Chapter 16 Precipitation equilibrium Solubility

l All dissolving is an equilibrium. l If there is not much solid it will all dissolve. l As more solid is added the solution will become saturated. l Soliddissolved l The solid will precipitate as fast as it dissolves. l Equilibrium

General equation l M + stands for the cation (usually metal). l Nm - stands for the anion (a nonmetal). l M a Nm b (s) aM + (aq) + bNm - (aq) l K = [M + ] a [Nm - ] b /[M a Nm b ] l But the concentration of a solid doesn’t change. l K sp = [M + ] a [Nm - ] b l Called the solubility product for each compound.

Watch out l Solubility is not the same as solubility product. l Solubility product is an equilibrium constant. l it doesn’t change except with temperature. l Solubility is an equilibrium position for how much can dissolve. l A common ion can change this.

l The solubility of iron(II) oxalate FeC 2 O 4 is 65.9 mg/L (MM= )

Calculating K sp l The solubility of iron(II) oxalate FeC 2 O 4 is 65.9 mg/L (MM= )

l The solubility of Li 2 CO 3 is 5.48 g/L (MM=73.88)

Calculating Solubility l The solubility is determined by equilibrium. l Its an equilibrium problem. l Calculate the solubility of SrSO 4, with a Ksp of 3.2 x in M.

l Calculate the solubility of Ag 2 CrO 4, with a Ksp of 9.0 x in M.

l Fe(OH) 3 (4 ions), what’s the molarity of Fe +3 in a saturated solution. l Ksp= 1.8E-15

l Calculate molarity of the salt and ions in Sr 3 (PO 4 ) 2 Ksp = 1.0E-31

l Calculate the solubility of Co(OH) 2 (Ksp = 2.5E-16) at pH 11.50

Relative solubilities l Ksp will only allow us to compare the solubility of solids that fall apart into the same number of ions. l The bigger the Ksp of those, the more soluble. l If they fall apart into different number of pieces you have to do the math.

Common Ion Effect l If we try to dissolve the solid in a solution with either the cation or anion already present less will dissolve.

l Calculate the solubility of SrSO 4, with a Ksp of 3.2 x in M in a solution of M Na 2 SO 4.

l Calculate the solubility of SrSO 4, with a Ksp of 3.2 x in M in a solution of M SrNO 3.

pH and solubility l OH - can be a common ion. l More soluble in acid. l For other anions if they come from a weak acid they are more soluble in a acidic solution than in water. l CaC 2 O 4 Ca +2 + C 2 O 4 -2 l H + + C 2 O 4 -2 HC 2 O 4 - l Reduces C 2 O 4 -2 in acidic solution.

Precipitation l Ion Product, Q =[M + ] a [Nm - ] b l If Ksp<Q a precipitate forms. l If Ksp>Q No precipitate. l If Ksp = Q equilibrium.

l A solution of mL of 4.00 x M Ce(NO 3 ) 3 is added to mL of 2.00 x M KIO 3. Will Ce(IO 3 ) 3 (Ksp= 1.9 x M)precipitate and if so, what is the concentration of the ions?

Selective Precipitations l Used to separate mixtures of metal ions in solutions. l Add anions that will only precipitate certain metals at a time. l Used to purify mixtures. l Often use H 2 S because in acidic solution Hg +2, Cd +2, Bi +3, Cu +2, Sn +4 will precipitate.

Selective precipitation l Follow the steps first with insoluble chlorides (Ag, Pb, Ba) l Then sulfides in Acid. l Then sulfides in base. l Then insoluble carbonate (Ca, Ba, Mg) l Alkali metals and NH 4 + remain in solution.

Complex ion Equilibria l A charged ion surrounded by ligands. l Ligands are Lewis bases using their lone pair to stabilize the charged metal ions. l Common ligands are NH 3, H 2 O, Cl -,CN - l Coordination number is the number of attached ligands. l Cu(NH 3 ) 4 +2 has a coordination # of 4

The addition of each ligand has its own equilibrium l Usually the ligand is in large excess. l And the individual K’s will be large so we can treat them as if they go to equilibrium. l The complex ion will be the biggest ion in solution.

l Calculate the concentrations of Ag +, Ag(S 2 O 3 ) -, and Ag(S 2 O 3 ) -3 in a solution made by mixing mL of AgNO3 with mL of 5.00 M Na 2 S 2 O 3 l Ag + + S 2 O 3 -2 Ag(S 2 O 3 ) - K 1 =7.4 x 10 8 l Ag(S 2 O 3 ) - + S 2 O 3 -2 Ag(S 2 O 3 ) -3 K 2 =3.9 x 10 4

l If you’re not part of the solution l Then you’re part of the precipitate!