Solubility study of the initial formation of calcium orthophosphates from aqueous solutions at pH 5–10

Archs oral Bid. Vol. 31, No 9, pp. 565-572, Printed in Great Britain. All rights reserved

1986

ooO3-9969/86 $3.00 + 0.00

Copyright 0 1986Pergamon JournalsLtd

SOLUBILITY STUDY OF THE INITIAL FORMATION OF CALCIUM ORTHOPHOSPHATES FROM AQUEOUS SOLUTIONS AT pH 5-10 M. JOOST LARSEN and S. J. JENSEN Department of Dental Pathology and Operative Dentistry and Department of Technology, Royal Dental College, Vennelyst Boulevard, 8000 Aarhus C, Denmark Summary-In unsaturated stock solutions with various concentrations of calcium and phosphate and pH of 4-5, supersaturation with respect to brushite, hydroxy- and fluorapatite, octacalcium phosphate and amorphous tricalcium phosphate was established at 20°C by increasing pH in two series of portions, one containing low concentrations of fluoride, one fluoride-free. The precipitate as identified by X-ray diffraction was compared to the ion-activity products in the liquid. Brushite and amorphous tricalcium phosphate were formed at low supersaturation with respect to the salt in question; apatite formation required a high degree of supersaturation. Octacalcium phosphate was not observed presumably due to moderate supers.aturations with respect to this salt. Fluoride had no effect on brushite and tricalciumphosphate formation; only apatite was influenced slightly through a fluorapatite precipitation. The salts were often formed in mutual competition.

INTRODUCTION

Much of the informa:tion on the formation of solid calcium phosphates is based on X-ray diffraction and infra-red spectroscopy coupled with chemical kinetics. These techniques have demonstrated that amorphous tricalcium phosphate acts as a precursor and is transformed to apatite when pH in the aqueous phase is high (Eaner:, Gillesen and Posner, 1967; Termine and Eanes, 1972; Boskey and Posner, 1973) and to brushite when pH is low (Kaufman and Kleinberg, 1977). The factors which control the rate of the conversion process, e.g. pH, and the stabilizing agents such as various phosphate compounds, carbonate, magnesium (Bachra, Trautz and Simon, 1964; Termine, Pekauskas and Posner, 1970; West, 1971; West and Storey, 1972; Robertson, 1973; Boskey and Posner, 1977) have attracted interest particularly after the discovery of solid amorphous calcium phosphates inside cells and vesicles, where mineralization is occurring (Bourne, 1976; Larsen, Thorsen and Jensen, 1985) and in bone mineral (Termine and Posner. 1967; Posner and Betts, 1975). Fluoride has been reported to enhance amorphous calcium-phosphate formation and to inhibit the conversion to apatite (Termine et al., 1970; Hasvold and Dahm, 1977). On the other hand, other studies have shown that fluoride increases the formation of apatite and reduces the half-‘life of precursors as amorphous calcium phosphate and octacalcium phosphate (West, 1971; West and Storey, 1972; Eanes and Meyer, 1978; Aoba (~2al., 1979; Eanes, 1980). The improvement of the apatite crystallinity induced by fluoride in vitro (Bachra et al., 1964) and in bone in vivo (Zipkin, Posner and Eanes, 1962) is interesting. That amorphous tricalcium phosphate is unstable and can be hydrolysed to other salts does not mean that the formation of these salts necessarily requires a precursor. In contrast to studies in which high calcium and phosphate concentrations led to immediate amorphous calcium-phosphate precipitation,

Boskey and Posner (1976) avoided the amorphous salt precipitation by using moderate concentrations and produced apatite directly, although a considerable lag-period preceded formation. There have been few studies of the effect of the composition of the aqueous phase. The solubility products and the ion-pair constants involved have been determined with great accuracy at several occasions but they have rarely been applied in experimental work to explain a precipitation of a salt or a conversion of one solid phase to another (Kaufman and Kleinberg, 1977; Eanes and Meyer, 1977; Meyer and Eanes, 1978). Our aim was to compare the degree of saturation with respect to various calcium phosphates in a wide variety of supersaturated aqueous calciumphosphate-fluoride solutions, together with the nature of the spontaneous precipitate when these were formed.

