The equivalent weight of ${\mathrm{Cl}}_2$ is $47.33$ grams in the following disproportionation reaction

${\mathrm{Cl}}_2 +{\mathrm{OH}}^{-}\to {\mathrm{Cl}}^{-}+{\mathrm{H}}_2\mathrm{O}+\text{oxidised product} \left(\text{molar mass of} {\mathrm{Cl}}_2=71 \mathrm{g}\right)$

The oxidised species is

$\mathrm{O}=\underset{1}{\mathrm{C}}=\underset{2}{\mathrm{C}}=\underset{3}{\mathrm{C}}=\mathrm{O}$

What are the oxidation states of ${\mathrm{C}}_1$, ${\mathrm{C}}_2$ and ${\mathrm{C}}_3$, respectively?

$3\mathrm{Cu}+8{\mathrm{HNO}}_3\to 3\mathrm{Cu}({\mathrm{NO}}_3{)}_2+4{\mathrm{H}}_2\mathrm{O}+2\mathrm{NO}$

How much mass of ${\mathrm{HNO}}_3$ is required to make $1 \mathrm{L}$ of $2 \mathrm{N}$  solution to be used as an oxidising agent in a reaction?