Solubility Equilibria

1. Solubility Equilibria • Solubility is a very important phenomenon. • We can flavor foods due to the solubility of salt and sugar in water. • In this section, we will consider sparingly soluble substances and quantitative ways to determine “How soluble?”. • According to the solubility rules previously learned, PbCl2 and AgCl are both insoluble salts. • However, if chloride ion is added to a solution containing both Pb2+ and Ag+ ions, nearly all of the Ag+ is precipitated as AgCl before any Pb2+ separates from the solution as PbCl2. • This occurs because AgCl is much less soluble than PbCl2. • To explain these differences in solubility, solubility equilibrium must be examined quantitatively.

2. When placed in water, a small amount of AgCl dissolves and the following equilibrium is established once the solution becomes saturated: • AgCl(s) ⇌Ag+(aq) + Cl-(aq) • Equilibrium expression: • Ksp = [Ag+][Cl-] • Ksp = solubility product constant • Kspequals the product of concentration terms for the ions dissolved in a saturated solution of a sparingly soluble substance. • The solubilities of salts change with temperature so a value of Ksp applies only to solutions only at the temperature at which its value was determined. If reactants or products have a coefficient other than 1 concentrations must be raised to that power.

3. Ksp can be obtained from a salt’s molar solubility in water – the number of moles of solute dissolved in one liter of its saturated solution. • Examples • Molar solubility can also be computed (estimated) from values of Ksp. • Examples

4. Common Ion Effect • Suppose we stir some calcium carbonate in water long enough to establish the following equilibrium: • CaCO3 (s) ⇌ Ca2+ (aq) + CO32- (aq) • Then we add to the solution a very solublesalt of calcium, like CaCl2. • This puts Ca2+ into solution, and it upsets the above equilibrium. • The ion product is no longer equal to Ksp. • Remember from Le Chatelier’s principle, if we add then equilibrium shifts opposite. • The above equilibrium shifts to the left causing CO32- to precipitate as CaCO3.

5. Eventually equilibrium is reestablished, but with a lower concentration of CO32+ in solution. • In this new system, there are two sources of Ca2+, the added CaCl2 and the CaCO3 still in solution. • Because Ca2+ is common to both sources, it is called a common ion. • The addition of the common ion lowers the solubility of CaCO3; it is less soluble in the presence of CaCl2 (or any other soluble calcium salt) than it is in pure water. • This lowering of the solubility of an ionic compound by the addition of a common ion is called the common ion effect. • The common ion effect can dramatically lower the solubility of a salt.

6. Will a Precipitate Form? • So far we have considered solids dissolving in solutions. • Now we will consider the reverse process – the formation of a solid from a solution. • We will use the ion product. • Ion product (Q) is defined like expression for Ksp but uses initialconcentrations instead of equilibrium concentrations. • Precipitate will form Ion product >Ksp (supersaturated) • No precipitate will form Ion product =Ksp (saturated) • No precipitate will form Ion product < Ksp (unsaturated)

7. pH and Solubility • The pH of a solution can affect a salt’s solubility. • Example: • Mg(OH)2(s) ⇌ Mg2+(aq) + 2OH-(aq) Add OH- (increase basicity) Add H+ (increase acidity -removes OH- by reaction to produce H2O) Decrease Solubility of Mg(OH)2 Increases solubility (lose OH- and must replace) • General rule – If the anion X- is an effective base (HX is a weak acid) the salt MX will show increased solubility in an acidic solution.

8. Complex Ion Equilibria • A complex ionis a charged species consisting of a metal ion surrounded by ligands. • A ligand is simply a Lewisbase. • Recall a Lewis base is an ion or molecule having a lone electron pair that can be donated to an empty orbital on the metal ion to form a covalent bond. • Some common ligands are H2O, NH3, Cl-, and CN-.

9. Metal ions add ligands one at a time in steps characterized by equilibrium constants called formation constants or stability constants. • For example, when solutions containing Ag+ and NH3 molecules are mixed, the following reactions take place: • Ag+ (aq) + NH3 (aq) ⇌ Ag(NH3)+ (aq) K1 = 2.1 x 103 • Ag(NH3)+(aq) + NH3(aq)⇌ Ag(NH3)2+(aq) K2 = 8.2 x 103 • where K1 and K2 are the formation constants for the two steps. • In a solution containing Ag+ and NH3, all the species NH3, Ag+, Ag(NH3)+, and Ag(NH3)+ exist at equilibrium.

10. When we write the formula for a complex, we follow two rules: • The symbol for the metal ion is always given first, followed by the ligands. • The charge on the complex is the algebraic sum of the charge on the metal ion and the charges on the ligands. • For example, the formula of the complex ion of Cu2+ and H2O is written Cu(H2O)42+ with the Cu first followed by the ligands. • The charge on the complex is 2+ because the copper ion has a charge of 2+ and the water molecules are neutral.

11. Metal ions that commonly form complex ions (or coordination compounds): Al3+, Cu2+, Zn2+, Fe2+ (or 3+), Ni2+, Ag+. (All Curiously Colored Zebras Felt Nicely Agreeable) • Common ligands: NH3, OH-, Cl-, SCN-, CN-, H2O. • Most common coordination number: twice the charge of the metal ion. • That means that Ag+ and NH3→ Ag(NH3)2+. • Formation of complex ions is a common means to dissolve otherwise insoluble salts.

12. For example, in a solution with only water present, AgCl only dissociates (dissolves) slightly to form Ag+. • When ammonia is added, the Ag+ complexes with the ammonia, and the removal of the Ag+ from the solution as it converts to Ag(NH3)2+ pulls the AgCl dissociation equilibrium to the right (LeChatelier’s Principle). • If sufficient ammonia is added to complex all of the silver ions, the AgCl will completely dissolve. • AgCl(s) ⇌ Ag+ (aq) + Cl- (aq)