565

MATERIALS AND METHODS

Procedures

Unsaturated stock solutions were made from Ca(OH),, neutralized by HCl, and KH*PO,, pH being adjusted to 4-5 (Fig. 1). Only reagent-grade chemicals were used. The water was distilled, filtered through Millipore filter (pore size 0.45 pm) and boiled to remove CO,. From each stock solution, two series of variously supersaturated samples of 100-300ml were prepared, one containing fluoride, the other being fluoride free. Supersaturation was established by increasing pH with KOH (Table 1). Each sample was kept in a tightly-closed flask at 20°C under gentle magnetic stirring until a precipitate was observed. The samples were watched throughout the first 6 h after initiation and terminated immediately after formation of a precipitate. None of the experiments were terminated between 18.00 h and the following morning. Immediately after observation of the precipitate, it was isolated by centrifugation

M. JWST LARSENand S. J. JENSEN

566

(Davies, 1962). The saturation with respect to the individual solid calcium-phosphates was estimated by comparing the ionic activity-products with the solubility products. The following constants were used:

1000

.

t

.

Dissociation

constants:

Phosphate-pK,

7.21

0 pK,

12.34

. Water-pKw

. 01

CaH,PO,+-K

3.67

(Gregory, Moreno and Brown, 1970). 264 (Gregory et al., 1970). 2.9 x lo6 (Chughtai, Mar&all and Nancollas, 1968).

CaHPO,-K

1

001

. CaPO;-K

I__

0

05

1

SoIubiIity products :

[Cal /[Cal+tPI

Fig. I. Ca x P products of the experimental stock solutions used plotted versus their Ca/Ca + P ratio. The open circle denotes

14.00.

constants :

Stability

t

the 3 mmol/l calcium and phosphate Tables 2 and 3 and Fig. 2.

solution

in

Ca, (PO,),OH-pK

54.6

Ca, (PO,), F-pK CaHPO,.2H,O-pK

59.6 6.60

/?-Ca, (PO,),-pK (10,OOOg) for 5 min, washed twice with water, once with ethanol and subsequently 20°C. When no solid formation occurred days, the experiments were terminated.

distilled dried at after 30

The nature of the deposit was determined by X-ray diffraction techniques and by chemical analysis of the dissolved powder. Calcium was measured by the method of Willis (1960); phosphorus was determined by that of Chen, Toriba and Warner (1956). Fluoride and pH were measured electrometrically. Before and after formation of the precipitates, the lO(r300ml samples and supernatants were analysed for calcium, phosphate and fluoride and pH was determined. The concentrations were corrected for complexes (CaH,PO:, CaHPO, and CaPO;) and the activities were calculated using the Davies activity-coefficients

PH

1. Composition PI HAP

PI BSH

of solutions PI TCP

Ca,(PO,),

(Moreno, Gregory and Brown, 1968). (McCann, 1968). (Gregory et al., 1970). (Gregory et al., 1974).

28.9

(amorphous) -pK 25.2

Ca,H(PO,),-pK CaF,-pK

Analysis

Table

(Bates and Acree, 1943), (Vanderzee and Quist, 1961).

(Meyer and Eanes, 1978). (Bjerrum, 1949). (McCann, 1968).

