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

${\mathrm{Br}}_2$ undergoes disproportionation to ${\mathrm{Br}}^{-} \text{and} {\mathrm{BrO}}_3^{-}$, in the alkaline medium. Which of the following can be correct for the above reaction?

a. The total electron change is $10$.

b. There are $10{\mathrm{OH}}^{-}$ in the balanced equation on the left-hand side.

c. The equivalent weight of ${\mathrm{Br}}_2$ as reductant is $\mathrm{M}/10$ $\left(\mathrm{M}=\right.$ Molar mass of $\left.{\mathrm{Br}}_2\right)$

d. $\mathrm{n}-$factor of ${\mathrm{Br}}_2$ is $5/3$.

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

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?