The figure shows two electrolytic cells connected in series
(a) How much electricity is required for the reduction of 1 mole of `Ag^(+)` ions to Ag ?
(b) If three Faradays of electricity are passed through these cells, what is the ratio of the cations (`Ag^(+)` and `Cu^(2+)`) deposited at the cathodes ?

Explanation

(a) The electrode reaction involving `Ag^(+)` ions is :
`Ag^(+)+e^(-) to Ag`
Quantity of electricity required for the reduction of 1 mol of `Ag^(+)` ions =1F-96500 C.
(b) The electrode reactions in the two cells are :
`{:(Ag^(+)+e^(-), to Ag, ,Cu^(2+)+2e^(-), to Cu),(" "1F," "1mol," "2F," "1mol),(" "3F," "3mol," "3F," "3//2mol):}`
Molar ratios of `Ag^(+)` and `Cu^(2+)` ions deposited in the two cells is =3 mol : 3//2 mol or 2:1.