How many coulombs of charge are required to deposit $108$ g of silver from a $\mathrm{AgNO_3}$ solution? (Atomic mass of Ag = 108 g/mol, $F = 96{,}500$ C/mol)

**Setup.** $\mathrm{Ag^+ + e^- \to Ag}$ means one mole of electrons (= 1 Faraday) deposits one mole of Ag. **Moles of Ag** = $\dfrac{108}{108} = 1$ mol. **Charge required** = $1 \times 96{,}500 = 96{,}500$ C (= 1 F). If the metal had been Cu ($\mathrm{Cu^{2+} + 2e^- \to Cu}$), the same mass would need 2 F instead, because each Cu needs 2 electrons. Option C ($1{,}93{,}000$ C) is the trap if you mistakenly use $n = 2$ for Ag.