Three electrolytic cells A,B and C containing solutions of zinc sulphate, silver nitrate and copper sulphate, respectively are connected in series.
A steady current of 1.5 A was passed through them until 1.45g of silver deposited at the cathode of cell B. How long did the current flow? What mass of copper and zinc were deposited in the concerned cells? ( Atomic mass of Ag= 108, Zn = 65.4, Cu = 63.5)

Explanation

`Ag^(+) + underset( 1F) ( e^(-)) rarr underset("1 mol")( Ag)`
`:.` 108 g of Ag is deposited by 96500C
1.45g of Ag will be deposited by `= ( 96500 xx 1.45)/( 108) C`
= 1295.6C
Q = it or 1295.6 C `= 1.5 xx t`
`t = ( 1295.6C )/( 1.5)`
`rArr t = 863.7s`
Since, the relation related to the deposition of copper is
`Cu^(2+) + 2e^(-) rarr underset( 63.5g) ( Cu)`
`:. 2 xx 96500 C` electricity deposits 63.5 g of Cu
`:.` 1295.6 C electricity will deposits
` ( 63.5 xx 1295.6)/( 2 xx96500 ) g ` of Cu, i.e. of 0.426 g of Cu
Since, the reaction related to the deposition of zinc is
`Zn^(2+) + underset( 2 xx 96500C )( 2e^(-)) rarr underset( 65.4g) ( Zn)`
`:. 2 xx 96500C` electricity deposits 65.4g of Zn
`:.` 1295.6C electricity will deposit
`( 65.4 xx 1295.6)/( 2 xx 96500) g ` of Zn, i.e. 0.44 g of Zn.