Topic 4.2: Energy Cycles in Reactions Questions

Explain why a value of ΔH calculated from average bond enthalpies usually differs slightly from the value measured experimentally.

Calculate the standard enthalpy change for the combustion of ethanol, C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l), using ΔH_f⊖: C₂H₅OH(l) = −277.0, CO₂(g) = −393.5, H₂O(l) = −285.8 kJ mol⁻¹.

Explain how Hess's law allows the enthalpy change of a reaction to be determined even when it cannot be measured directly.

Carbon monoxide reacts with chlorine to form phosgene: CO(g) + Cl₂(g) → COCl₂(g).

Using the bond enthalpies C≡O = 1077, Cl–Cl = 242, C=O = 745 and C–Cl = 324 kJ mol⁻¹, determine the enthalpy change for the reaction.

The standard enthalpy of combustion of methanol, CH₃OH(l), is −726 kJ mol⁻¹ and of hydrogen, H₂(g), is −286 kJ mol⁻¹.

Use these values to calculate ΔH⊖ for CO₂(g) + 3H₂(g) → CH₃OH(l) + H₂O(l).

(CO₂ and H₂O do not combust.)

Determine the enthalpy change for the reaction C(s) + 2H₂(g) → CH₄(g), given the combustion data: C(s) + O₂(g) → CO₂(g), ΔHc = −394 kJ mol⁻¹; H₂(g) + ½O₂(g) → H₂O(l), ΔHc = −286 kJ mol⁻¹; CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), ΔHc = −890 kJ mol⁻¹.