47.1 10.5

X-ray diffraction films were made with a Nonius Guinier camera (FeKx, 30 kV, 10 mA, I5 h, Kodirex films) and tracings were produced by means of a Joyce-Loebl double-beam densitometer. The tracings of the precipitates were compared with standards of hydroxyapatite (Bio Rad Laboratories, California, U.S.A.), Brushite (Merck, p.a.) and amorphous calcium phosphate prepared by the method of Meyer and Eanes (1978) and showing the characteristic broad maximum at 0 = 19” not coinciding with the main maximum of apatite at 0 = 20.2” (Bienenstock

prior

to precipitation

Lag time

X-ray

and nature PI FAP

of precipitates

Lag time

X-ray

4.91 5.10 5.21 5.31 5.89 6.31 8.36

56.7 55.5 54.7 54.1 50.6 48.5 41.9

(A) C,, = 2.85 mmol/l; C, = 132.3 mmol/l; 30.7 6.39 30.0 24 h BSH 6.21 29.6 24h BSH 6.11 29.2 6h BSH 6.02 27.3 imm BSH 5.57 26. I imm BSH 5.35 22.8 imm ATCP 5.43

C, = 0.18 mmol/l 51.5 50.4 49.8 24h 49.3 5h 46.2 imm 44.7 imm 40.1 imm

BSH BSH BSH BSH ATCP

6.39 6.54 6.71 6.90 7.51 8.00

51.1 50.3 49.4 48.6 46.1 44.4

(B) Cc, = 0.804 mmol/l; C, = 13.3 mmol/l; 27.9 6.28 27.4 10d BSH 6.19 27.0 8d BSH 6.11 26.5 6d BSH 6.04 25.2 30min AP 5.94 24.4 30min ATCP 5.97

C, = 0.16 mmol/l 47.4 46.7 lld 46.0 8d 45.3 6d 43.5 30min 42.3 30min

BSH BSH BSH AP ATCP

Formation

of calcium

Table

PH

PI HAP

PI BSH

PI TCP

phosphates

I-continued

Lag time

X-ray

6.82 7.03 7.30 7.60 7.79 8.09

50.8 50.0 48.9 47.8 47.2 46.3

(C) C, = 0.32 mmol/l; C, = 22.6 mmol/l; 6.36 27.8 6.31 27.3 18d BSH 6.28 26.2 10d BSH 6.28 26.8 10d BSH 6.30 25.9 47t AP 6.35 25.5 2t AP

8.24 8.70 9.20 9.54 10.31

48.1 47.0 46.5 46.3 46.8

(D) Cc, = 0.102 mmol/l; C, = 2.98 mmol/l; 7.10 26.8 7.24 26.3 14d AP 7.52 26.1 3d AP 7.78 26.2 1d AP 8.48 26.8 1d AP

52.0 50.1 48.6 46.3 44.3 40.1

(E) C, = 2.98 mmol/l; C, = 6.54 28.5 6.30 27.4 17d 6.12 26.5 1d 5.90 25.3 90min 5.78 24.2 15min 5.88 22.1 imm

3.02 mmol/l;

6.02 6.32 6.57 6.99 7.42 8.62

55.9 52.7 50.2 47.6 45.8 40.0

(F) Cca = 26.4 mmol/l; C, = 6.37 30.3 5.93 28.5 45 min 5.58 27.0 1 min 5.24 25.6 imm 5.02 24.6 imm 4.63 21.5 imm

25.2 mmol/l;

4.56 5.02 5.40 5.80 6.10 7.30

7.00 7.53 8.16 8.30 8.72 9.25

51.0 48.3 45.7 45.2 43.8 42.5

9.97 10.54 11.02

45.1 44.3 43.9

5.34 5.59 5.72 5.97 6.22 6.80 7.31

52.6 51.0 50.1 48.5 46.9 43.8 41.1

(I) Cc, = 38 mmol/l; 6.42 28.7 6.19 21.7 6.08 27.3 5.87 26.3 5.69 25.5 5.40 23.7 5.30 22.5

6.52 7.02 7.44 7.69 8.02 9.28

50.7 47.8 45.8 44.7 43.4 39.8

(J) Cc, = 2.73 mmol/l; 6.76 27.9 6.47 26.3 6.34 25.2 6.30 24.7 6.29 24.0 6.75 22.4

