The molecular weight of a monoprotic acid HX was to be determined. A sample of 15.126 grams of HX was dissolved in distilled water and the volume were brought to exactly 250mL in a volumetric flask. Several 50mL portions of this solution were titrated against NaOH solution, requiring an average of 38.21mL of NaOH.
The NaOH solution was standardized against oxalic acid dihydratem H2C2O4(2H2O) (molecular weight: 126.066 grams per mol). The volume of NaOH solution required to neutralize 1.2596 grams of oxalic acid dihydrate was 42.21mL
a) Calculate the molarity of the NaOH solution
pH = pKa + log (A/HA)
That's what I think is going on here, no?
The Attempt at a Solution
I'm stuck at the second paragraph of the problem, am I supposed to assume that oxalic acid dihydrate completely dissociate into the hydromium ions and the conjugate base? If it is then this turns into a simple problem.
But what if not all of the oxalic acid dihydrate dissociate completely? How am I supposed to calculate the concentration of the NaOH if I don't have the Ka of the oxalic acid dihydrate?