Archs oral Bid. Vol. 31, No. 7, pp. 58S591, Printed in Great Britain
0003~9969/92 $5.00 + 0.00
1992
Pergamon Press Ltd
THE EFFECTS OF MAGNESIUM AND FLUORIDE ON THE HYDROLYSIS OF OCTACALCIUM PHOSPHATE* M. S. TUNG, B. TOMAZICand W. E. BROWN American Dental Association Health Foundation Paffenbarger Research Center, National Institute of Standards and Technology, Gaithersburg, MD 20899, U.S.A. (Accepted 4 February 1992) Summary-The adsorption of Mg ions on octacalcium phosphate (GCP) and its effect on OCP hydrolysis, with and without F, were studied. The Mg adsorption isotherm was fitted by the Langmuir model with an affinity constant of 0.74 ml/pmol and maximum number of sites, 31.19 pmol/g. The hydrolysis rates were measured in a pH stat by titration of base and were strongly temperature dependent. The products were examined by X-ray diffraction and chemical analysis. GCP hydrolysis takes place in two stages: the fast initial process, which is attributed to the surface topotactical conversion, followed by the main, slower process, which involves the nucleation and crystal growth. Mg ions, as 1 mmol/l MgCl,, prevented the initial surface reaction and decreased the nucleation rate dramatically and the growth rate slightly; F increased the rates of surface reaction and both the nucleation and crystal growth processes. The Ca/P ratio (1.53) and the line broadening in the X-ray diffraction patterns of the apatitic products were not significantly affected by the F. Mg also did not affect the Ca/P ratio and the line broadening at (002) diffraction, but decreased the line broadening at (310) diffraction. Key words: adsorption, biomineralization,
hydrolysis, fluoride, magnesium, octacalcium phosphate.
INTRODUCITON To understand better the effects of F and Mg ions on the formation and maturation of calcified tissues, these effects have been studied in systems of synthetic apatite formation involving precursors such as amorphous calcium phosphate (Boskey and Posner, 1974; Cheng, Grabher and LeGeros, 1988), dicalcium phosphate dihydrate (Tung and Brown, 1985) and octacalcium phosphate (Ca,H,(PO,), .5H,O) (Le Geros, 1991). F and Mg have opposite effects on the formation and maturation of calcium phosphates (Brown et al., 1962; Amjad, Koutsoukos and Nancollas, 1984) and enamel (Brown, Eidelman and Tomazic, 1987). Fluoride has been shown to promote the formation of apatite and the hydrolysis of dicalcium phosphate dihydrate (Tung and Brown, 1985); it greatly curtails or eliminates the appearance of the octacalcium phosphate precursor in the spontaneous precipitation of apatite (Eanes and Meyer, 1978) as well as the formation of dental caries (reviewed in Leach, 1986; Driessens and Woltgens, 1988). On the other hand, Mg ions have been implicated in the early stages of caries (Hallsworth, Robinson and Weatherell, 1972). These ions retard the growth rates of hydroxyapatite, fluorapatite and octacalcium phosphate (Amjad et al., 1984; Salimi, Heughebaert and Nancollas, 1985) and also lengthen the lifetime of the *Certain commercial materials and equipment are identified in this paper to specify the experimental procedure. In no instance does such identification imply recommendation or endorsement by the National Institute of Standards and Technology or the ADA Health Foundation or that the material or equipment identified is necessarily the best available for the purpose.
precursors, amorphous calcium phosphate (Boskey and Posner, 1974; Cheng et al., 1988) and octacalcium phosphate, during the spontaneous precipitation of apatite (Eanes and Meyer, 1978). It has been suggested that these effects of Mg ions on the formation of calcium phosphates are due to the blocking of active growth sites through adsorption of Mg ions at the crystal surfaces (Eanes and Rattner, 1981; Amjad et al., 1984). Our purposes now were to study the adsorption of Mg ions on synthetic, well-crystallized octacalcium phosphate and the effects of F and Mg ions on the hydrolysis of octacalcium phosphate at physiological pH. Octacalcium phosphate is structurally similar to hydroxyapatite and its hydrolysis does not further involve an intermediate precursor. Formation and transformation of octacalcium phosphate have been reported under conditions similar to those in physiological environments (Tomson and Nancollas, 1978; LeGeros, Kijkowska and LeGeros, 1984). Octacalcium phosphate is also frequently present as one of the crystalline components of dental calculi (Le Geros, 1974) and, on occasions, is associated with urinary stones (MacGregor, Robertson and Nordin, 1965) and other pathological calcifications (Tung and Brown, 1985). It is believed to be a precursor in the formation of calcified tissues such as bone, teeth and pathological deposits (Brown et al., 1987; Tomazic, Etz and Brown, 1987). Therefore, hydrolysis of octacalcium phosphate provides a model for maturation processes in biomineralization. MATERIALSANDMETHODS Reagent-grade chemicals and distilled water were used throughout the experiments. Octacalcium
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phosphate was prepared by two methods. One method consisted of the hydrolysis of dicalcium phosphate dihydrate, as reported by Chickerur, Tung and Brown (1980). G&calcium phosphate prepared in this way was used throughout unless specified otherwise. It has a Ca/P ratio of 1.35 + 0.02 and a surface area of 25.16 m2/g as determined by the N adsorption BET method. The crystals of this preparation yield a characteristic X-ray diffraction pattern for octacalcium phosphate and appear as a homogeneous phase under the polarizing microscope. The second method involved the direct crystallization as reported by LeGeros (1985) with minor modifications; the phosphate so prepared is designated as octacalcium phosphate-L (Tomazic et al., 1989). Adsorption measurements
5 Ca,H,(PO,),
.5&o
+ 8 Ca,(PO&OH
Accurately weighed samples of octacalcium phosphate (0.1-0.5 g) were transferred to plastic bottles, and 2 ml of water were added. The slurry was shaken for 2 h and was saturated with octacalcium phosphate with no detectable hydrolysis at 23°C (Tung et al., 1988). A Mg-containing solution was then added, with final solution volume of 4-6 ml. The adsorption equilibration was carried out in the shaker for 1 h at 23°C. (A preliminary study had indicated that the adsorption reached equilibrium in 5 min, with no adsorption on the walls of the plastic bottles, and the pH of the slurries was 7.5, independent of the Mg concentration in the slurry.) After 1 h, the slurry was filtered through a 0.22 pm filter (Millipore, Bedford, MA, U.S.A.), and the concentration of Mg in the filtrate was analysed by atomic absorption spectroscopy. The difference between the initial and filtrate Mg concentrations is the amount adsorbed on the surface of the octacalcium phosphate crystals. The adsorption isotherm was obtained from the amounts of Mg adsorbed as a function of equilibrium concentration. A general interpretation of the isotherm is achieved by the use of the Langmuir adsorption model (Adamson, 1960) described by C/m = l/Km, + C/m,
in solution. F was added as solid CaF,. We used CaF, because it was the main product in the topical fluoridation; it released the F and Ca slowly as the reaction proceeded and it dissolved. The released Ca partially compensated the released phosphate during hydrolysis. The slurry was stirred by a magnetic bar at approx. 100 rev/mm. Several of the hydrolysis reactions were made at 50 and 60°C at pH 7.4, with 1 g of octacalcium phosphate/l to study the Mg incorporation. Periodic samplings of solutions and solids were done to determine the progression of hydrolysis and incorporation. The hydrolysis of octacalcium phosphate produces phosphoric acid by the following reaction:
(1)
where C @mol/ml) is the equilibrium concentration of adsorbate (Mg) in the solution, m bmol/g) is the amount of adsorbate and m, (pmol/g) is the maximum amount (maximum number of sites) adsorbed per gram of substrate (octacalcium phosphate), and K, the affinity constant, is a constant reflecting the affinity that the adsorbate molecules have for the adsorption sites. The Langmuir plot, C/m versus C, should be a straight line where the slope and intercept give the information about the K and mOparameters. The correlation coefficient of the least-squares linear regression was calculated to indicate the difference between the Langmuir model and the experimental data. Hydrolysis A weighed sample of octacalcium phosphate (0.87-0.89 g) and 35 ml of water (25 g of octacalcium phosphate/l) with or without Mg and/or F in a 120 ml polypropylene bottle was stirred (Mg and/or F were added before the phosphate), and maintained at pH 7.5 and 37 or 25°C. Mg was added as 1 mmol/l MgCl,
+ 6 H,PO, + 17 H,O.
