The $\mathrm{p}K_\mathrm{a}$ values for $\ce{H3AsO4}$ and $\ce{H3PO4}$ are $2.30$ and $2.12$ which indicate that $\ce{H3PO4}$ is a stronger acid compared to $\ce{H3AsO4}$. I looked up their electronegativity values to find that they are almost the same for both arsenic and phosphorus. I feel $\ce{H3AsO4}$ 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?

  • 2
    $\begingroup$ Charge is mostly on oxygens, BTW such a difference as this is negligible . $\endgroup$
    – Mithoron
    Commented Aug 3, 2018 at 22:09
  • 4
    $\begingroup$ When comparing acidic strength for oxoacids with the same number of oxygens, typically the acid with the central atom of higher electronegativity - in this case P, will be more acidic. This is because the electrons are slightly withdrawn from the OH bond as the electronegativity of the central atom increases. As bonding electrons are pulled closer to the OH bond, the molecule becomes more polar, and so the molecule becomes a stronger acid. $\endgroup$ Commented Aug 3, 2018 at 23:00
  • 2
    $\begingroup$ Wikipedia lists the electronegativity of phosphorous as 2.253 and arsenic as 2.211 on the Allen scale. Wikipedia's values for the pKa1, pKa2 and Ka3 of phosphoric acid are 2.148, 7.198 and 12.319. For arsenic acid the values are 2.19, 6.94, and 11.5. Consider the pka2 and pKa3 values I find the electronegativity argument dubious. $\endgroup$
    – MaxW
    Commented Aug 4, 2018 at 0:35

2 Answers 2


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.

Now consider $\ce{H3PO4}$. The structure has one $\ce{P=O}$ bond and $\ce{3 P-OH}$ bonds.

Phosphoric acid

The conjugate base would be one Hydrogen short from an $\ce{OH}$ group. Similarly with $\ce{H3AsO4}$. However, the conjugation (resonance) will be better in the case of $\ce{P=O}$ bond, since it is an orbital overlap of oxygen's $\mathrm{2p}$ and phosphorus' $\mathrm{3p}$ as against arsenic's $\mathrm{4p}$. $\mathrm{2p}$ and $\mathrm{3p}$ 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, $\ce{H3PO4}$ is better than $\ce{H3AsO4}$ as an acid.


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.

  • 1
    $\begingroup$ The OP has in mind $\ce{H3AsO4}$, not $\ce{H3AsO3}$. $\endgroup$
    – Poutnik
    Commented Apr 6, 2019 at 11:55

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service and acknowledge you have read our privacy policy.

Not the answer you're looking for? Browse other questions tagged or ask your own question.