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?
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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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1$\begingroup$ Rule with oxo groups works for every central atom separately, and accounts rather for mesomeric than inductive effects. $\endgroup$ – Mithoron Jul 10 '15 at 22:43
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$$\ce{H2SO4 ⇌ H2O + SO3}$$
$$\ce{SO3 + H2SO4 ⇌ H2S2O7}$$
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
includes not just questions from actual homework assignments, but also self-study problems, puzzles, etc.$\endgroup$ – user15489 Jun 13 '15 at 5:04