(2)
The rate of hydrolysis was followed by titration with 0.18 mol/l KOH as a function of time in a pH stat. The pH-stat system included a METROHM pH meter 632, Impulsomat 614 and the titrator, Dosimat 655 (Brinkmann Instruments Inc., Westbury, NY, U.S.A.). The pH measurements were made with a glass/calomel combination electrode, which was calibrated daily. The reaction was stopped when the consumption of KOH ceased. The solids were filtered, washed with water, and dried at room temperature. The filtered solutions were analysed for Ca (Willis, 1960; McDowell, Brown and Sutter, 1971) and phosphate (Brabson et al., 1958). Because there is a net release of phosphate during hydrolysis of octacalcium phosphate, the decrease in solution Ca/P ratio is an indicator of the progression in the hydrolysis. The X-ray diffraction analysis of the solids were made with a Norelco Philips unit with Ni-filtered CuKa radiation. The crystallite size and/or crystal strain of the solid reaction products was assessed in terms of line broadening at half-peak heights, /3;, of the 002 and 310 reflections in the apatite (Klug and Alexander, 1974). RESULTS
Table 1 summarizes the results for the adsorption of Mg on octacalcium phosphate from aqueous solution at 23°C with the initial concentrations of Mg ranging from 0.016 to 7.85 mmol/l. The data in Table 1 were plotted according to the Langmuir equation and are shown in Fig. 1, together with the adsorption isotherm. The linearity of the experimental data in Fig. 1 indicates that the Langmuir-type model is reasonable for Mg adsorption, with an affinity constant of 0.74ml/pmol, a maximum number of sites at 31.19 pmol/g, and correlation coefficient of 0.979. The effects of Mg and of F on the hydrolysis rates of octacalcium phosphate at pH 7.5 and 37°C are shown in Figs 2 and 3 as the variations in the volume of 0.18 N KOH titrated with time. The control is the hydrolysis of 0.87 g of octacalcium phosphate at pH 7.5 and 37°C and is shown as curve 1 in Fig. 2; the hydrolysis rate in the first 100 h is also shown as curve 1 in Fig. 3. Figure 2 compares the effects of 1 mmol/l of MgCl, on 0.89 g of octacalcium phosphate slurry at 37°C (curve 2) with the control (curve 1). Figure 3
Hydrolysis of octacalcium phosphate
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Table 1. Adsorption isotherm of magnesium on octacalcium phosphate (OCP) at 23°C Amount Concentration
“&ICI,
of Mg
Equilibrium C @mol/ml)
Initial OtmoW) 0.016 0.038 0.086 1.64 1.59 2.45 3.76 4.91 4.91 5.89 7.36 7.36 7.85
0.012 0.025 0.054 1.13 1.28 1.66 3.33 3.56 3.93 4.88 5.84 6.20 6.57
Adsorbate adsorbed m (mg/g of CCP) Ocmol/g) 0.274 0.629 1.71 14.43 17.38 14.26 22.14 22.44 19.52 23.55 28.68 23.95 28.02
OCP (g) 0.092 0.105 0.104 0.212 0.107 0.221 0.101 0.241 0.202 0.215 0.212 0.194 0.228
compares the effects of 0.42 g CaF, on the hydrolysis of 0.87 g of octacalcium phosphate slurry in the presence (curve 3) and absence (curve 2) of 1 mmol
MgCl, together with the control (curve 1). The hydrolysis of octacalcium phosphate alone or in the presence of CaF, (Fig. 3 and curve 1 in Fig. 2) appeared to take place in two stages: an initial rapid process, which consumed 0.35 ml of KOH in the first 4 h in the control (as indicated by arrow) and in 0.5 h in the presence of CaF,, followed by the main, slower process with a typical sigmoidal curve. The initial rapid process was attributed to the surface reaction on octacalcium phosphate; the F accelerated and the adsorption of Mg prevented this process. Therefore, only the slower process (with sigmoidal curve) was observed in the presence of Mg ions and without F. The sigmoidal curve (which is above the 0.35 ml of KOH baseline in the case of the initial rapid process) can be divided ideall;y into three regions that usually
*
YO
100
2
3
Concentration
4
5
6
500
600
700
600
Time (h) Fig. 2. The effect of Mg on hydrolysis of octacalcium phosphate as indicated by consumption of KOH as a function of time at pH 7.5. Curve 1: octacalcium phosphate alone at 37°C; the arrow indicates the end of the initial surface reaction. Curve 2: octacalcium phosphate in the presence of 1 mmol/l MgCl, at 37°C. Curve 3: octacalcium phosphate alone at 25°C.
indicate the nucleation, growth and completion processes in heterogeneous crystallization (Tung, Chow and Brown, 1985). The nucleation time (Tn), which is related reciprocally to the nucleation rate, was obtained by the intersection of the tangent at the steepest part of the curve with the abscissa at which the nucleation started (in the case of the initial rapid process, the abscissa with 0.35 ml of KOH; in the case of the sigmoidal curve only, the abscissa is the X axis). The growth rate, expressed as pmol of KOH titrated per hour and per gram of octacalcium phosphate (pmol/h - g), was obtained by the slope at the steepest part of the curve divided by the amount of octacalcium phosphate. Table 2 and Fig. 2 show that Mg ions, as 1 mmol/l of MgClr, decreased the rate of nucleation dramatically, as indicated by the increase in the nucleation time by a factor of 7.7, and the rate of crystal growth only slightly, as indicated by the slope and KOH titration rate. Table 2 and Fig. 3 show that the F, as CaF, solid in the solution, decreased the nucleation time and increased the crystal growth rates by factors of 9.4 and 5.4, respectively. When both Mg and F ions were present, the
*
0
1
3OO’n400
200
7
(pmollml)
Fig. 1. Adsorption isotherm of Mg on octacalcium phosphate at 23°C: (*) isotherm; (A) Langmuir plot. The straight line was obtained by linear regression of the Langmuir model,
’ 10
20
30
40
50
60
70
60
90
100
Time (h) Fig. 3. The effect of Mg and F on the OCP hydrolysis as indicated by consumption of KOH as a function of time at 37°C and pH 7.5. Curve 1: OCP alone. Curve 2: OCP in the presence of CaF, solid. Curve 3: OCP in the presence of 1 mm01 MgCl, and CaF,.
