How many faraday of electricity is required to produce 0.25 mole of copper?

Analysis

To solve this problem, we need to understand the relationship between the amount of electricity (in Faradays) and the number of moles of metal deposited during electrolysis. This is based on Faraday's First Law of Electrolysis.

Concept

The reduction of copper ions ($Cu^{2+}$) to copper metal ($Cu$) occurs according to the following half-equation:

From this equation, we can see that:

• 1 mole of copper requires 2 moles of electrons to be deposited.
• Since 1 mole of electrons carries a charge of 1 Faraday (F), it follows that 1 mole of copper requires 2 Faradays of electricity.

Solution

Step 1: Identify the stoichiometry. As shown in the equation above, 1 mole of $Cu$ requires 2 Faradays of electricity.

Step 2: Set up the calculation. We are asked to find the electricity required for 0.25 mole of copper. We use the ratio:

Step 3: Calculate the result.

Conclusion

The amount of electricity required to produce 0.25 mole of copper is 0.50 Faraday. This matches option D.

Exam Tip

Always write out the half-equation for the metal being deposited. The coefficient of the electrons ($e^-$) tells you exactly how many Faradays are needed to produce 1 mole of that metal. For monovalent ions (like $Ag^+$), it's 1F; for divalent ions (like $Cu^{2+}$ or $Zn^{2+}$), it's 2F; and for trivalent ions (like $Al^{3+}$), it's 3F.