Calcium Phosphate Dissolution
The solubility of calcium salts is highly dependent on pH. For example, let’s look at tricalcium phosphate, which presents the solubility equilibrium:
$$\ce{Ca3(PO4)2(s) + H2O &lt;=&gt; 3Ca^2+(aq) + 2PO4^3-(aq)}$$
The solubility product constant for this equilibrium ($K_\mathrm{sp}$) is vastly varied from source to source. Some listed it as $\pu{2.0 × 10^–29 mol^5L{–5}}$, and another listed it as $\pu{2.07 × 10^−33 mol^5L^{–5}}$ at $\pu{25 ^{\circ}C}$ and $\mathrm{pH}~7.00$. However, I found a reliable $K_\mathrm{sp}$ value in literature, which reported that $K_\mathrm{sp}$ of $\beta$-$\ce{Ca3(PO4)2}$ in the $\mathrm{pH}$ range $6.0-7.5$ is $\pu{1.2 × 10^–29 mol^5L^{–5}}$ at $\pu{25 ^{\circ}C}$ and $\pu{2.83 × 10^–30 mol^5L^{–5}}$ at $\pu{37 ^{\circ}C}$ (a negative thermal coefficient of solubility; Ref.1). Therefore, the quantitative analysis of $\ce{Ca3(PO4)2}$ is near impossible.

Still, we can qualitatively look at the problem: If you decrease the $\mathrm{pH}$ of the solution in above equilibrium by adding acid to the solution, some of the phosphate ions would protonate and transform into $\ce{HPO4^2–}$ ions, as what happens in phosphate buffer solutions:
$$\ce{PO4^3-(aq) + H3O+ (aq) &lt;=&gt; HPO4^2-(aq) + H2O}$$
As a result, the concentration of $\ce{PO4^3-}$ ion would be reduced. The system would respond to this reduction by producing more $\ce{PO4^3-}$ ions, according to the Le Chatelier principle. To do so, some solid $\ce{Ca3(PO4)2}$ would dissolve in first equilibrium, and the equilibrium will be shifted to the right.

Keep in mind that in first equilibrium, $\ce{Ca/P}$ ratio is 1.5. The solubility of calcium phosphates is also depend on $\ce{Ca/P}$ (Ref.2). For example, both monocalcium phosphate ($\ce{Ca(H2PO4)2}$) and monocalcium phosphate monohydrate ($\ce{Ca(H2PO4)2.H2O}$) are highly soluble in water at $\pu{25 ^{\circ}C}$ ($\ce{Ca/P}$ is 0.5). On the other hand, $K_\mathrm{sp}$s of dicalcium phosphate ($\ce{CaHPO4}$) and dicalcium phosphate dihydrate ($\ce{CaHPO4.2H2O}$) are $1.26 × 10^{–7}$ and $\pu{2.57 × 10^{–7} mol^2L^{–2}}$, respectively at $\pu{25 ^{\circ}C}$ (Ref.2). Based on these values, when $\mathrm{pH}$ is gradually reduced to a point such that all calcium phosphates would dissolve.

References:


T. M. Gregory, E. C. Moreno, J. M. Patel, W. E. Brown, “Solubility of $\beta$-$\ce{Ca3(PO4)2}$ in the system $\ce{Ca(OH)2-H3PO4-H2O}$ at $5, 15, 25,$ and $\pu{37 ^{\circ}C}$,” Journal of Research of the National Bureau of Standards–A. Physics and Chemistry 1974, 78A(6), 667–674 (http://dx.doi.org/10.6028/jres.078A.042).
L.C. Chow, “Solubility of Calcium Phosphate,” In Octacalcium Phosplrate; L. C. Chow, E. D. Eanes, Vol. Eds., Volume 18 of Monographs in Oral Science (G. M. Whitford, Ed.) Karger Publishers: Basel, Switzerland, 2001, pp. 94-111 (DOI:10.1159/000061650).