H2O2 (aq) = O2 (g) in acidic solution

(1) Assign oxidation states. In this case, the H is + 1, the O in peroxide is - 1 and in O2 is 0. Since you've only got something being oxidized here, you're only going to get an oxidation half-reaction.

(2) Write the half-reaction skeleton, with just the species being oxidized or reduced, and the electrons in the right place: H2 O2 (aq) → O2 (g) + 2 e -

Notice I put two electrons being released, because two oxygen atoms are being oxidized from - 1 to 0.

(3) Balance elements other than O and H. Already done.

(4) Balance oxygen by adding water to the appropriate side. Don't need to.

(5) Balance hydrogen by adding H+: H2 O2 (aq) →2 H + (aq) + O2 (g) + 2 e-

(6) If your reaction is in acid solution, you're done. If it's in basic solution, add OH - to both sides to cancel the H+: H2 O2 (aq) + 2 OH - (aq) →2 H2 O (l) + O2 (g) + 2 e-