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Electrochemistry — Chemistry STD 12 Science — Question

Rajasthan BoardEnglish MediumSTD 12 ScienceChemistryElectrochemistry5 Marks
Question
Three electrolytic cells A, B, C containing solutions of ZnSO4, AgNO3 and CuSO4, 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?
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Answer

The following reaction is used to deposit silver on the cathode of cell B :
$Ag ^{+}+ e ^{-} \rightarrow Ag$
Atomic mass of 1 mole of Ag = 108 g
Amount of electricity required deposition of 108 g Ag = 96500 C
Therefore, for deposition of 1.45 g silver Ag
$=\frac{96500}{108} \times 1.45$
$=1295.6 C$ amount of electricity is required.
Charge $Q = It , t =\frac{ Q }{ I }=\frac{1295.6}{1.5}=863.73$ second.
Hence, time in minutes $=\frac{863.73}{60}=14.39$
$=14.40$ minutes.
Hence, electric current flowed for 14.40 minutes.
The reaction for deposition of Cu :
$Cu ^{2+}+2 e ^{-} \rightarrow Cu$
Atomic mass of 1 mole Cu = 63.5 g
Hence, copper obtained from 2 × 96500 C = 63.5 g
Copper obtained from 1295.6 C electricity
$\begin{array}{l}=\frac{63.5 \times 1295.6}{2 \times 96500} \\ =0.426 g\end{array}$
Hence, mass of deposited copper = 0.426 g
Reaction required for deposition of Zn :
$Zn ^{2+}+2 e ^{-} \rightarrow Zn$
Atomic mass of 1 mole Zn = 65 g
Zinc obtained from 2 × 96500C = 65 g
Zn obtained from 1295.6 C electricity
$=\frac{65 \times 1295.6}{2 \times 96500}=0.436 g$
Therefore, mass of deposited zinc = 0.436 g.

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