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 1.45 g of silver deposited at the cathode of cell R 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)
${{Ag}^{+}}$ + ${{e}^{-}}$ -----> Cu
108 g of Ag is deposited by 96500 C
1.45 of Ag will be deposited by = 96500x1.45/108 C = 1295.6 C
Q = it or 1295.6 C = 1.5 A x t
t = 1295.6C/1.5 =====> 863.7 s
Since, the reaction related to the deposition of copper is
${{Cu}^{2+}}$ + ${{2e}^{-}}$ -----> Cu
2 x 96500 C electricity deposits 63.5 g of Cu
1295.6 C electricity will deposits 63.5 x 1295.6/2 x 96500 g of Cu, i.e. 0.426 g of Cu.
Since the reaction related to the deposition of zinc is
${{Zn}^{2+}}$ + ${{2e}^{-}}$ -----> Zn
2 x 96500 C electricity deposits 65.4 g of Zn.
1295.6 C electricity will deposit 65.4 x 1295.6/2 x 96500 g of Zn, i.e. 0.44 g of Zn.