If a springly soluble salt is placed in water, after some time an equilibrium is established when the rate of dissolution of ions form the soid equal to the rate of precipitation of ions from the saturated solution at a particular temperature. Thus, a dynamic equilibrium exists between the undissociated solid species and the dissolved ionic species in a saturated and the dissolved ionic species in a saturated solution at a particular temperature. For example, in AgCl, we have the following equilibrium: AgCl_((aq.)) Ag_((aq))^(+) + Cl_((aq))^(-) The equilibrium constant K_(eq) = ([Ag^(+)][Cl^(-)])/([AgCl]) K_(eq) xx [AgCl] = [Ag^(+)] [Cl^(-)] rArr K_(sp) (AgCl) = [Ag^(+)][Cl^(-)]"........"(A)" :' [AgCl] is constant If there would not have been a saturated solution, then from equation (A), Keq. [AgCl] ne K_(sp), but K_(eq).[AgCl] = Q_(AgCl), where Q is ionicproduct, it implies that for a saturated solution, Q = K_(sp) K_(sp) is temperature dependent. When Q lt K_(sp), then the solution is unsaturated and there will be no precipitate formation. When Q = K_(sp), then solution will be saturated, no and ppt. will be formed When Q gt K_(sp), the solution will be supersaturated and there will be formation precipitate. At 25^(0)C, will a precipitate of Mg(OH)_(1) form when a 0.0001 M solution of Mg(NO_(3))_(2) is adjusted to a pH of 9.0 ? At what minimum value of pH will precipitation start? [Given : K_(sp) (Mg(OH)_(2)) = 10^(-11) M^(3)]