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 ampere was passed through them until 1.45 g of silver were deposited at the cathode of cell B. How long did the current flow? What mass of copper and what mass of zinc were deposited in the concerned cells? (Atomic masses of Ag = 108, Zn = 65.4, Cu = 63.5)

Explanation

`Ag^(+)+e^(-)rarrAg`
`because` 108 g of Ag is deposited by 1 F = 96500C
`therefore` 1.45g of Ag is deposited by 1 F = `(96500)/(108)xx1.45 = 1295.6 C`
`therefore Q = It`
`therefore t = (Q)/(I) = (1295.6)/(1.5) = 863.75 sec`.
Now, `Cu^(2+) + 2e^(-) rarr Cu`
`therefore` Copper, Cu deposited = `(63.5)/(2xx96500)xx1295.6 = 0.426g`
Also, `Zn^(2+)+2e^(-)rarrZn`
`therefore` Zinc, Zn deposited = `(65.4)/(2xx96500)xx1295.6 = (84732.24)/(193000) = 0.439g`