(G) C, 7.08 6.88 6.83 6.84 6.90 7.10

= 0.55 mmol/l; 28.2 26.8 25.4 25.2 24.5 24.0

BSH BSH BSH +AP AP ATCP

BSH BSH BSH BSH + AP ATCP*

C, = 0.62 mmol/l;

14d 44 h 30min 10min

567

AP AP ATCP ATCP

(H) Cc, = 0.105 mmol/l; C, = 0.102 mmol/l; 8.50 26.0 8.88 25.8 8d AP 9.27 25.8 48h AP

PI FAP C, = 0.072 47.9 47.2 46.5 45.7 45.2 44.7

Lag time mmol/l

20d 48t 2t

C, = 0.16 mmol/l 46.2 45.5 9d 45.4 1d 45.7 1d 46.9 1d C, = 0.14 48.0 46.3 45.0 43.2 41.6 38.5 C, = 0.10 50.6 47.6 45.7 43.5 42.0 37.4

X-ray

BSH AP AP

AP AP AP AP

mmol/l 16d 3d 45min 15min imm

BSH BSH + AP AP AP ATCP

mmol/l 2 min imm imm imm imm

BSH BSH BSH BSH + AP ATCP*

C, = 0.16 mmol/l 47.8 45.7 43.7 43.3 42.3 41.6

20h 2h 30min 10min

AP AP ATCP ATCP

C, = 0.16 mmol/l 44.9 44.7 lid AP 44.8 24h AP

C, = 2.95 mmol/l; 18d 22h 30min imm imm imm

C, = 0.037 mmol/l 48.5 BSH 47.1 20d BSH 46.4 20h BSH 45.0 6h BSH + AP 43.7 imm AP 41.1 imm ATCP 39.3 imm

C, = 0.71 mmol/l; 10d 1h 1h 1h imm

AP AP AP AP ATCP

C, = 0.054 47.5 46.2 43.5 42.7 41.7 39.4

BSH BSH BSH BSH + AP AP ATCP

mmol/l 8d 2h 15min 10min imm

AP AP AP AP AP

Cc,, C,, and C, are total concentrations of calcium, phosphate and fluoride in solutions before precipitation, p1 the negative logarithm of the ion-activity product, HAp = Ca,(PO,), OH, BSH = CaHPO,, TCP = Ca, (PO,), , FAp = Ca, (PO,), F, lag time is the time from addition of KOH to observation of precipitate (imm = immediately). In the columns headed X-ray, the nature of the precipitates is given: BSH = brushite, AP = apatite, ATCP = amorphous tricalcium phosphate, columns 556 represent fluoride-free solutions, 7-9 fluoride-containing solutions. *Traces of BSH and AP.

M. JOUSTLARSEN and S. J. JENSEN

568

and Posner, 1965). Precipitates were also considered as amorphous calcium phosphate when this maximum could not be detected and no other diffuse or weak bands or peaks could be seen. Precipitates of apatite showed a line broadening compared with the standard due either to less perfect crystallinity or to smaller crystal size. RESULTS Only a single experimental series will be presented in detail, namely that containing 3 mmol calcium and 3 mmol phosphate per litre followed by a shorter account of the remaining experiments. Table 2 shows the effect of the pH changes on the so-called free concentrations of calcium, of phosphate ion species and of their various ion pairs together with the ion-activity products derived from these data. As pH increased, the supersaturation with respect to all salts also increased, as shown by a comparison of pI (the negative logarithm of the ion product), with the pK value as given in Materials and Methods (PI= pK indicates saturation, p1 > pK undersaturation and pI < pK supersaturation; Fig. 2). With increasing pH, the supersaturation with respect to brushite steadily increased. No solid was formed in the first-row experiment presumably because the supersaturation was not sufficient. In the second-row experiment, brushite appeared. In third and fourth row, apatite was mixed with brushite, indicating that the supersaturation with respect to apatite was sufficient for a competing formation of this salt. In the last rows, no brushite was formed, although the supersaturation with respect to this salt was higher than before. Apparently, the formation of apatite and, to increasing extent also of amorphous tricalcium phosphate, outpaced that of brushite. The supersaturation with respect to the two latter salts became high which may explain their formation and the speed by which they were formed. Chemical analysis of the solids showed that the Ca:P ratio was consistent with the nature of the solids assuming that the Ca:P ratio of brushite is unity, that of the amorphous tricalcium phosphate is 3:2, and that of apatite is 5:3. Considerable variation was expected as the salts were poorly crystallized. For the fluoride-containing samples, when apatite was formed, the Ca:F ratio was low, indicating a high-fluoride content. When brushite or amorphous tricalcium phosphate was formed alone, the ratio was higher. When apatite was formed, the fluoride concentration in the aqueous phase fell to 1: 3 or less; virtually no changes of the fluoride concentrations were observed when brushite or amorphous calcium phosphate was precipitated (Table 3). The experimental series in which fluoride was added to the aqueous phase differed little from that without fluoride. The lag-time of brushite and amorphous calcium phosphate formation was entirely unaltered but a faintly increased tendency to formation of apatite was observed. Thus, the high supersaturation with respect to fluorapatite, provided by the addition of fluoride (Fig. 2) was not reflected in increased initiation of apatite formation. A direct comparison of the crystals of the various apatite materials does not seem justified because they (Table 2) were not formed at equal rates. Nevertheless, the

