Discussion
In an electrochemical cell, Ag+ ions in AgNO3 are reduced to Ag metal at the cathode and Cu is oxidized to Cu2+ at the anode. A current of 0.7 A is passed through the cell for 10 min. The mass (in grams) of silver deposited and copper dissolved, respectively, are:
[Faraday Constant = 96,485 C mol−1, Atomic Weight of Ag = 107.9, Atomic Weight of Cu = 63.55]
Answer: A
The masses are 0.469 g of silver and 0.138 g of copper.
Charge passed and moles of electrons:
\[ \begin{aligned} Q &= 0.7 \times 600 = 420\ \mathrm{C} \\ n_{e} &= \frac{420}{96485} = 4.353 \times 10^{-3}\ \mathrm{mol} \end{aligned} \]
Ag+ takes one electron and Cu gives two:
\[ \begin{aligned} m_{\mathrm{Ag}} &= 4.353 \times 10^{-3} \times 107.9 = 0.469\ \mathrm{g} \\ m_{\mathrm{Cu}} &= \frac{4.353 \times 10^{-3}}{2} \times 63.55 = 0.138\ \mathrm{g} \end{aligned} \]
0.069 g of copper divides by 2 twice, and 0.235 g of silver treats Ag+ as if it needed two electrons per atom.
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