Caries Management - Science and Clinical Practice. Группа авторов

above shows the arrangement of calcium, phosphate, and hydroxyl groups in hydroxyapatite, as if looking down the long axis of a crystal. The outline of one unit cell, which has the formula Ca₁₀(PO₄)₆(OH)₂, is marked by a bold dashed line. The dotted triangles in middle mark groups of three Ca²⁺ ions which are rotated 60° in successive layers of the structure, forming a “channel” in which lie the OH⁻ ions (cf. Fig. 2.8). The remaining calcium ions form two columns of single ions. For clarity, the ions are not shown to scale. In reality, the oxygen atoms would be much larger and would fill most of the space within the structure. (Modified from Brown WE, et al: Ann Rev Mater Sci 1976;6:213–235.) The table compares lattice ions and impurity ions that can take their place, to show the disparities in charge and size (except for fluoride).

      The solubility of a particular form of calcium phosphate, as defined by the concentration of dissolved solid in a saturated solution, is not constant but varies according to the solution composition. The dominant factor in calcium phosphate chemistry is pH: the solubility of all calcium phosphates increases as the pH falls below 7. The solubility of hydroxyapatite increases particularly rapidly with pH, by a factor of ca. 10 per pH unit fall (Fig. 2.7).

      Fluoride and Calcium Phosphate Chemistry

      Fluoride has profound effects on solubility of hydroxyapatite and dental minerals. It is readily incorporated into the apatite structure because F⁻ ions can replace OH⁻ ions. Fluorapatite, in which all OH⁻ is replaced by F⁻, is less soluble than hydroxyapatite at pH7 or below (Fig. 2.7). In hydroxyapatite crystals, OH⁻ ions are located within channels formed by triangular groups of Ca²⁺ ions, but lie between the triangles^25,26 (A in Fig. 2.8). In fluorapatite, the slightly smaller F^- ion fits within the triangles (B in Fig. 2.8) and this is associated with a denser, more stable crystal structure. Partial substitution of F⁻ for OH⁻ produces fluorhydroxyapatites, which are thought to be stabilized by hydrogen-bonding between adjacent F⁻ and OH⁻ ions²⁴ (C in Fig. 2.8).

      Important though they are, the effects of fluoride on crystal structure are probably much less significant than the reactions which take place between crystal surfaces of apatites and fluoride dissolved in the bathing liquid. If fluorhydroxyapatite, or even pure hydroxyapatite—which of course contains no fluoride ions—is placed in an acidic solution containing low concentrations of fluoride, the rate of dissolution is lower than in one that is fluoride-free^28,29 (Fig. 2.9). This phenomenon is due to replacement of OH⁻ ions by F⁻ ions at the crystal surfaces, probably not much deeper than one unit cell thickness.³⁰ The F⁻ ions stabilize the surrounding Ca²⁺ ions.³⁰ Regions of the crystal surface in which F⁻ has replaced OH⁻ are in effect converted into fluorapatite, so are much less soluble than nonfluoridated regions.^28,31,32 When exposed to acidic conditions, nonfluoridated regions of the crystal surfaces dissolve at the rate normal for hydroxyapatite, while the fluoridated regions will dissolve more slowly or not at all. The overall rate of dissolution is therefore slower than in the absence of fluoride, and the higher the concentration of fluoride in the ambient solution the greater the effect, because a greater proportion of the crystals surfaces is converted to fluorapatite. Thus, dissolution of tooth mineral can be partly or wholly prevented by low concentrations of fluoride in the environment of the tooth. A similar effect is thought to be responsible for the reduced solubility of fluorhydroxyapatites in fluoride-free acid.³³ Here, partial dissolution of the solid produces enough fluoride ions to convert the crystal surfaces to fluorapatite.³²

      Fig. 2.7 Variation of solubility of hydroxyapatite and fluorapatite with pH.

      Fig. 2.8 The crystal lattice of hydroxyapatite (see Fig. 2.6) contains channels formed by stacked triangles of Ca²⁺ ions (orange), with successive triangles being rotated by 60°. In hydroxyapatite (A), the OH⁻ ions (large white circles representing the oxygen + small black circle representing the hydrogen) are too large to fit within the Ca triangles. In fluorapatite (B), the F⁻ ions (green) fit within the triangles, because they are smaller than the OH⁻ ion, and a more compact structure results. In fluorhydroxyapatites (C), F⁻ ions occupy some of the OH⁻ sites and can form hydrogen bonds with adjacent OH⁻ ions (dashed lines); this helps to stabilize the structure.

      Fig. 2.9 Effects of fluoride (F) on hydroxyapatite (HA) dissolution. HA and hydroxyapatites modified with fluoride (FHA, SFA) were added to acetate buffer, pH 5.0, containing 0, 0.1, or 5mg/L of F. The dissolution process was followed by measuring the release of Ca with time. FHA1 and FHA2 contained, respectively, 44 and 602mg/g F in the crystal lattice, whereas SFA1 and SFA2 were samples of HA powder that had been treated with F solution and contained respectively 60 and 706 mg/g F, which was concentrated at the crystal surfaces. The results show, first, that fluoride incorporated in the crystal structure or adsorbed to the crystal surfaces reduces dissolution rate, since the rates of FHA and SFA in fluoride-free buffer were lower than that of HA; and second, that the presence of fluoride in solution reduced dissolution rate, even that of HA, which contained no fluoride at all. (Data from ref.²⁹.)

      The lower solubility of fluorapatite also affects crystal growth. If fluoride ions are available in a solution supersaturated with respect to hydroxyapatite, crystal growth is accelerated. The solid that forms will not be hydroxyapatite but fluorapatite or a fluorhydroxyapatite, depending on how much fluoride is available, so the product of remineralization in the presence of fluoride tends to be less soluble than the mineral which had been lost.⁷

      If teeth are exposed to high concentrations of fluoride, for example, by treatment with fluoride varnish, a form of calcium fluoride (CaF₂), probably combined with phosphate, may be precipitated on the tooth surfaces.³¹ The Ca²⁺ ions required for precipitation of the CaF₂-like material are derived from the tooth mineral, so more of the precipitate is formed at lower pH. Although its formation removes Ca²⁺ ions from the tooth surface, CaF₂-like material could have a beneficial effect: since it is relatively soluble, it can act as a fluoride reservoir, maintaining raised concentrations of Ca²⁺ and F⁻ ions in the tooth environment.

      On the basis of this evidence, the prevailing current opinion is that strategies for caries prevention using fluoride should be aimed at maintaining low but sufficient concentrations of fluoride ions in the environment of the tooth, rather than at increasing the fluoride concentration in the tooth mineral.³¹ This is achieved by topical methods of fluoride administration such as toothpastes, mouth rinses, varnishes, and