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(B) During the electrolysis of molten $NaCl$,the oxidation reaction at the anode is:$2Cl^{-} \longrightarrow Cl_2 + 2e^{-}$From the stoichiometry of t

Vedclass · Accepted Answer

$19300$

Vedclass · Answer

(B) During the electrolysis of molten $NaCl$,the oxidation reaction at the anode is:$2Cl^{-} \longrightarrow Cl_2 + 2e^{-}$From the stoichiometry of the reaction,$1 \ mole$ of $Cl_2$ gas is produced by the transfer of $2 \ moles$ of electrons.Therefore,to liberate $0.1 \ mole$ of $Cl_2$,the required number of moles of electrons is $0.1 	imes 2 = 0.2 \ mole$.The quantity of electricity $Q$ is given by $Q = n 	imes F$,where $n$ is the number of moles of electrons and $F$ is Faraday's constant $(96500 \ C \ mol^{-1})$.$Q = 0.2 \ mole 	imes 96500 \ C \ mol^{-1} = 19300 \ C$.

Calculate the quantity of electricity required to liberate $0.1 \ mole$ of chlorine gas during electrolysis of molten sodium chloride. (in $C$)

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