The value of the standard electrode potential for the oxidation of $Cl^{-}$ ions is more positive than that of water. Even then,in the electrolysis of aqueous sodium chloride,why is $Cl^{-}$ oxidized at the anode instead of water?

(N/A) The oxidation reactions at the anode are:
$(i)$ $Cl_{(aq)}^{-} \rightarrow \frac{1}{2} Cl_{2(g)} + e^{-}$; $E^{\circ} = 1.36 \ V$
$(ii)$ $2H_{2}O_{(l)} \rightarrow O_{2(g)} + 4H_{(aq)}^{+} + 4e^{-}$; $E^{\circ} = 1.23 \ V$
Although the standard oxidation potential of water $(1.23 \ V)$ is lower than that of $Cl^{-}$ $(1.36 \ V)$,which theoretically suggests water should be oxidized,the oxidation of $Cl^{-}$ is observed.
This occurs due to the 'overpotential' of oxygen. The evolution of oxygen gas at the anode requires an extra potential (overpotential) beyond the standard value,making the effective potential for water oxidation significantly higher than $1.36 \ V$. Consequently,$Cl^{-}$ ions are preferentially oxidized to $Cl_{2}$ gas.