentration changes during seeded growth experiments 7-10 at 10 0 C. Time indicated as zero is the time of the calcite seed addition. Figure 3. Reaction kinetics of seeded growth experiments (< 2000 min) using different seed concentrations to initiate the reaction (experiments 8, 11-13). Conditions are shown in Table 3. Figure 4. Calculated saturation indices for experiments without calcite seed, illustrating supersaturations required for homogenous precipitation. Limits of the initial saturation indices, with respect to HAP, for experiments 7-13 (Table 3) are indicated by the dashed “tramlines”.