Formation of calcium phosphates FKS - pl

pKr - pl

(a) 12

569

( b)

16

pH = 6.02

pH = 6.32 12 8 4 * /lldLIl HBT AS PHP

HBTF ASCA PHPP

* HBTF ASCA P HP

C

pKr - pl

IIKS - pl

16

16

*

pH = 6 57

PH = 6.99

12

*

*

HBT AS C P HP

HBTF AS C P HP

*

*

Ii 6 T ASCE PHPP

HBT AS C P HP

A P

*d

(d)

20

(c)

P

F

PKS - PI

(f 1

pKs - pl 20

(e)

20-l

pH = 6.62 16

HBT AS PHP

C

H 6 T ASCA P HP

HBT ASC P HP

F P

HBl ASCA P HP

P

Fig. 2. Degrees of saturation in fluoride-free (left) and fluoride-containing aqueous solutions of 3 mmol/l calcium and phosphate at various pH prior to precipitation. The ordinate shows the difference between the negative logarithm of the solubility product and the ion product so that positive values indicate a supersaturation with respect to the individual salt. The open part of the tricalcium phosphate refers to well-crystallized fi-tricalcium phosphate, the hatched part to the amorphous salt. The asterisks indicate the nature of the salt formed.

apatite of Ca:P = 1.68 and Ca:F = 8.9 showed the best apatite pattern. The stability of the various salts depended on the comlposition of the aqueous phase after the precipitation (Table 3). When brushite was formed, the aqueous phase approached saturation and the supersaturation with respect to apatite and tricalcium phosphate decreased. After the brushite formation, the ion products of apatite and tricalcium phosphate differed little from those of the experiment in Table 2 first row, in which no solid was found within 30 days. Thus, no competing formation of another salt could be expected after the brushite formation. When both brushite and apatite (hydroxy- and fluorapatite) were formed, the aqueous phase became saturated with respect to brushite, while it was supersaturated with apatite. Thus, the formation of apatite

Table 3. Composition of solutions E after precipitation.

pH (a) 6.21 6.27 6.49 6.81 7.55 (b) 6.24 6.35 6.31 7.06 1.45

PI

PI

PI

Cc, C,

HAp BSH TCP

2.48 2.45 1.64 1.40 0.50

2.69 2.38 2.30 2.02 1.28

50.9 51.1 50.6 49.2 48.0

6.45 6.50 6.51 6.42 6.72

21.9 28.0 21.1 27.0 26.5

2.51 2.22 1.50 1.55 0.59

2.71 2.45 2.14 2.14 1.30

51.0 50.7 51.9 47.6 48.1

6.46 6.47 6.71 6.23 6.67

21.9 27.8 28.5 26.1 26.6

PI FAp

PI CaF,

C,

47.2 47.0 50.1 45.2 45.5

10.6 10.7 14.8 12.2 11.2

0.14 0.14 0.0012 0.028 0.14

(a) Fluoride-free. (b) Fluoride-containing. ations see Table I.

