You have 400 mL of a solution of 0.20 M Na3PO4 buffer at pH 6.5. You add 0.7 g
of NaOH to this solution. What will be the new pH?
I suppose vital givens will be:
pH = p_Ka + log[(A-)/(HA)]
p_Ka values for H3PO4 = 2.15, 6.78, 12.4
The Attempt at a Solution
Concerns that I have are:
1) Did I use the correct p_Ka value? (This is if the Henderson-Hasselbalch equation is of any use.)
2) Did I find each species’ mols accordingly?
3) As a weak triprotic acid/system is featured, have I setup and represented the involved species for final pH?
0.20 M Na3PO4 * 0.4 L Na3PO4 = 0.08 mol Na3PO4
Base will react fully with acid.
The question that I have before typing out my workings is the following.
At pH 6.5, is this valid to propose: [H3PO4] + [H2PO4 (-)] + [HPO4 (2-)] + [PO4 (3-)] = 0.08 mol of phosphate? I thought that perhaps the first two species would be predominant, or not negligible, at this pH.
Any advice is appreciated.