Chemistry 2001 JAMB Past Questions - Question 26: What current in amperes will deposit 2.7g of aluminium in 2 hours?

Correct Answer

C

Explanation

To determine the current required to deposit a certain amount of a substance during electrolysis, you can use Faraday's laws of electrolysis. The equation that relates the amount of substance deposited, the current, the time, and the molar mass is given by:

\[ \text{Amount of substance deposited (in moles)} = \frac{\text{Current (in amperes)} \times \text{Time (in seconds)}}{\text{Faraday's constant}} \]

For aluminum, the molar mass is approximately \(27 \, \text{g/mol}\), and Faraday's constant is \(96485 \, \text{C/mol}\). Given that you want to deposit \(2.7 \, \text{g}\) of aluminum in \(2 \, \text{hours}\) (which is \(7200 \, \text{seconds}\)), you can rearrange the formula to solve for the current:

\[ \text{Current} = \frac{\text{Amount of substance deposited} \times \text{Faraday's constant}}{\text{Time}} \]

Substitute the known values:

\[ \text{Current} = \frac{(2.7 \, \text{g}) \times (1 \, \text{mol/27 \, g}) \times (96485 \, \text{C/mol})}{7200 \, \text{s}} \]

Calculating this expression will give you the current in amperes required to deposit \(2.7 \, \text{g}\) of aluminum in \(2 \, \text{hours}\).