What is the difference between balancing a redox reaction in acidic solution vs basic solution? I'm asked to balance the potassium permanganate and hydrogen peroxide reaction. I misread the problem and thought that the reaction was in acidic solution. This is the result that I've got:  

$$\ce{2MnO4 + 5H2O2 + 6H+ -&gt; 2Mn^2+ +5O2 +8H2O}$$
What do I have to change to balance the equation in a basic solution?.
You basically just have to add hydroxide ions to either sides of your equilibrium.

$$\ce{2MnO4 + 5H2O2 + 6H+ + \color{blue}{6OH-} -&gt; 2Mn^2+ +5O2 +8H2O + \color{blue}{6OH-}}$$

$\ce{H+}$ and $\ce{OH-}$ ions can be collectively re-written as $\ce{H2O}$, so we have:

$$\ce{2MnO4 + 5H2O2 + \color{blue}{6H2O} -&gt; 2Mn^2+ +5O2 +8H2O + 6OH-}$$

Finally, we can cancel out the excess water molecules which leaves us with:

$$\ce{2MnO4 + 5H2O2  -&gt; 2Mn^2+ +5O2 +2H2O + 6OH-}$$