I've read that for an oxyacid, the oxidation number of the central atom shows the power of that acid, but here, the oxidation numbers of sulphur in sulphuric acid and pyrosulphuric acid are both the same, i.e. +6. That is confusing! How should I proceed now?

  • $\begingroup$ Welcome to Chemistry.SE. Take the tour to get familiar with this site. This appears to be a homework question, please share your thoughts and attempts towards the solution. It'll make us certain that ‎we aren't doing your homework for you. $\endgroup$
    – user15489
    Jun 13, 2015 at 3:43
  • $\begingroup$ Thank you. That's not my homework. I've read dat for an oxi-acid the oxidising number of the centre atom shows the power of dat acid. But here in this case,oxidising numbers of sulpher in sulphuric acid and pirosulphuric acid both are same,+6. That makes me confused! :/ $\endgroup$ Jun 13, 2015 at 3:55
  • $\begingroup$ From the homework policy link above includes not just questions from actual homework assignments, but also self-study problems, puzzles, etc. $\endgroup$
    – user15489
    Jun 13, 2015 at 5:04
  • $\begingroup$ H2S2O7 is stonger than H2SO4 due to conjugate base of olium is more stable than conjugate base of sulfuric acid.Negative charge is delocalize by two S=O bond for olium but in sulfuric acid it one . $\endgroup$
    – user16806
    Jun 13, 2015 at 17:28
  • $\begingroup$ @Totan Nandi It is more stable but mesomeric stabilization is practically the same. It's inductive effect that matters. $\endgroup$
    – Mithoron
    Jun 13, 2015 at 20:46

2 Answers 2


To answer this question at a more relatable level, typically the strength of oxoacids can be determined from the number of "lone" oxygen atoms in the compound--that is, the number of oxygen atoms that do not have a corresponding hydrogen. What usually happens is the central atom in an oxoacid will draw electrons from the bonded oxygen atoms, thereby making them more electronegative. This increases the polarity of the O-H bond, increasing acidity. In your example, H$_2$SO$_4$ has 2 lone oxygen atoms; while in H$_2$S$_2$O$_7$, 5 lone oxygens exist. We therefore suspect that the latter is more acidic. Although there are exceptions to this rule, for a general chemistry course it typically is valid.

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    $\begingroup$ Rule with oxo groups works for every central atom separately, and accounts rather for mesomeric than inductive effects. $\endgroup$
    – Mithoron
    Jul 10, 2015 at 22:43

$$\ce{H2SO4 ⇌ H2O + SO3}$$

$$\ce{SO3 + H2SO4 ⇌ H2S2O7}$$

The mutual electron-withdrawing effects of each sulfuric acid unit on its neighbour causes a marked increase in acidity. Disulfuric acid is strong enough to protonate "normal" sulfuric acid in the (anhydrous) sulfuric acid solvent system.

From the Wikipedia entry on Disulfuric acid


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