Answer:
A disproportionation redox reaction
Explanation:
1. Assign an oxidation number to every atom in the equation
[tex]2\stackrel{\hbox{$\mathbf{+4}$}}{\hbox{N}}\stackrel{\hbox{-2}}{\hbox{O }}_{2} + 2\stackrel{\hbox{-2}}{\hbox{O}}\stackrel{\hbox{+1}}{\hbox{H}^{-}} \longrightarrow \, (\stackrel{\hbox{$\mathbf{+3}$}}{\hbox{N}}\stackrel{\hbox{-2}}{\hbox{O }}_{2})^{-} + (\stackrel{\hbox{$\mathbf{+5}$}}{\hbox{N}}\stackrel{\hbox{-2}}{\hbox{O}_{3}})^{-} +\stackrel{\hbox{+1}}{\hbox{H}}_{2}\stackrel{\hbox{-2}}{\hbox{O}}[/tex]
2. Identify the atoms that change their oxidation number
N in NO₂: +4 ⟶ +3 in NO₂⁻
N in NO₂: +4 ⟶ +5 in NO₃⁻
3. Identify the type of change
This is a disproportionation — a reaction in which one substance is simultaneously oxidized and reduced.
