1. Solubility & SOLUBILITY PRODUCT CONSTANTS

2. Solubility Rules All Group 1 (alkali metals) and NH 4 + compounds are water soluble. All nitrate, acetate, and chlorate compounds are water soluble. All Br -, Cl -, I - (except mercury (I), Pb +2, and Ag + ) compounds are water soluble. All sulfate (except Pb +2, Ba +2, Sr +2 ; Ag + and Ca +2 are sparingly soluble) compounds are water soluble.

3. Solubility Rules All oxides, hydroxides, phosphates, carbonates, and sulfides form water INSOLUBLE compounds except Group 1 (alkali metals) and ammonium compounds.

4. Solubility Products Copper (I) bromide has a solubility of 0.02869 g/L at 25°C. What is the equilibrium constant for this reaction? The equation for the solubility of this compound is written:  CuBr (s)  Cu + Br

5. Solubility Products The equilibrium constant is written: K sp =

6. Solubility Products Convert mass to moles: 0.02869 g ( ) = 2 x 10 -4 moles [Cu + ] = 2 x 10 -4 moles/L [Br - ] = 2 x 10 -4 moles/L

7. Solubility Products Calculate the K sp K sp = (2 x 10 -4 moles/L) (2 x 10 -4 moles/L) K sp = 4 x 10 -8

8. SOLUBILITY PRODUCT CONSTANTS  Consider the following reaction  The equilibrium constant expression is K sp = [Pb 2+ ][Cl - ] 2  K sp is called the solubility product constant or simply solubility product  For a compound of general formula, M y X z

9. K sp = [M z+ ] y [X y- ] z K sp = [Mg 2+ ][NH 4 + ][PO 4 3- ] K sp = [Zn 2+ ][OH - ] 2 K sp = [Ca 2+ ] 3 [PO 4 3- ] 2

10.  Molar solubility: the number of moles that dissolve to give 1 liter of saturated solution  As with any equilibrium constant the numerical value must be determined from experiment  The K sp expression is useful because it applies to all saturated solutions - the origins of the ions are not relevant  Consider that K sp BaSO 4 = 1.1 x 10 -10

11. FACTORS AFFECTING SOLUBILITY The common ion effect  This effect affects solubility equilibria as it does other ionic equilibria  The solubility of a compound is less in a solution that contains a common ion than in pure water  Consider again the reaction  Adding Pb 2+ or Cl - precipitates PbCl 2

12.  Can be explained using LeChatelier’s principle  Likewise adding PbCl 2 to aqueous NaCl results in [Pb 2+ ][Cl - ] 2 > K sp  So for instance MgF 2 is less soluble in 0.10 M NaF than in pure water by a factor of 35000 pH of the solution  Solubility of some solids in water may be increased by changing the pH

13. e.g. Al(OH) 3 is more soluble in acidic solutions than in pure water  Salts that contain basic anions such as CN -, F -, and PO 4 3- will react in this fashion  Salts that contains anions of strong acids do not show this reactivity

14. Complex ion formation  Solubility of an ionic compound increases in a solution that contains a coordinating Lewis base  Typical examples are NH 3, CN -, OH - and Cl -

15.  Similarly,  The reaction decrease the metal ion concentration so that [M 2+ ][OH - ] 2 < K sp and the M(OH) 2 dissolves

16. THE REACTION QUOTIENT  The reaction quotient (called ion product) may be applied to solubility equilibria - determines if a substance will precipitate from solution Q sp < K sp Forward process occurs No precipitation occurs Q sp = K sp Solution is just saturated Q sp > K sp Reverse process occurs Precipitation occurs

17. SELECTIVE PRECIPITATION  This is a method which allows separation of metal ions or anions based on solubility e.g. Cl -, Br -, and I - can be separated by selective precipitation with silver ions K sp AgCl = 1.8 x 10 -10 AgBr = 3.3 x 10 -13 AgI = 1.5 x 10 -16