For abbrevi-

M. JOOSTLARSEN and S. J. JENSEN

570

(hydroxy- and fluorapatite) continued whereas brushite by dissolution maintained the liquid saturated with respect to brushite until all of this salt was dissolved: brushite is unstable in such a system. In principle, the trends of the experiment with 3 mmol calcium and 3 mmol phosphate per htre appeared also in the other experiments (Table 1). The formation of brushite did not require a high degree of supersaturation. Compared with the solubility product of 2.5 x lo-’ (pKs: 6.60), the brushite ion product was about 10 x lo-’ (PI: 6.0); this salt was formed within 24 h, whereas an increase of the lagtime was observed when the ion product was only 5 x IO-’ (PI: 6.30). A smaller ion product did not initiate brushite formation. Formation of apatite required a higher supersaturation; a ratio between the ion products and the solubility product of 2.5 x lo-” ranging around 10’ was required. When the hydroxyapatite ion product was 10m4’, apatite was formed, albeit after a lag period of 1 or 2 weeks. Higher ion products led to shorter lag periods. Thus, apatite formed within a few hours when the ion product was 10-45-10-44. Amorphous tricalcium phosphate was only formed when the supersaturation with hydroxyapatite was high. Addition of fluoride increased the supersaturation with respect to fluorapatite, indicating an increased tendency for formation of apatite. Only occasionally was the formation of apatite accelerated to an extent that another salt was replaced by apatite when fluoride was added to the liquid. Neither was fluoride taken up in brushite or amorphous tricalcium phosphate during their formation nor was a major alteration of lag time or other influence on the formation of these salts of fluoride addition observed. Almost invariably, when pH was sufficiently high, amorphous calcium phosphate was formed and often rapidly after the preparation of the solutions. Rapid formation of amorphous calcium phosphate required a tricalcium-phosphate ion-product of at least 10-24.5 with delayed formation when ion products were lower. The amorphous salt was always formed in competition with the other calcium phosphates when these solutions were highly supersaturated with respect to these salts; they were often found as precipitates at lower degrees of supersaturation. DISCUSSION

The experiments were arranged so that the liquid was always unsaturated with respect to calcium fluoride, except for the experiment of Table 2, in which the liquid was exactly saturated and that of Table 1 F, in which it was slightly supersaturated. It is, therefore, not surprising that no calcium fluoride was formed. In none of the experiments was octacalcium phosphate observed. The supersaturation with respect to this salt was only slight. Usually the ratio between the octacalcium phosphate ion products and the solubility product ranged below lo3 and often for shorter than a few hours. Although 37°C would have been a temperature more relevant to physiological processes, 20°C was chosen partly for convenience but also because the relevant constants are the more accurately determined at 20°C. The washing procedure prior to the ethanol rinse was

pKs -

pl pH = 5.21

n

F

-8 AS PHP

C

ASCA P HPP

Fig. 3. Degrees of saturation in a fluoride-free (left) and a fluoride-containing (right) solution, Cca, 2.85 mmol/l; C,, 132 mmoljl and pH, 5.21 (Table 1 A 3). Abbrevations as in Table 1. Positive values indicate a supersaturation, negative values an undersaturation. Brushite was formed in both experiments despite the low supersaturation with respect to this salt. The liquid was unsaturated with respect to both amorphous and a-tricalcium phosphate which excludes these sofids as precursors.

short to avoid the solid phase ethanol tends to induce (Larsen et al., 1985). Transformation of amorphous tricalcium phosphate

