Q.What happens when Cl2 reacts with dry slaked lime?

Chlorine, $Cl_2$, is prepared in the laboratory or industrially by oxidising chloride ion, for example by heating sodium chloride with concentrated sulphuric acid in the presence of manganese dioxide, $2NaCl + MnO_2 + 3H_2SO_4 \rightarrow 2NaHSO_4 + MnSO_4 + 2H_2O + Cl_2$; it also readily oxidises iodide ion, so passing chlorine gas into aqueous sodium iodide liberates iodine, $Cl_2 + 2NaI \rightarrow 2NaCl + I_2$. Chlorine reacts with hot, concentrated sodium hydroxide in a disproportionation reaction, being simultaneously oxidised (to $+5$) and reduced (to $-1$) from its starting oxidation state of $0$: $3Cl_2 + 6NaOH(hot,\ conc.) \rightarrow 5NaCl + NaClO_3 + 3H_2O$. Its bleaching action arises because chlorine dissolved in water generates nascent (atomic) oxygen via hypochlorous acid, $Cl_2 + H_2O \rightarrow HCl + HOCl$, followed by $HOCl \rightarrow HCl + [O]$, and this nascent oxygen oxidatively destroys coloured organic matter. Hydrogen chloride, $HCl$, reacts with finely powdered iron to give ferrous chloride rat …