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Table 2. Effect of magnesium and fluoride on the nucleation time and crystal growth rate of hydroxyapatite during hydrolysis of octacalcium phosphate
Control MgCl, (1 mmol/l) CaF, (as solid) MgCl, + CaF,
Nucleation time fh)
Growth rate (umol/h . d
41 362 5 21
3.57 3.31 19.2 16.2
time was between the control and that of octacalcium phosphate in the presence of CaF, alone, owing to the countervailing effects of Mg and F ions; yet the growth rate was only slightly less than that of octacalcium phosphate in the presence of CaF, alone. The final hydrolysis products obtained at 37°C were relatively well-crystallized, Ca-deficient apatites, as shown by X-ray diffraction patterns (Fig. 4), with a Ca/P ratio of 1.53 as compared to 1.67 for stoichiometric apatite. The line broadening, /Ii, which reflects the crystal size and/or strain, was 0.23 + 0.01 along the c axis obtained by the (002) diffraction at 26 = 25.88”; and 0.63 + 0.01 along the a and b axes obtained by the (310) diffraction at 26 = 39.82”. The CaF, solid in the solution did not significantly affect the crystal size and/or strain or the Ca/P ratios of the final apatitic products. The 1 mmol/l of MgCl, did induction
not affect the crystal size and/or strain along the c axis or the Ca/P ratios of the tial products, but decreased the crystal size and/or strain slightly along the a and b axes with /I+of 0.65 at (310) diffraction. At 25°C and pH 7.5, the hydrolysis of octacalcium phosphate was very slow and ceased within 80 days, with products containing unreacted octacalcium phosphate and apatite; the hydrolysis for the first 800 h is shown as curve 3 in Fig. 2. During the 6rst 3 h, the slurry was saturated with octacalcium phosphate, with no detectable hydrolysis. This contirms that our Mg adsorption study at 23°C was not influenced by the hydrolysis. Higher temperature (37°C) increased the rate dramatically. Therefore, higher temperature was used to accelerate the reaction of octacalcium phosphate in the presence of Mg. Table 3 gives the compositions of solutions and solids during the hydrolysis of octacalcium phosphate in the presence of Mg at 6O”C, together with control data at 50°C. It shows that equilibrium was reached within 1.5 h both for control samples and for samples with added Mg ions. This indicates that the dissolution is a fast process, as reported by Tung et al. (1988). At 50°C and in the absence of Mg, hydrolysis of octacalcium phosphate was complete in 26 h. The Ca/P ratios in solutions decreased with time, indicating the hydrolysis of octacalcium phosphate, which released the phosphate into the solution. The higher values of
Fig. 4. X-ray powder diffraction patterns of products of hydrolysis of o&calcium phosphate at 37°C and pH 7.5. Curve 1: octacalcium phosphate in the presence of 1 mM MgCI,. Curve 2: o&calcium phosphate alone.
Hydrolysis of octacalcium phosphate
589
Table 3. Composition of solution and solid during hydrolysis of octacalcium phosphate in the presence of MgCl, at 50 and 60°C Reaction Initial concentration of M&I, (mmol) 0 0.2 2
Time h
Temperature “C
0.5 1.7 26 1.5 24 1.3 24
50 60 60
Solution Ca (mmol/l) 0.11 0.08 0.02 0.183 0.09 0.22 0.15
P (mmol/l) 0.161 0.189 0.693 0.172 0.135 0.176 0.142
WP 0.68* 0.42 0.03 1.06* 0.67 1.25’ 1.06
Solid (Mg/Ca) 0 0.003 0.009 0.009 0.012
*The solution was saturated with respect to octacalcium phosphate [for calculation, see Tomazic et al. (1989)]. the Ca/P ratio in the Mg solutions demonstrate the inhibitory effect of Mg on the hydrolysis, and the hydrolysis of octacalcium phosphate was not complete after 24 h. The solids contained small amounts of Mg when the initial solutions contained Mg ions. The Mg content in the solid increased with time and the initial Mg concentration in the solutions, indicating uptake of Mg during hydrolysis. Although whitlockite was not detected in the final apatitic products by the X-ray diffraction powder pattern, its presence was not ruled out wben the Mg was in the solution. DEXXJSSION
Hydrolysis of well-crystallized octacalcium phosphate is very slow and involves two processes: a rapid dissolution of octacalcium phosphate crystals and a slow formation of apatitic products [which have been referred to as calcium-deficient apatite, non-stoichiometric apatite, tricalcium phosphate hydrate, or octacalcium phosphate-hydrolysate (OCPH)]. Because the dissolution of octacalcium phosphate is rapid and equilibrium is reached in 30 min (Tung et al., 1988; also Table 3), it is not the rate-controlling process. Therefore, the observed effects of F, Mg and temperature on the hydrolysis are not due to the effects on the dissolution of octacalcium phosphate, but on the formation of apatite. The formation of apatitic products may occur in two ways (LeGeros er al., 1989; Tomazic et al., 1989): (i) by in situ solid-state (topotactical) transformation as calcium ions diffuse into the crystal lattice of the octacalcium phosphate and/or (ii) by precipitation from the solution. The first way occurs because of the structural similarity between the apatite and the octacalcium phosphate. We propose that the initial rapid titration of 0.35 ml KOH is due to this topotactical transformation on the surface. Although the kinetic phenomena of this solid-solid transformation are unknown, it is expected to be slow inside the large crystals and fast on the surface of the crystals in the presence of the calcium and absence of the inhibitor. As shown in Fig. 3, this transformation is accelerated by the F and inhibited by the Mg. The second way occurs because the apatite is less soluble than the octacalcium phosphate under the conditions studied. These precipitation kinetics, indicated by the sigmoidal titration curves after the initial titration, are consistent with those of a mechanism based on precipitation from solution involving heterogeneous crystallization that
is affected by the Mg, F and temperature. The effect of F is due to the decrease in the solubility of fluoridated hydroxyapatite as compared to hydroxy apatite (Moreno, Kresak and Zahradnik, 1977) and, therefore, the increase in the thermodynamical driving force to form the F-containing apatite. Also, the presence of F in the solution that has a low degree of supersaturation with respect to hydroxyapatite increases the rate of direct precipitation of apatite (Moreno et al., 1977). The effect of Mg is probably due to the blocking of the active growth sites through adsorption of Mg, as suggested by Eanes and Rattner (1981). This adsorption of Mg on forming nuclei increases the nucleation time until the numbers of growth sites are large enough to overcome the effect of Mg ions, implying that Mg ions are taken up and their effect decreases with time. The above mechanisms that caused F to increase and Mg to decrease the nucleation and crystal growth processes seem to be operative in the hydrolysis and maturation of other calcium phosphates (LeGeros et al., 1980, 1982; Tung and Brown, 1985). These opposite effects are probably related to the cariostatic and cariogenic effects. F promotes the hydrolysis of acidic calcium phosphates and is readily incorporated into the structure with lower solubility (LeGeros et al., 1982). Mg slows down the hydrolysis of acidic calcium phosphates and therefore increases the acidic calcium phosphate content and solubility in the enamel (LeGeros et al., 1980). These results also offer a mechanism explaining observations in vivo that fluoride reduces and Mg increases the amount of acidic calcium phosphate in rat enamel (Brown et al., 1987) and that the increase in Mg uptake delays enamel mineralization, with line broadening on X-ray diffraction of the enamel (Spencer et al., 1989). Therefore, these in vitro observations may be extrapolated to biological systems where the calcified tissues precipitate in a medium relatively rich in Mg ions (Wuthier, 1977) and contain a trace amount of Mg (Shaw and Yen, 1972; Le Geros, 1984). The Mg ions appear to be involved in the regulation of mineralization with an effect opposite to that of fluoride. It is also pertinent that hydrolysis of octacalcium phosphate is very slow compared to that of dicalcium phosphate dihydrate (Tung et al., 1985) or of the octacalcium phosphate precursor during the precipitation of hydroxyapatite from solutions containing physiological concentrations of calcium and phosphate (Tomson and Nancollas, 1978). At 37”C, the
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M. S. TIJNG
hydrolysis of octacalcium phosphate at pH 7.5 has a nucleation time of 47 h and a growth rate of 3.57 pmol/h *g (Table 2), and the hydrolysis of dicalcium phosphate dihydrate at pH 7.4 a nucleation time of 1 h and a growth rate of 429 pmol/h *g (47 and 120 times faster than that of octacalcium phosphate, respectively); and in the hydrolysis of the octacalcium phosphate precursor, one in every three molecules of the precursor transforms into hydroxyapatite within minutes of precipitation from solutions (Tomson and Nancollas, 1978). This is due to the slow direct formation rate of hydroxyapatite compared to the fast rate of formation of octacalcium phosphate, and the fast, in situ, one layer at a time, transformation of the precursor. This in situ, one layer at a time, surface transformation is considered to be a topotactical solid-solid transformation, as discussed above. Hydrolysis of well-crystallized octacalcium phosphate, without going through the precursor phase, involves mainly the former slow direct formation of hydroxyapatite. Hydrolysis of dicalcium phosphate dihydrate and precipitation from solutions containing physiological concentrations of calcium and phosphate involve the latter fast formation of octacalcium phosphate and subsequent in situ, layer-bylayer, topotactical transformation to hydroxyapatite. In the former direct precipitation, well-crystallized hydroxyapatite is formed, and in the latter poorly crystallized hydroxyapatite is formed because in situ conversion incorporates the imperfections and strains. Acknowledgements-This investigation was supported, in part, by USPHS Research Grant DE08916 to the American Dental Association Health Foundation from the National Institutes of Health-National Institute of Dental Research and is part of the dental research program conducted by the National Institute of Standards and Technology in cooperation with the American Dental Association Health Foundation.
