Question

In my research I try to prepare trimagnesium dicitrate aqueous solution. Initially it stays clear, but after some time white insoluble precipitate is formed. Initial pH was about 4,5-5,0. If pH is lowered, there is less precipitate. But why? From solubility equation pH should be >8 to form Mg(OH)2 precipitate at the concentration I am using. Is there some other Mg compound formed?

Answer 1

The formation of the white insoluble precipitate in a trimagnesium dicitrate solution at pH around 4.5-5.0 may be due to an equilibrium reaction with H+ ions at lower pH, leading to more Mg(OH)2 dissolving and thus less precipitate.

Step-by-step explanation:

When preparing a solution of trimagnesium dicitrate, you've observed the formation of a white insoluble precipitate over time, despite the initial clear solution. The pH when properly balanced around 4.5 to 5.0 should not cause precipitation of Mg(OH)2; only at pH levels above 8 should this occur given the concentration in your experiment.

However, if there is less precipitation when the pH is lowered, then it's likely that the white insoluble precipitate forming in your solution is due to a secondary equilibrium where Mg(OH)2(s) reacts with the H+ ions present in the acidic solution, forming more soluble Mg2+ ions and water, as depicted in the equation Mg(OH)2(s) + 2H+ (aq) = Mg2+ (aq) + 2H2O(l).

Consequently, by lowering the pH, you're providing more H+ ions which help to dissolve the magnesium hydroxide precipitate by pushing the equilibrium to the right and increasing the solubility of Mg(OH)2.