Question

electrolytic cell

6A Practical Application of the electrolysis process. How much does it cost to generate the mass of aluminum (14.0 g) in one beverage can The aluminum in the can is produced by reducing Al+ to Al(s). The reaction is run commercially at 50,000 A and 4.0 V (4.0 J/C), 1 kWh ofelectricity cost about 10 cents and 1 kWh is 3.60 x 10s J. Energy charge x voltage.

Answer 1

Q6. The chemical equation for formation of solid aluminum from reduction of Al³⁺ is :

Al³⁺ + 3 e^- $\small \rightarrow$ Al (s)

mass Al = 14.0 g

moles Al = (mass Al) / (molar mass Al)

moles Al = (14.0 g) / (26.98 g/mol)

moles Al = 0.519 mol

moles of electron = 3 * (moles Al)

moles of electron = 3 * (0.519 mol)

moles of electron = 1.5567 mol e^-

Charge consumed = (moles of electron) * (Faraday's constant)

Charge consumed = (1.5567 mol e^-) * (96490 C/mol e^-)

Charge consumed = 150206.82 C

Energy consumed = (Charge consumed) * (voltage)

Energy consumed = (150206.82 C) * (4.0 J/C)

Energy consumed = 600827.3 J

electricity consumed = (Energy consumed) / (3.60 x 10⁶ J/KWh)

electricity consumed = (600827.3 J) / (3.60 x 10⁶ J/KWh)

electricity consumed = 0.1669 kWh

Cost of electricity = (electricity consumed) * (electricity rate)

Cost of electricity = (0.1669 kWh) * (10 cents / 1 kWh)

Cost of electricity = 1.67 cents