REFERENCES Adamson A. W. (1960) Physical Chemistry of Surface, p. 574. Interscience, New York. Amjad Z., Koutsoukos P. G. and Nancollas G. H. (1984) The crystallization of hydroxyapatite and fluorapatite in the presence of magnesium ions. J. Coil. Interface Sci. 101, 250-256.
Boskey A. L. and Posner A. S. (1974) Magnesium stabilization of amorphous calcium phosphate: a kinetic study. Mater. Res. BUN. 9, 907-916.
Brabson J. A., Dunn R. L., Epps E. A., Hoffman W. M. and Jacob K. D. (1958) Report on phosphorus in fertilizers: nhotometric determination of total nhosphorus. J. Ass. bff. Analyt. Chem. 41, 517-524. _ Brown W. E.. Smith J. P.. Lehr J. R. and Frazier A. W. (1962) Octacalcium phosphate and hydroxyapatite. Nature l%,
1048-1055.
Brown W. E., Eidelman N. and Tomazic B. (1987) Octacalcium phosphate as a precursor in biomineral formation. Adv. dent. Res. 1, 306-313. Cheng P.-T., Grabher J. J. and LeGeros R. Z. (1988) Effects of magnesium on calcium phosphate formation. Magnesium J. 7, 123-132. Chickerur N. S., Tung M. S. and Brown W. E. (1980) A mechanism for incorporation of carbonate into apatite. Calc. Tiss. Znt. 37, 194-197.
et
al.
Driessens F. C. M. and Woltgens J. H. M. (Eds) (1988) Tooth Develonment and Cartes. CRC Press. Boca Raton. FL. Eanes E. D. and Meyer J. I. (1978) The effect of fluoride on apatite formation from unstable supersaturated solutions at pH 7.4. J. dent. Res. 60, 1719-1723. Eanes E. D. and Rattner S. L. (1981) The effect of magnesium on apatite formation in seeded supersaturated solutions at pH 7.4. J. dent. Res. 60, 1719-1723. Hallsworth A. S., Robinson C. and We&here11 J. A. (1972) Mineral and magnesium distribution within the approxima1 carious lesion of dental enamel. Caries Res. 6, 156168.
Klug H. P. and Alexander L. E. (1974) X-ray Diffraction Procedures for Poiycrystalline and Amorphous Materials,
2nd edn. Wiley, New York. Leach S. A. Ed. (1986) Factors Relating to Demineralisation and Remineralisation of the Teeth. IRL Press, Oxford. LeGeros R. Z. (1974) Variations in the crystalline components of human dental calculus: I. Crystallographic and spectroscopic method of analyses. J. dent. Res. 53,45-50. LeGeros R. Z. (1984) Incorporation of magnesium in synthetic and in biological apatites. In Tooth Enamel IV (Eds Fearnhead R. W. and Suga S.), pp. 32-36. Elsevier Science, Amsterdam. LeGeros R. Z. (1985) Preparation of octacalcium phosphate (OCP): a direct fast method. Calc. Tiss. Znt. 37. 194-197. LeGeros R. Z. (1991) Calcium Phosphates in Oral Biology and Medicine, Chap. 4, p. 68. S. Karger, Basel. LeGeros R. Z., Taheri M. H., Quirolgico G. M. and LeGeros J. P. (1980) Formation and stability of apatites: effects of some cationic substituents. Proc. 2hd Znt. Congr. Phosphorus Compounds, pp. 89-103. Boston, MA. LeGeros R. Z., Singer L., Ophaug R. and Quirlogico G. (1982) The effect of fluoride on the stability of synthetic and biological (bone mineral) apatites. In- Osteaporosis (Eds Menczel J.. Robin G. C. and Makin M.). pp. 327-341. Wiley, New York. LeGeros R. Z., Kijkowska R. and LeGeros J. P. (1984) Formation and transformation of octacalcium phosphate, OCP: a preliminary report. Scann. Elect. Microsc. IV, 1771-1777. LeGeros R. Z., Daculsi G., Orly I. and Abergas T. (1989) Solution-mediated transformation of octacalcium phosphate (OCP) to apatite. S.E.M. 3, 129-138. McDowell H., Brown W. E. and Sutter J. R. (1971) Solubility study of calcium hydrogen phosphate: ion pair formation. Inorganic Chem. 10, 1638-1643. MacGregor J., Robertson, W. G. and Nordin B. E. C. (1965) Octacalcium phosphate: the governing phase in the solubility equilibrium in apatite calculi. Brit. J. Vrol. 37, I.
518-525.
Moreno E. G., Kresak M. and Zahradnik R. T. (1977) Physicochemical aspects of fluoride-apatite systems relevant to the study of dental caries. Caries Res. 11 (suppl I), 142-171. Salimi M. H., Heughebaert J. C. and Nancollas G. H. (1985) Crystal growth of calcium phosphate in the presence of magnesium ions. Langmuir. 1, 119-122. Shaw J. H. and Yen P. K. J. (1972) Sodium, potassium, and magnesium concentrations in the enamel and dentin of human and rhesus monkey teeth. J. dent. Res. 51, 95-100. Spencer P., Barnes C., Martini J., Garcia R., Elliott C. and Doremus R. (1989) Incornoration of magnesium into rat dental enamel and its inguence on crys~llization. Archs oral Biol. 34, 767-771.
Tomazic B. B., Etz E. S. and Brown W. E. (1987) Nature and properties of cardiovascular deposits. S.E.M. 1, 95-10s.
Tomazic B. B., Tung M. S., Gregory T. M. and Brown W. E. (1989) Mechanism of hydrolysis of octacalcium phosphate. Scanning Microsc. 3, 119-127.
Hydrolysis of octacalcium phosphate Tomson M. B. and Nancollas G. H. (1978) Mineralization kinetics: a constant composition approach. Science 280, 1059-1060. Tung M. S. and Brown W. E. (1985) The role of octacalcium phosphate in subcutaneous heterotopic calcification. Calc. Tbs. Int. 37, 329-331. Tung M. S., Chow L. C. and Brown W. E. (1985) Hydrolysis of dicalcium phosphate in the presence or absence of calcium fluoride. J. dent. Res. 64, 2-5.
591
Tung M. S., Eidelman N., Sieck B. and Brown W. E. (1988) Gctacalcium phosphate solubility product from 4 to 37°C. J. natn. Inst. Stand. Tech. 93, 613-624. Willis J. B. (1960) The determination of metals in blood serum by atomic absorption spectroscopy. I. Calcium. Spectrochim Acta. 16, 259. Wuthier R. E. (1977) Electrolytes of isolated epiphyseal chondrocytes, matrix vesicles, and extracellular fluids. Calc. Tiss. Res. 23, 125-133.
