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Answer :
Final answer:
The current required to produce 20.0 g of aluminum from an AlCl3 solution in 100 minutes is approximately 11.9 A.
Explanation:
To calculate the current required to produce 20.0 g of aluminum from an AlCl3 solution in 100 minutes, we can use Faraday's laws of electrolysis. Faraday's first law states that the amount of substance produced at an electrode is directly proportional to the quantity of electricity passed through the electrolyte. Faraday's second law states that the amount of substance produced at an electrode is directly proportional to the molar mass of the substance.
First, we need to calculate the number of moles of aluminum using its molar mass. The molar mass of aluminum is 26.98 g/mol, so:
Number of moles of aluminum = mass of aluminum / molar mass of aluminum
Number of moles of aluminum = 20.0 g / 26.98 g/mol
Next, we need to calculate the charge required to produce this amount of aluminum. Each aluminum ion (Al3+) requires 3 electrons to be reduced to aluminum. Therefore, the charge required can be calculated using the equation:
Charge required = number of moles of aluminum * 3 * Faraday's constant
Faraday's constant is 96,485 C/mol. Substituting the values, we get:
Charge required = (20.0 g / 26.98 g/mol) * 3 * 96,485 C/mol
Finally, we can calculate the current required using the equation:
Current = Charge required / time
Substituting the values, we get:
Current = [(20.0 g / 26.98 g/mol) * 3 * 96,485 C/mol] / 100 minutes
Calculating this, we find that the current required is approximately 11.9 A.
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