Once upon a time, a common method for determining phosphorous was to convert it to phosphate, precipitate it as magnesium ammonium phosphate, MgNH4PO4, and ignite it to form magnesium pyrophosphate, Mg2P2O7. The solubility reaction for MgNH4PO4 is
MgNH4PO4(s) ↔ Mg2+(aq) + NH4+(aq) + PO43-(aq)
for which the Ksp is 2.51 x 10-13. Additional equilibrium constants of interest are listed below.
H3PO4: pKa1 = 2.33, pKa2 = 7.22, pKa3 = 12.33
NH4+: pKa = 9.25
(a) Le Châtelier’s principle suggests that the solubility of MgNH4PO4 will decrease if we add an excess of NH4+. Experimental work, however, shows that the solubility actually increases. Explain why the solubility increases when NH4+ is added in excess. You may approach this using a ladder diagram, or by combining appropriate equilibrium constant expressions.
(b) The procedure calls for dissolving the sample in HCl and HNO3. Solutions of NH4+ and Mg2+ are added along with additional HCl to ensure an acidic solution. The precipitate is then formed by adding NH3 dropwise until the pH is neutral, followed by adding 2 mL in excess. Clearly explain why the solution is initially made acidic, why the precipitant is added dropwise, and why the precipitant is added in excess.