Do all the precipitates undergo an equilibrium with its respective ions, why can't they dissociate fully?
In the two examples you mentioned, we have an equilibrium between the solid and the small quantity of the solid that is soluble in water and completely dissociated:
$$\ce{Ca3(PO4)2_{(s)}&lt;=&gt; 3Ca^2+_{(aq)}  + 2 PO4^{3-}_{(aq)} }$$ 
$$\ce{AgCl_{(s)}&lt;=&gt; Ag^+(aq)_{(aq)}  + Cl^{-}_{(aq)} }$$ 

I think it's wrong to write a solid forward reaction.
It's an heterogeneous equilibrium between the solid state (the precipitate) and the aqueous state (ions constituents of the precipitate). The constant of this equilibrium is the solubility product $K_{\mathrm{sp}}$ of the precipitate.