0003~9969/92 $5.00 + 0.00
1992
Pergamon Press Ltd
THE EFFECTS OF MAGNESIUM AND FLUORIDE ON THE HYDROLYSIS OF OCTACALCIUM PHOSPHATE* M. S. TUNG, B. TOMAZICand W. E. BROWN American Dental Association Health Foundation Paffenbarger Research Center, National Institute of Standards and Technology, Gaithersburg, MD 20899, U.S.A. (Accepted 4 February 1992) Summary-The adsorption of Mg ions on octacalcium phosphate (GCP) and its effect on OCP hydrolysis, with and without F, were studied. The Mg adsorption isotherm was fitted by the Langmuir model with an affinity constant of 0.74 ml/pmol and maximum number of sites, 31.19 pmol/g. The hydrolysis rates were measured in a pH stat by titration of base and were strongly temperature dependent. The products were examined by X-ray diffraction and chemical analysis. GCP hydrolysis takes place in two stages: the fast initial process, which is attributed to the surface topotactical conversion, followed by the main, slower process, which involves the nucleation and crystal growth. Mg ions, as 1 mmol/l MgCl,, prevented the initial surface reaction and decreased the nucleation rate dramatically and the growth rate slightly; F increased the rates of surface reaction and both the nucleation and crystal growth processes. The Ca/P ratio (1.53) and the line broadening in the X-ray diffraction patterns of the apatitic products were not significantly affected by the F. Mg also did not affect the Ca/P ratio and the line broadening at (002) diffraction, but decreased the line broadening at (310) diffraction. Key words: adsorption, biomineralization,
hydrolysis, fluoride, magnesium, octacalcium phosphate.
INTRODUCITON To understand better the effects of F and Mg ions on the formation and maturation of calcified tissues, these effects have been studied in systems of synthetic apatite formation involving precursors such as amorphous calcium phosphate (Boskey and Posner, 1974; Cheng, Grabher and LeGeros, 1988), dicalcium phosphate dihydrate (Tung and Brown, 1985) and octacalcium phosphate (Ca,H,(PO,), .5H,O) (Le Geros, 1991). F and Mg have opposite effects on the formation and maturation of calcium phosphates (Brown et al., 1962; Amjad, Koutsoukos and Nancollas, 1984) and enamel (Brown, Eidelman and Tomazic, 1987). Fluoride has been shown to promote the formation of apatite and the hydrolysis of dicalcium phosphate dihydrate (Tung and Brown, 1985); it greatly curtails or eliminates the appearance of the octacalcium phosphate precursor in the spontaneous precipitation of apatite (Eanes and Meyer, 1978) as well as the formation of dental caries (reviewed in Leach, 1986; Driessens and Woltgens, 1988). On the other hand, Mg ions have been implicated in the early stages of caries (Hallsworth, Robinson and Weatherell, 1972). These ions retard the growth rates of hydroxyapatite, fluorapatite and octacalcium phosphate (Amjad et al., 1984; Salimi, Heughebaert and Nancollas, 1985) and also lengthen the lifetime of the *Certain commercial materials and equipment are identified in this paper to specify the experimental procedure. In no instance does such identification imply recommendation or endorsement by the National Institute of Standards and Technology or the ADA Health Foundation or that the material or equipment identified is necessarily the best available for the purpose.
precursors, amorphous calcium phosphate (Boskey and Posner, 1974; Cheng et al., 1988) and octacalcium phosphate, during the spontaneous precipitation of apatite (Eanes and Meyer, 1978). It has been suggested that these effects of Mg ions on the formation of calcium phosphates are due to the blocking of active growth sites through adsorption of Mg ions at the crystal surfaces (Eanes and Rattner, 1981; Amjad et al., 1984). Our purposes now were to study the adsorption of Mg ions on synthetic, well-crystallized octacalcium phosphate and the effects of F and Mg ions on the hydrolysis of octacalcium phosphate at physiological pH. Octacalcium phosphate is structurally similar to hydroxyapatite and its hydrolysis does not further involve an intermediate precursor. Formation and transformation of octacalcium phosphate have been reported under conditions similar to those in physiological environments (Tomson and Nancollas, 1978; LeGeros, Kijkowska and LeGeros, 1984). Octacalcium phosphate is also frequently present as one of the crystalline components of dental calculi (Le Geros, 1974) and, on occasions, is associated with urinary stones (MacGregor, Robertson and Nordin, 1965) and other pathological calcifications (Tung and Brown, 1985). It is believed to be a precursor in the formation of calcified tissues such as bone, teeth and pathological deposits (Brown et al., 1987; Tomazic, Etz and Brown, 1987). Therefore, hydrolysis of octacalcium phosphate provides a model for maturation processes in biomineralization. MATERIALSANDMETHODS Reagent-grade chemicals and distilled water were used throughout the experiments. Octacalcium
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phosphate was prepared by two methods. One method consisted of the hydrolysis of dicalcium phosphate dihydrate, as reported by Chickerur, Tung and Brown (1980). G&calcium phosphate prepared in this way was used throughout unless specified otherwise. It has a Ca/P ratio of 1.35 + 0.02 and a surface area of 25.16 m2/g as determined by the N adsorption BET method. The crystals of this preparation yield a characteristic X-ray diffraction pattern for octacalcium phosphate and appear as a homogeneous phase under the polarizing microscope. The second method involved the direct crystallization as reported by LeGeros (1985) with minor modifications; the phosphate so prepared is designated as octacalcium phosphate-L (Tomazic et al., 1989). Adsorption measurements
5 Ca,H,(PO,),
.5&o
+ 8 Ca,(PO&OH
Accurately weighed samples of octacalcium phosphate (0.1-0.5 g) were transferred to plastic bottles, and 2 ml of water were added. The slurry was shaken for 2 h and was saturated with octacalcium phosphate with no detectable hydrolysis at 23°C (Tung et al., 1988). A Mg-containing solution was then added, with final solution volume of 4-6 ml. The adsorption equilibration was carried out in the shaker for 1 h at 23°C. (A preliminary study had indicated that the adsorption reached equilibrium in 5 min, with no adsorption on the walls of the plastic bottles, and the pH of the slurries was 7.5, independent of the Mg concentration in the slurry.) After 1 h, the slurry was filtered through a 0.22 pm filter (Millipore, Bedford, MA, U.S.A.), and the concentration of Mg in the filtrate was analysed by atomic absorption spectroscopy. The difference between the initial and filtrate Mg concentrations is the amount adsorbed on the surface of the octacalcium phosphate crystals. The adsorption isotherm was obtained from the amounts of Mg adsorbed as a function of equilibrium concentration. A general interpretation of the isotherm is achieved by the use of the Langmuir adsorption model (Adamson, 1960) described by C/m = l/Km, + C/m,
in solution. F was added as solid CaF,. We used CaF, because it was the main product in the topical fluoridation; it released the F and Ca slowly as the reaction proceeded and it dissolved. The released Ca partially compensated the released phosphate during hydrolysis. The slurry was stirred by a magnetic bar at approx. 100 rev/mm. Several of the hydrolysis reactions were made at 50 and 60°C at pH 7.4, with 1 g of octacalcium phosphate/l to study the Mg incorporation. Periodic samplings of solutions and solids were done to determine the progression of hydrolysis and incorporation. The hydrolysis of octacalcium phosphate produces phosphoric acid by the following reaction:
(1)
where C @mol/ml) is the equilibrium concentration of adsorbate (Mg) in the solution, m bmol/g) is the amount of adsorbate and m, (pmol/g) is the maximum amount (maximum number of sites) adsorbed per gram of substrate (octacalcium phosphate), and K, the affinity constant, is a constant reflecting the affinity that the adsorbate molecules have for the adsorption sites. The Langmuir plot, C/m versus C, should be a straight line where the slope and intercept give the information about the K and mOparameters. The correlation coefficient of the least-squares linear regression was calculated to indicate the difference between the Langmuir model and the experimental data. Hydrolysis A weighed sample of octacalcium phosphate (0.87-0.89 g) and 35 ml of water (25 g of octacalcium phosphate/l) with or without Mg and/or F in a 120 ml polypropylene bottle was stirred (Mg and/or F were added before the phosphate), and maintained at pH 7.5 and 37 or 25°C. Mg was added as 1 mmol/l MgCl,
+ 6 H,PO, + 17 H,O.
