Balancing the equation when MnO₂ is reduced to MnO by H₂(g) I am studying for a final tomorrow and I was going through some problems on older exams and came across one which I wasn't really sure how to solve.

The problem is to write the balance equation when:

managnese(IV)oxide is reduced to manganese(II) oxide by hydrogen gas.

I took a few approaches to this, because at first I did not realize it say reduced "by" hydrogen gas. This part threw me off because I initially added $\ce{H2(g)}$ to the product side.

So far I have:
$\ce{MnO2(s) + H2(g)-&gt; MnO(s) + H2O(l)}$ which is not balanced yet. I noticed in this case the conventional type of balancing the equation does not work. Would this be a redox rxn in which I find each half reaction and then balance?

I am completely stuck at this point, and I don't think my formula is correct. Does anyone have any insight or good examples ?
This is balanced: I count the same number of elements on each side, which is the definition of balanced. What is not explicitly stated is that your hydrogen goes from $0$ to $+1$, and is oxidized, thus reducing your manganese from $+4$ to $+2$. But this all balances!