1. The problem statement, all variables and given/known data 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? 2. Relevant equations I suppose vital givens will be: pH = p_Ka + log[(A-)/(HA)] p_Ka values for H3PO4 = 2.15, 6.78, 12.4 3. 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. Thank you.