(2)
The rate of hydrolysis was followed by titration with 0.18 mol/l KOH as a function of time in a pH stat. The pH-stat system included a METROHM pH meter 632, Impulsomat 614 and the titrator, Dosimat 655 (Brinkmann Instruments Inc., Westbury, NY, U.S.A.). The pH measurements were made with a glass/calomel combination electrode, which was calibrated daily. The reaction was stopped when the consumption of KOH ceased. The solids were filtered, washed with water, and dried at room temperature. The filtered solutions were analysed for Ca (Willis, 1960; McDowell, Brown and Sutter, 1971) and phosphate (Brabson et al., 1958). Because there is a net release of phosphate during hydrolysis of octacalcium phosphate, the decrease in solution Ca/P ratio is an indicator of the progression in the hydrolysis. The X-ray diffraction analysis of the solids were made with a Norelco Philips unit with Ni-filtered CuKa radiation. The crystallite size and/or crystal strain of the solid reaction products was assessed in terms of line broadening at half-peak heights, /3;, of the 002 and 310 reflections in the apatite (Klug and Alexander, 1974). RESULTS
Table 1 summarizes the results for the adsorption of Mg on octacalcium phosphate from aqueous solution at 23°C with the initial concentrations of Mg ranging from 0.016 to 7.85 mmol/l. The data in Table 1 were plotted according to the Langmuir equation and are shown in Fig. 1, together with the adsorption isotherm. The linearity of the experimental data in Fig. 1 indicates that the Langmuir-type model is reasonable for Mg adsorption, with an affinity constant of 0.74ml/pmol, a maximum number of sites at 31.19 pmol/g, and correlation coefficient of 0.979. The effects of Mg and of F on the hydrolysis rates of octacalcium phosphate at pH 7.5 and 37°C are shown in Figs 2 and 3 as the variations in the volume of 0.18 N KOH titrated with time. The control is the hydrolysis of 0.87 g of octacalcium phosphate at pH 7.5 and 37°C and is shown as curve 1 in Fig. 2; the hydrolysis rate in the first 100 h is also shown as curve 1 in Fig. 3. Figure 2 compares the effects of 1 mmol/l of MgCl, on 0.89 g of octacalcium phosphate slurry at 37°C (curve 2) with the control (curve 1). Figure 3
Hydrolysis of octacalcium phosphate
587
Table 1. Adsorption isotherm of magnesium on octacalcium phosphate (OCP) at 23°C Amount Concentration
“&ICI,
of Mg
Equilibrium C @mol/ml)
Initial OtmoW) 0.016 0.038 0.086 1.64 1.59 2.45 3.76 4.91 4.91 5.89 7.36 7.36 7.85
0.012 0.025 0.054 1.13 1.28 1.66 3.33 3.56 3.93 4.88 5.84 6.20 6.57
Adsorbate adsorbed m (mg/g of CCP) Ocmol/g) 0.274 0.629 1.71 14.43 17.38 14.26 22.14 22.44 19.52 23.55 28.68 23.95 28.02
OCP (g) 0.092 0.105 0.104 0.212 0.107 0.221 0.101 0.241 0.202 0.215 0.212 0.194 0.228
compares the effects of 0.42 g CaF, on the hydrolysis of 0.87 g of octacalcium phosphate slurry in the presence (curve 3) and absence (curve 2) of 1 mmol
MgCl, together with the control (curve 1). The hydrolysis of octacalcium phosphate alone or in the presence of CaF, (Fig. 3 and curve 1 in Fig. 2) appeared to take place in two stages: an initial rapid process, which consumed 0.35 ml of KOH in the first 4 h in the control (as indicated by arrow) and in 0.5 h in the presence of CaF,, followed by the main, slower process with a typical sigmoidal curve. The initial rapid process was attributed to the surface reaction on octacalcium phosphate; the F accelerated and the adsorption of Mg prevented this process. Therefore, only the slower process (with sigmoidal curve) was observed in the presence of Mg ions and without F. The sigmoidal curve (which is above the 0.35 ml of KOH baseline in the case of the initial rapid process) can be divided ideall;y into three regions that usually
*
YO
100
2
3
Concentration
4
5
6
500
600
700
600
Time (h) Fig. 2. The effect of Mg on hydrolysis of octacalcium phosphate as indicated by consumption of KOH as a function of time at pH 7.5. Curve 1: octacalcium phosphate alone at 37°C; the arrow indicates the end of the initial surface reaction. Curve 2: octacalcium phosphate in the presence of 1 mmol/l MgCl, at 37°C. Curve 3: octacalcium phosphate alone at 25°C.
indicate the nucleation, growth and completion processes in heterogeneous crystallization (Tung, Chow and Brown, 1985). The nucleation time (Tn), which is related reciprocally to the nucleation rate, was obtained by the intersection of the tangent at the steepest part of the curve with the abscissa at which the nucleation started (in the case of the initial rapid process, the abscissa with 0.35 ml of KOH; in the case of the sigmoidal curve only, the abscissa is the X axis). The growth rate, expressed as pmol of KOH titrated per hour and per gram of octacalcium phosphate (pmol/h - g), was obtained by the slope at the steepest part of the curve divided by the amount of octacalcium phosphate. Table 2 and Fig. 2 show that Mg ions, as 1 mmol/l of MgClr, decreased the rate of nucleation dramatically, as indicated by the increase in the nucleation time by a factor of 7.7, and the rate of crystal growth only slightly, as indicated by the slope and KOH titration rate. Table 2 and Fig. 3 show that the F, as CaF, solid in the solution, decreased the nucleation time and increased the crystal growth rates by factors of 9.4 and 5.4, respectively. When both Mg and F ions were present, the
*
0
1
3OO’n400
200
7
(pmollml)
Fig. 1. Adsorption isotherm of Mg on octacalcium phosphate at 23°C: (*) isotherm; (A) Langmuir plot. The straight line was obtained by linear regression of the Langmuir model,
’ 10
20
30
40
50
60
70
60
90
100
Time (h) Fig. 3. The effect of Mg and F on the OCP hydrolysis as indicated by consumption of KOH as a function of time at 37°C and pH 7.5. Curve 1: OCP alone. Curve 2: OCP in the presence of CaF, solid. Curve 3: OCP in the presence of 1 mm01 MgCl, and CaF,.
