Question

The corrosion of rebar inside concrete structures is serious issue for the safety and stability of bridges and buildings. The chemistry of corrosion is complex, but a simplified chemical model (iron rusting) is: 4Fe (s) + 3O2 (g) ⟶ 2Fe2O3 (s) ∆H = −1.65 × 103 kJ a) How much heat is released when 1.00 pound of iron rusts? b) If 6.25 × 104 kJ is released, how many kilograms of iron have rusted?

Answer 1

(a) Mass of iron rusted = 1.00 lb

mass of iron rusted = 1.00 lb * (453.6 g / 1.00 lb)

mass of iron rusted = 453.6 g

moles of iron rusted = (mass of iron rusted) / (molar mass Fe)

moles of iron rusted = (453.6 g) / (55.845 g/mol)

moles of iron rusted = 8.12 mol

heat released = (moles of iron rusted) * (H)

heat released = (8.12 mol) * (-1.65 x 10³ kJ / 4 mol Fe)

heat released = (8.12 mol / 4 mol) * (-1.65 x 10³ kJ)

heat released = -3350.5 kJ

(negative sign indicates heat is released by the reaction)

(b) Heat released = -6.25 x 10⁴ kJ

moles of Fe = (heat released) / (H)

moles of Fe = (-6.25 x 10⁴ kJ) / (-1.65 x 10³ kJ / 4 mol Fe))

moles of Fe = 37.9 mol

mass of Fe = (moles of Fe) * (molar mass Fe)

mass of Fe = (37.9 mol) * (55.845 g/mol)

mass of Fe = 2115.34 g

mass of Fe = 2115.34 g * (1 kg / 1000 g)

mass of Fe = 2.12 kg