Do all the precipitates undergo an equilibrium with its respective ions, why can't they dissociate fully?

I’m really confused about this. If we consider $\ce{Ca3(PO4)2}$, it can be seen it undergoes an equilibrium. But why it’s an equilibrium?


This is the equilibrium it undergoes $$\ce{Ca3(PO4)2\_{(s)}<=> 3Ca^2+\_{(aq)} + 2 PO4\_{(aq)}^3-}$$


By the way I’m having a doubt of why some are using backward arrow for a dissociation of a particular precipitate while some are using a solid forward arrow for the same reaction. For an example let’s take $\ce{AgCl(s)}$


Now this is a place where a backward arrow is used,
$$\ce{AgCl\_{(s)}<=> Ag+\_{(aq)} + Cl\_{(aq)}-}$$ 


But I have seen in some books where there is a solid forward arrow instead of back-ward arrow for the dissociation of $\ce{AgCl(s)}$.




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:
$$\ce{Ca3(PO4)2\_{(s)}<=> 3Ca^2+\_{(aq)} + 2 PO4^{3-}\_{(aq)} }$$ 
$$\ce{AgCl\_{(s)}<=> Ag^+(aq)\_{(aq)} + Cl^{-}\_{(aq)} }$$ 


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.


 1