[GoldShadow]

This question appeared as a bonus on an exam:

The Ka of H₃PO₄ is 7.6 x 10^-3.  Calculate the equilibrium constant for the formation of H₃PO₄ From H₂PO₄^- and explain in words why we consider H₂PO₄^- to be an acid despite its amphiprotic nature.

Using Kw=Ka*Kb I calculated that for H₂PO₄^- + H₂O<--> H₃PO₄ + OH^-, Kb= 1.316 x 10^-12.  But other than that I don't know what to do using the information given in this problem.  I looked up the Ka2 for H₃PO₄ and found that it was 6.2 x 10^-8, which is bigger than the Kb of H₂PO₄^-, but is there any way to determine using only the information given in the problem?

[english]

If you know the relative Ka and Kb of H₂PO₄^-, then you can determine the average behavior of the compound in solution.  More prone to acidic character or basic?

What do high Ka and Kb values indicate?  What do the size of these constants say about the progress of their respective reactions? 

Strong acidic behavior is determined by favorable deprotonation of the acid, and strong basic character is determined by favorable protonation of the base.  This is the Bronsted-Lowry concept.

[Borek]

What is pH of NaH₂PO₄ solution?