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:
H2O2(aq)→O2(g)+2e−
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+:
H2O2(aq)→2H+(aq)+O2(g)+2e−
(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+:
H2O2(aq)+2OH−(aq)→2H2O(l)+O2(g)+2e−