M. S. TUNG et al.
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Table 2. Effect of magnesium and fluoride on the nucleation time and crystal growth rate of hydroxyapatite during hydrolysis of octacalcium phosphate
Control MgCl, (1 mmol/l) CaF, (as solid) MgCl, + CaF,
Nucleation time fh)
Growth rate (umol/h . d
41 362 5 21
3.57 3.31 19.2 16.2
time was between the control and that of octacalcium phosphate in the presence of CaF, alone, owing to the countervailing effects of Mg and F ions; yet the growth rate was only slightly less than that of octacalcium phosphate in the presence of CaF, alone. The final hydrolysis products obtained at 37°C were relatively well-crystallized, Ca-deficient apatites, as shown by X-ray diffraction patterns (Fig. 4), with a Ca/P ratio of 1.53 as compared to 1.67 for stoichiometric apatite. The line broadening, /Ii, which reflects the crystal size and/or strain, was 0.23 + 0.01 along the c axis obtained by the (002) diffraction at 26 = 25.88”; and 0.63 + 0.01 along the a and b axes obtained by the (310) diffraction at 26 = 39.82”. The CaF, solid in the solution did not significantly affect the crystal size and/or strain or the Ca/P ratios of the final apatitic products. The 1 mmol/l of MgCl, did induction
not affect the crystal size and/or strain along the c axis or the Ca/P ratios of the tial products, but decreased the crystal size and/or strain slightly along the a and b axes with /I+of 0.65 at (310) diffraction. At 25°C and pH 7.5, the hydrolysis of octacalcium phosphate was very slow and ceased within 80 days, with products containing unreacted octacalcium phosphate and apatite; the hydrolysis for the first 800 h is shown as curve 3 in Fig. 2. During the 6rst 3 h, the slurry was saturated with octacalcium phosphate, with no detectable hydrolysis. This contirms that our Mg adsorption study at 23°C was not influenced by the hydrolysis. Higher temperature (37°C) increased the rate dramatically. Therefore, higher temperature was used to accelerate the reaction of octacalcium phosphate in the presence of Mg. Table 3 gives the compositions of solutions and solids during the hydrolysis of octacalcium phosphate in the presence of Mg at 6O”C, together with control data at 50°C. It shows that equilibrium was reached within 1.5 h both for control samples and for samples with added Mg ions. This indicates that the dissolution is a fast process, as reported by Tung et al. (1988). At 50°C and in the absence of Mg, hydrolysis of octacalcium phosphate was complete in 26 h. The Ca/P ratios in solutions decreased with time, indicating the hydrolysis of octacalcium phosphate, which released the phosphate into the solution. The higher values of
Fig. 4. X-ray powder diffraction patterns of products of hydrolysis of o&calcium phosphate at 37°C and pH 7.5. Curve 1: octacalcium phosphate in the presence of 1 mM MgCI,. Curve 2: o&calcium phosphate alone.
Hydrolysis of octacalcium phosphate
589
Table 3. Composition of solution and solid during hydrolysis of octacalcium phosphate in the presence of MgCl, at 50 and 60°C Reaction Initial concentration of M&I, (mmol) 0 0.2 2
Time h
Temperature “C
0.5 1.7 26 1.5 24 1.3 24
50 60 60
Solution Ca (mmol/l) 0.11 0.08 0.02 0.183 0.09 0.22 0.15
P (mmol/l) 0.161 0.189 0.693 0.172 0.135 0.176 0.142
WP 0.68* 0.42 0.03 1.06* 0.67 1.25’ 1.06
Solid (Mg/Ca) 0 0.003 0.009 0.009 0.012
*The solution was saturated with respect to octacalcium phosphate [for calculation, see Tomazic et al. (1989)]. the Ca/P ratio in the Mg solutions demonstrate the inhibitory effect of Mg on the hydrolysis, and the hydrolysis of octacalcium phosphate was not complete after 24 h. The solids contained small amounts of Mg when the initial solutions contained Mg ions. The Mg content in the solid increased with time and the initial Mg concentration in the solutions, indicating uptake of Mg during hydrolysis. Although whitlockite was not detected in the final apatitic products by the X-ray diffraction powder pattern, its presence was not ruled out wben the Mg was in the solution. DEXXJSSION
Hydrolysis of well-crystallized octacalcium phosphate is very slow and involves two processes: a rapid dissolution of octacalcium phosphate crystals and a slow formation of apatitic products [which have been referred to as calcium-deficient apatite, non-stoichiometric apatite, tricalcium phosphate hydrate, or octacalcium phosphate-hydrolysate (OCPH)]. Because the dissolution of octacalcium phosphate is rapid and equilibrium is reached in 30 min (Tung et al., 1988; also Table 3), it is not the rate-controlling process. Therefore, the observed effects of F, Mg and temperature on the hydrolysis are not due to the effects on the dissolution of octacalcium phosphate, but on the formation of apatite. The formation of apatitic products may occur in two ways (LeGeros er al., 1989; Tomazic et al., 1989): (i) by in situ solid-state (topotactical) transformation as calcium ions diffuse into the crystal lattice of the octacalcium phosphate and/or (ii) by precipitation from the solution. The first way occurs because of the structural similarity between the apatite and the octacalcium phosphate. We propose that the initial rapid titration of 0.35 ml KOH is due to this topotactical transformation on the surface. Although the kinetic phenomena of this solid-solid transformation are unknown, it is expected to be slow inside the large crystals and fast on the surface of the crystals in the presence of the calcium and absence of the inhibitor. As shown in Fig. 3, this transformation is accelerated by the F and inhibited by the Mg. The second way occurs because the apatite is less soluble than the octacalcium phosphate under the conditions studied. These precipitation kinetics, indicated by the sigmoidal titration curves after the initial titration, are consistent with those of a mechanism based on precipitation from solution involving heterogeneous crystallization that
is affected by the Mg, F and temperature. The effect of F is due to the decrease in the solubility of fluoridated hydroxyapatite as compared to hydroxy apatite (Moreno, Kresak and Zahradnik, 1977) and, therefore, the increase in the thermodynamical driving force to form the F-containing apatite. Also, the presence of F in the solution that has a low degree of supersaturation with respect to hydroxyapatite increases the rate of direct precipitation of apatite (Moreno et al., 1977). The effect of Mg is probably due to the blocking of the active growth sites through adsorption of Mg, as suggested by Eanes and Rattner (1981). This adsorption of Mg on forming nuclei increases the nucleation time until the numbers of growth sites are large enough to overcome the effect of Mg ions, implying that Mg ions are taken up and their effect decreases with time. The above mechanisms that caused F to increase and Mg to decrease the nucleation and crystal growth processes seem to be operative in the hydrolysis and maturation of other calcium phosphates (LeGeros et al., 1980, 1982; Tung and Brown, 1985). These opposite effects are probably related to the cariostatic and cariogenic effects. F promotes the hydrolysis of acidic calcium phosphates and is readily incorporated into the structure with lower solubility (LeGeros et al., 1982). Mg slows down the hydrolysis of acidic calcium phosphates and therefore increases the acidic calcium phosphate content and solubility in the enamel (LeGeros et al., 1980). These results also offer a mechanism explaining observations in vivo that fluoride reduces and Mg increases the amount of acidic calcium phosphate in rat enamel (Brown et al., 1987) and that the increase in Mg uptake delays enamel mineralization, with line broadening on X-ray diffraction of the enamel (Spencer et al., 1989). Therefore, these in vitro observations may be extrapolated to biological systems where the calcified tissues precipitate in a medium relatively rich in Mg ions (Wuthier, 1977) and contain a trace amount of Mg (Shaw and Yen, 1972; Le Geros, 1984). The Mg ions appear to be involved in the regulation of mineralization with an effect opposite to that of fluoride. It is also pertinent that hydrolysis of octacalcium phosphate is very slow compared to that of dicalcium phosphate dihydrate (Tung et al., 1985) or of the octacalcium phosphate precursor during the precipitation of hydroxyapatite from solutions containing physiological concentrations of calcium and phosphate (Tomson and Nancollas, 1978). At 37”C, the
590
M. S. TIJNG
hydrolysis of octacalcium phosphate at pH 7.5 has a nucleation time of 47 h and a growth rate of 3.57 pmol/h *g (Table 2), and the hydrolysis of dicalcium phosphate dihydrate at pH 7.4 a nucleation time of 1 h and a growth rate of 429 pmol/h *g (47 and 120 times faster than that of octacalcium phosphate, respectively); and in the hydrolysis of the octacalcium phosphate precursor, one in every three molecules of the precursor transforms into hydroxyapatite within minutes of precipitation from solutions (Tomson and Nancollas, 1978). This is due to the slow direct formation rate of hydroxyapatite compared to the fast rate of formation of octacalcium phosphate, and the fast, in situ, one layer at a time, transformation of the precursor. This in situ, one layer at a time, surface transformation is considered to be a topotactical solid-solid transformation, as discussed above. Hydrolysis of well-crystallized octacalcium phosphate, without going through the precursor phase, involves mainly the former slow direct formation of hydroxyapatite. Hydrolysis of dicalcium phosphate dihydrate and precipitation from solutions containing physiological concentrations of calcium and phosphate involve the latter fast formation of octacalcium phosphate and subsequent in situ, layer-bylayer, topotactical transformation to hydroxyapatite. In the former direct precipitation, well-crystallized hydroxyapatite is formed, and in the latter poorly crystallized hydroxyapatite is formed because in situ conversion incorporates the imperfections and strains. Acknowledgements-This investigation was supported, in part, by USPHS Research Grant DE08916 to the American Dental Association Health Foundation from the National Institutes of Health-National Institute of Dental Research and is part of the dental research program conducted by the National Institute of Standards and Technology in cooperation with the American Dental Association Health Foundation.
