How many Faradays of electricity are required to deposit $10g$ of calcium from Calcium Chloride (${\mathrm{CaCl}}_2$) using inert electrodes? (Molar mass of Calcium= $40gmo{l}^{-1}$)

Step 1: Given Data

• Given mass of Calcium= $10\mathrm{g}$
• Molar mass of Calcium= $40\mathrm{g}{\mathrm{mol}}^{-1}$

Step 2: Calculation of number of moles of Calcium

Formula for number of moles= $\frac{\mathrm{Given}\mathrm{Mass}}{\mathrm{Molar}\mathrm{Mass}}$

Substituting the given data in above formula:

$$\begin{gathered} \frac{10\mathrm{g}}{40\mathrm{g}{\mathrm{mol}}^{-1}} \\ =0.25\mathrm{mol} \end{gathered}$$

Number of moles of calcium= $0.25mol$

Step 3: Calculation of Faradays of Electricity

The equation is written as: $\underset{\mathrm{Calcium}\mathrm{ion}}{{\mathrm{Ca}}^{2+}\left(\mathrm{aq}\right)}+2e^{-}\to \underset{\mathrm{Calcium}}{\mathrm{Ca}\left(\mathrm{s}\right)}$

$1\mathrm{mole}\mathrm{Ca}$ requires= $2\mathrm{moles}$ of electrons= $2$ Faradays of electricity

$0.25\mathrm{mole}\mathrm{Ca}$= $2\times 0.25$= $0.50\mathrm{moles}\mathrm{of}\mathrm{electrons}$= $0.5$ Faradays of Electricity

Hence, $0.5$ faradays of electricity are required to deposit $10g$ of calcium from molten Calcium Chloride (${\mathrm{CaCl}}_2$) using inert electrodes.