Question

How much time would it take for 336 mg of copper to be plated at a current of 5.6 A ? Express your answer using two significant figures.

Answer 1

Answer:

1.8 × 10² s

Explanation:

Let's consider the reduction that occurs upon the electroplating of copper.

Cu²⁺(aq) + 2 e⁻ ⇒ Cu(s)

We will establish the following relationships:

• 1 g = 1,000 mg
• The molar mass of Cu is 63.55 g/mol
• When 1 mole of Cu is deposited, 2 moles of electrons circulate.
• The charge of 1 mole of electrons is 96,486 C (Faraday's constant).
• 1 A = 1 C/s

The time  that it would take for 336 mg of copper to be plated at a current of 5.6 A is:

$336mgCu \times \frac{1gCu}{1,000mgCu} \times \frac{1molCu}{63.55gCu} \times \frac{2mole^{-} }{1molCu} \times \frac{94,486C}{1mole^{-}} \times \frac{1s}{5.6C} = 1.8 \times 10^{2} s$