The $pKa$ values for H3AsO4 and H3PO4 are $2.30$ and $2.12$ which indicate that H3PO4 is a stronger acid compared to H3AsO4. I looked up their electronegativity values to find that they are almost the same for both arsenic and phosphorus. I feel H3AsO4 should be the stronger acid because arsenic atom is larger in size and the charge is spread out in much more volume which should make the conjugate base more stable. What am I missing out?
H3PO4 has more EN atoms than H3AsO3 first of all. So it wins that first (checkmark more ability to stabalize a negative charge therefore more stable and that means stronger acid. P + 4 oxygens > As +3 oxygens. Secondly, there are more terminal oxygens on the H3PO4 as well. All of the oxygen atoms on the H3AsO3 are bonded to the 3 H atoms. IF they had an even number of electonegative atoms bonded to the central atom and equal (or even closely similar) number of terminal (not bonded to H) oxygens... then we can go ahead and compare their size since they are in the same column on the periodic table. When you have binary acids bonded to one other anion... go ahead and jump to electonegativity comparisons ( for comparisons in the same period) or size of anion comparisons (if in the same column or group.) first. With oxoacids like this first compare which one has more electronegative atoms .... then compare their electronegatively/ size.
The stability of an acid is judged by the strength of it's conjugate base. If the conjugate base is stable, it implies that it is a weak base and hence the corresponding acid is stronger.
The conjugate base would be one Hydrogen short from an OH group. Similarly with H3AsO4. However, the conjugation(resonance) will be better in the case of P=O bond, since it is an orbital overlap of Oxygen's 2p and Phosphorus' 3p as against Arsenic's 4p. 2p and 3p are of comparable size and account for a better overlap and better resonance, distribution of electronic charge and hence a more stable conjugate base.
Consequently, H3PO4 is better than H3AsO4 as an acid.