[shafaifer]

Let us say you have this salt: KH₂PO₄. To this salt you add HCl(aq):

KH₂PO₄(s) + HCl(aq) --> KCl(aq) + H₂PO₄^-(aq) + H⁺(aq)     (equation 1)
 
Now phosphate reacts with added molybdate in an acidic enviroment:
 
H₂PO₄^-(aq) + H⁺(aq) + (NH₄)₂MoO₄(aq)  --> (NH₄)₃[PMo₁₂O₄₀](s) (equation 2)
 
 
According to a lab manual, the product on the right side of the reaction in Eq. 2 is correct. The manual only states that phosphate reacts with molybdate and creates the product in equation 2 - not what the reactions in question look like. Have I made something correct? Is the balancing of this equation "only" a matter of stoichiometry skills?
 
Best regards,
 
Shafaifer

[Borek]

Be consistent about ions - whatever is dissolved is dissociated, whatever is solid, is not. Mixing KCl(aq) and H₂PO₄^-(aq) and H⁺(aq) in one equation doesn't make sense.

Final reaction equation should be just that of free ions combining into a single precipitate.