Our samples were monitored throughout the first 6 h, thus the salt observed may not always be the salt first formed. There might have been a short-lived precursor, especially with X-ray diffraction uncharacteristic amorphous calcium phosphate. However, the suggestion that brushite is formed after a transformation of amorphous calcium phosphate was not supported; brushite was formed in a series of solutions unsaturated with respect to the amorphous salt (Table 1, line A 226, B 2-4, C 224, E 2-3, F 24, I 24) but was formed even when the solution was unsaturated with respect to crystalline /?-tricalcium phosphate, pK = 28.9 (Gregory et al., 1974) (Table 1, line A 24, Fig. 3). The undersaturation indicates that no amorphous calcium phosphate could possibly be formed prior to brushite and thus that formation of brushite does not necessarily require an amorphous precursor. A similar argument applies to apatite formation. The chemical conditions, most clearly favouring the formation of apatite over that of the amorphous salt, occurred in experiments D and J in which the apatiteforming solutions were unsaturated with amorphous tricalcium phosphate but slightly supersaturated with crystalline fi-tricalcium phosphate. Under these conditions, it is unlikely that the amorphous salt acts as a precursor necessary to apatite formation. Furthermore, when the amorphous salt was formed, the liquid phase was invariably supersaturated with hydroxyapatite, and often the amorphous salt was formed in competition with apatite. In reports of conversion of amorphous calcium phosphate into apatite (Boskey et al., 1973), the amorphous salt has always been prepared quickly with high concentrations of calcium and phosphate and at pH around 7-10 and thus in an aqueous phase highly super-

571

Formation of calcium phosphates saturated immature, within the instability

also with respect to apatite. Therefore, early apatite seeds may well be present amorphous salt produced adding to the of the salt.

Transformation

of brushite

When brushite at high pH is transformed to apatite, the conversion may partly be due to the composition of the aqueous phase and partly to seeding properties of the salt. In an aqueous brushite suspension the liquid phase becomes saturated with respect to brushite, provided that formation of other salts does not tend to decrease the concentrations. In our study, some solutions were close to saturation or were slightly undersaturated with respect to brushite. For example, suspending brushite at pH 8.16 gives a composition of the aqueous phase not different from that in Table 1 line G 3 and an apatite would be formed if no inductive or other influence from the brushite powder occurred. At pH 6.02, the composition of the liquid phase would be approximately that of Table 1 line E 1 and no precipitate would occur. A number of our precipitates in solutions saturated with respect to brushite (e.g. G S-6, J 6) were amorphous calcium phosphate. It seems unlikely that brushite can be transformed to the amorphous salt. On the contrary, apatite is formed when brushite is suspended at high pH. Apparently, because the dissolution and conversion of brushite is a slow process and the aqueous phase becomes highly supersaturated with respect to apatite its transformation to amorphous calcium phosphate is precluded. Transformation

of apatite

Few apatite A 2-3). would

solutions were only saturated with respect to during formation of another salt (Table 1 line The data suggest that at pH4.S5.5 apatite be transformed to brushite if direct seeding effects of the apatite crystals are disregarded. On the other hand, no salt had ,a solubility sufficiently low to induce its formation in a high pH solution saturated with respect to apatite; this explains the stability of this salt not only in most inorganic solutions but also in the organic milieu in living organisms. The slight effect of fluoride on apatite formation contrasts strongly with its effect on growth of apatite crystals (Eanes, 1980); only a slight supersaturation with respect to fluorapatite induces apatite growth even when the aqueous phase is considerably undersaturated with respect to hydroxyapatite (Larsen, van der Fehr and Birkeland, 1976; Larsen and Thorsen, 1984). Transformation

of @-tricalcium phosphate

A similar argument applied to /?-tricalcium phosphate transformation; at low pH the fi-tricalcium phosphate may transform to brushite (Table 1, line A 24). Higher pH does not create supersaturation sufficient to induce a precipitation in solutions saturated with respect to /I-tricalcium phosphate.

Acknowledgement-The

Medical Research 512-16170.

study was supported by the Danish

Council,

grant

nos

512-8087 and

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