Calculate the current that must be passed into a solution of Aluminum salt for 1hr 30min. in order to deposit 1.5g of Aluminum (Al =27)
To calculate the current required to deposit 1.5g of Aluminum (Al), we need to use Faraday's laws of electrolysis.
First, we need to determine the moles of Aluminum deposited using its molar mass:
Molar mass of Aluminum (Al) = 27 g/mol
Moles of Aluminum (Al) = mass / molar mass = 1.5g / 27 g/mol = 0.0556 mol
Next, we need to determine the total charge required to deposit this amount of Aluminum. Since Aluminum has a charge of +3, each mole of Aluminum requires 3 moles of electrons.
Total charge (Coulombs) = moles of Aluminum (Al) * 3 * Faraday's constant
= 0.0556 mol * 3 * 96485 C/mol
= 16063 C
Since the time given is 1 hour 30 minutes, which is equal to 90 minutes, we need to convert it to seconds:
Time (s) = 90 minutes * 60 s/min = 5400 s
Finally, we can calculate the current required using the equation I = Q/t, where I is the current, Q is the charge, and t is the time:
Current (A) = Total charge (Coulombs) / Time (s)
= 16063 C / 5400 s
= 2.979 A
Therefore, the current that must be passed into the solution of Aluminum salt for 1hr 30min in order to deposit 1.5g of Aluminum is approximately 2.979 Amperes.