The oxidation of iron has a change of enthalpy ΔH⁰_rxn = -1648 kJ.mol^-1 and may be represented as:

4 Fe (s) + 3 O₂(g) → 2 FeO₃(s)

1) Is this reaction endothermic or exothermic?

2) Calculate the energy released as heat when 25.0 g of iron undergo oxidation at constant pressure.

Question

The oxidation of iron has a change of enthalpy ΔH⁰_rxn = -1648 kJ.mol^-1 and may be represented as:

4 Fe (s) + 3 O₂(g) → 2 FeO₃(s)

1) Is this reaction endothermic or exothermic?

2) Calculate the energy released as heat when 25.0 g of iron undergo oxidation at constant pressure.

| General Chemistry 2

Accepted Answer

1) ΔH⁰_rxn < 0 ⇒ the reaction is exothermic

2) ΔH⁰_rxn = - 1648 kJ.mol^-1 ⇒ energy released per mole of the stated reaction equation

4 molecules of iron react during the reaction

⇒ -1648 kJ per 4 moles of Fe(s) is released as heat

n_Fe=m_Fe/ M_Fe = 25.0 / 55.8 = 4.48 x 10^-1 mol

Energy released = ΔH = $\frac{∆ {H^{0}}_{rxn}}{4 n_{Fe}}$

⇒ ΔH =

\frac{- 1648 \times 10^{3}}{4 \times 4.48 \times 10^{- 1}}

⇒ ΔH = - 1.85 x 10⁵ J = - 185 kJ

Exercise 9 | Thermochemistry