REFERENCES Adamson A. W. (1960) Physical Chemistry of Surface, p. 574. Interscience, New York. Amjad Z., Koutsoukos P. G. and Nancollas G. H. (1984) The crystallization of hydroxyapatite and fluorapatite in the presence of magnesium ions. J. Coil. Interface Sci. 101, 250-256.
Boskey A. L. and Posner A. S. (1974) Magnesium stabilization of amorphous calcium phosphate: a kinetic study. Mater. Res. BUN. 9, 907-916.
Brabson J. A., Dunn R. L., Epps E. A., Hoffman W. M. and Jacob K. D. (1958) Report on phosphorus in fertilizers: nhotometric determination of total nhosphorus. J. Ass. bff. Analyt. Chem. 41, 517-524. _ Brown W. E.. Smith J. P.. Lehr J. R. and Frazier A. W. (1962) Octacalcium phosphate and hydroxyapatite. Nature l%,
1048-1055.
Brown W. E., Eidelman N. and Tomazic B. (1987) Octacalcium phosphate as a precursor in biomineral formation. Adv. dent. Res. 1, 306-313. Cheng P.-T., Grabher J. J. and LeGeros R. Z. (1988) Effects of magnesium on calcium phosphate formation. Magnesium J. 7, 123-132. Chickerur N. S., Tung M. S. and Brown W. E. (1980) A mechanism for incorporation of carbonate into apatite. Calc. Tiss. Znt. 37, 194-197.
et
al.
Driessens F. C. M. and Woltgens J. H. M. (Eds) (1988) Tooth Develonment and Cartes. CRC Press. Boca Raton. FL. Eanes E. D. and Meyer J. I. (1978) The effect of fluoride on apatite formation from unstable supersaturated solutions at pH 7.4. J. dent. Res. 60, 1719-1723. Eanes E. D. and Rattner S. L. (1981) The effect of magnesium on apatite formation in seeded supersaturated solutions at pH 7.4. J. dent. Res. 60, 1719-1723. Hallsworth A. S., Robinson C. and We&here11 J. A. (1972) Mineral and magnesium distribution within the approxima1 carious lesion of dental enamel. Caries Res. 6, 156168.
Klug H. P. and Alexander L. E. (1974) X-ray Diffraction Procedures for Poiycrystalline and Amorphous Materials,
2nd edn. Wiley, New York. Leach S. A. Ed. (1986) Factors Relating to Demineralisation and Remineralisation of the Teeth. IRL Press, Oxford. LeGeros R. Z. (1974) Variations in the crystalline components of human dental calculus: I. Crystallographic and spectroscopic method of analyses. J. dent. Res. 53,45-50. LeGeros R. Z. (1984) Incorporation of magnesium in synthetic and in biological apatites. In Tooth Enamel IV (Eds Fearnhead R. W. and Suga S.), pp. 32-36. Elsevier Science, Amsterdam. LeGeros R. Z. (1985) Preparation of octacalcium phosphate (OCP): a direct fast method. Calc. Tiss. Znt. 37. 194-197. LeGeros R. Z. (1991) Calcium Phosphates in Oral Biology and Medicine, Chap. 4, p. 68. S. Karger, Basel. LeGeros R. Z., Taheri M. H., Quirolgico G. M. and LeGeros J. P. (1980) Formation and stability of apatites: effects of some cationic substituents. Proc. 2hd Znt. Congr. Phosphorus Compounds, pp. 89-103. Boston, MA. LeGeros R. Z., Singer L., Ophaug R. and Quirlogico G. (1982) The effect of fluoride on the stability of synthetic and biological (bone mineral) apatites. In- Osteaporosis (Eds Menczel J.. Robin G. C. and Makin M.). pp. 327-341. Wiley, New York. LeGeros R. Z., Kijkowska R. and LeGeros J. P. (1984) Formation and transformation of octacalcium phosphate, OCP: a preliminary report. Scann. Elect. Microsc. IV, 1771-1777. LeGeros R. Z., Daculsi G., Orly I. and Abergas T. (1989) Solution-mediated transformation of octacalcium phosphate (OCP) to apatite. S.E.M. 3, 129-138. McDowell H., Brown W. E. and Sutter J. R. (1971) Solubility study of calcium hydrogen phosphate: ion pair formation. Inorganic Chem. 10, 1638-1643. MacGregor J., Robertson, W. G. and Nordin B. E. C. (1965) Octacalcium phosphate: the governing phase in the solubility equilibrium in apatite calculi. Brit. J. Vrol. 37, I.
518-525.
Moreno E. G., Kresak M. and Zahradnik R. T. (1977) Physicochemical aspects of fluoride-apatite systems relevant to the study of dental caries. Caries Res. 11 (suppl I), 142-171. Salimi M. H., Heughebaert J. C. and Nancollas G. H. (1985) Crystal growth of calcium phosphate in the presence of magnesium ions. Langmuir. 1, 119-122. Shaw J. H. and Yen P. K. J. (1972) Sodium, potassium, and magnesium concentrations in the enamel and dentin of human and rhesus monkey teeth. J. dent. Res. 51, 95-100. Spencer P., Barnes C., Martini J., Garcia R., Elliott C. and Doremus R. (1989) Incornoration of magnesium into rat dental enamel and its inguence on crys~llization. Archs oral Biol. 34, 767-771.
Tomazic B. B., Etz E. S. and Brown W. E. (1987) Nature and properties of cardiovascular deposits. S.E.M. 1, 95-10s.
Tomazic B. B., Tung M. S., Gregory T. M. and Brown W. E. (1989) Mechanism of hydrolysis of octacalcium phosphate. Scanning Microsc. 3, 119-127.
Hydrolysis of octacalcium phosphate Tomson M. B. and Nancollas G. H. (1978) Mineralization kinetics: a constant composition approach. Science 280, 1059-1060. Tung M. S. and Brown W. E. (1985) The role of octacalcium phosphate in subcutaneous heterotopic calcification. Calc. Tbs. Int. 37, 329-331. Tung M. S., Chow L. C. and Brown W. E. (1985) Hydrolysis of dicalcium phosphate in the presence or absence of calcium fluoride. J. dent. Res. 64, 2-5.
591
Tung M. S., Eidelman N., Sieck B. and Brown W. E. (1988) Gctacalcium phosphate solubility product from 4 to 37°C. J. natn. Inst. Stand. Tech. 93, 613-624. Willis J. B. (1960) The determination of metals in blood serum by atomic absorption spectroscopy. I. Calcium. Spectrochim Acta. 16, 259. Wuthier R. E. (1977) Electrolytes of isolated epiphyseal chondrocytes, matrix vesicles, and extracellular fluids. Calc. Tiss. Res. 23, 125-133.
