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

1.8 × 10² s

Step-by-step 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 * (1gCu)/(1,000mgCu) * (1molCu)/(63.55gCu) * (2mole^(-) )/(1molCu) * (94,486C)/(1mole^(-)) * (1s)/(5.6C) = 1.8 * 10^(2) s`