Ammonium nitrate comproportionates into dinitrogen monoxide in the following reaction:
$\ce{NH4+ (aq) + NO3- (aq) -> N2O(g) + 2H2O (l)}$
I attempted to write the half-reactions by calculating the oxidation state for nitrogen in the ammonium and nitrate ions, and then finding the electrons needed for each half-reaction.
For example, for the ammonium half-reaction, I noticed that $\ce{N}$ is in a -3 oxidation state in $\ce{NH4+}$ and in a +1 oxidation state in $\ce{N2O}$. I concluded that four electrons must reduce ammonia and wrote the following preliminary equation:
$\ce{NH4+ (aq) + 4e- -> N2O (g)}$
Which I balanced the elements using standard acidic redox into:
$\ce{2NH4+ (aq) + H2O (l) + 4e- -> N2O(g) + 6H+ (aq)}$
However, I realized that my chemical equation is not balanced with respect to charge, indicating I have a flawed half-reaction.
What would be the correct half-reactions for this chemical reaction?
