A current of 1.50 amp was passed through an electrolytic cell containing AgNO3 solution with inert electrodes. The weight of Ag deposited was 1.50g. How long did the current flow? (2marks)
(b) Write the reactions taking place at the anode and cathode in the above cell.
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12th
Chemistry
Electrochemistry
Electrolytic Cells and Electrolysis
Three electrolytic cells A,...
CHEMISTRY
Three electrolytic cells A,B,C containing solutions of ZnSO
4
,AgNO
3
and CuSO
4
, respectively are connected in series. A steady current of 1.5 amperes was passed through them until 1.45 g of silver deposited at the cathode of cell B. How long did the current flow? What mass of copper and zinc were deposited?
November 22, 2019avatar
Puneetha Adak
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VIDEO EXPLANATION
ANSWER
Cell B: Ag
+
+e
−
⇌Ag at cathode.
1 mole (108 g) of Ag is deposited by 96500 C.
1.45 g of Ag will be deposited by
108
96500×1.45
=1295.6C.
Q=It
1295.6=1.5×t
t=864s
Cell A: Zn
2+
+2e
−
→Zn
2 moles of electrons (2×96500 C of current) produces 1 mole (63.5 g) of zinc.
1295.6 C of electricity will deposit
2×96500
65.3
×1295.6=0.438 g of zinc
Cell C: Cu
2+
+2e
−
→Cu
2 moles of electrons (2×96500 C) of current will produce 1 mole (63.5 g) of Cu.
1295.6 C of current will deposit
2×96500
63.5×1295.6
=0.426g of copper