Chemistry · General chemistry II · Worked example

Find the mass plated by electrolysis

A current of 2.00 A flows for 1.00 h through a solution of Cu²⁺. What mass of copper is deposited?

Find the charge

1.00 h = 3600 s, and an ampere is one coulomb per second.

q = (2.00 A) (3600 s) = 7200 C

Convert to moles of electrons

Divide by the Faraday constant.

n_{e^{-}} = \frac{7200 C}{96485 C/mol} = 0.07462 mol

\begin{array}{c} n_{e^{-}} = \frac{7200 C}{96485 C/mol} \\ = 0.07462 mol \end{array}

Use the half-reaction

Cu²⁺ + 2e⁻ → Cu: two moles of electrons deposit one mole of copper.

n_{Cu} = \frac{0.07462 mol}{2} = 0.03731 mol

Convert to mass

Copper’s molar mass is 63.55 g/mol.

m = (0.03731 mol) (63.55 g/mol) = 2.371 g

\begin{array}{c} m = (0.03731 mol) (63.55 g/mol) \\ = 2.371 g \end{array}

Result

2.37 g of copper.

Your turn

How long must a current of 1.50 A flow to plate 1.00 g of silver (Ag⁺ + e⁻ → Ag)?

Show the answer and explanation

About 596 s, or 9.94 min.

1.00 g ÷ 107.87 g/mol = 9.270 × 10⁻³ mol of silver, which needs as many moles of electrons: 894 C. Then t = q/I = 894 C ÷ 1.50 A = 596 s.

\frac{1.00 (96485)}{107.87 (1.50)} \approx 596.3

Keep exploring

In Electrochemistry & charge, change the product to silver, with one electron per atom. The same charge plates 8.05 g of silver: twice the atoms, each heavier.

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