By Ted Culbertson and Scott Becker

 Always Soluble  NO 3 -, C 2 H 3 O 2 -, Alkali Metal Cations, NH 4 +  Usually Soluble  Halogen Anions, SO 4 2-  Usually Insoluble  S 2-, CO 3 2-, PO 4 3-, OH -

 Leave out Spectator Ions  Check each one to see if it stays dissolved throughout  Cancel them if they do  EX:  2NaOH + Ag(NO 3 ) 2 → Ag(OH) 2 + 2Na + + 2NO 3 -  should be written as OH - + Ag +2 → Ag(OH) 2

 Is Dynamic  For the reaction aA + bB → cC + dD  K = [C] c [D] d /[A] a [B] b  K is determined by the Gibbs Free Energy of the reaction  Only include Gasses and Aqueous things  At Equilibrium = K eq  For Salts = K sp  For Complex Ions = K f  When not at Equilibrium, replace K with Q

 If Q > K, reaction reverses  If Q < K, reaction proceeds  If Q = K, reaction has reached Equilibrium  EX:  K sp for CaCO 3 is 3.8e-9.  If [Ca +2 ][Co 3 -2 ] > 3.8e-9, a precipitate will form

 If Equilibrium is disturbed, reaction will shift to reverse the change  It helps to use a Disturbance Chart: What Happened A ↔B +CShift More C was added ↑↓ ↓↓Left

 If Q > K sp, the salt precipitates  Two salts with a shared ion are less likely to dissolve  To separate:  Lower Solubility precipitates first  Get Q to just equal the K sp of that salt, without reaching the K sp of the other salt

 K f are usually much greater than other K types  EX:  Initially have AgCl at equilibrium in a one liter solution.  K sp = 1.6e-10. [Ag + ] = [Cl - ] = 1.3e-5 M.  Adds.10 moles of NaCl.  [Cl - ] is approximately.10 M.  1.6e-10 = [Ag + ](.10). [Ag + ] = 1.6e-9