Standard Enthalpy of Combustion ($ΔH_c^\circ$) and Formation ($ΔH_f^\circ$)
Standard Enthalpy of Combustion ($ΔH_c^\circ$)
Standard enthalpy of combustion
The standard enthalpy of combustion ($ΔH_c^\circ$) is the energy change when one mole of a substance is completely burned in oxygen under standard conditions (298 K and 1 atm).
The products of combustion must also be in their standard states (e.g., CO₂ as a gas and H₂O as a liquid).
ExampleThe combustion of methane (CH₄) is represented as:
$$\text{CH}_4(g) + 2\text{O}_2(g) \to \text{CO}_2(g) + 2\text{H}_2O(l) \quad \Delta H_c^\circ = -890 \, \text{kJ mol}^{-1}$$
This equation tells us that burning 1 mole of methane releases 890 kJ of energy.
Key Features of $ΔH_c^\circ$
- Exothermic Nature: Combustion reactions are always exothermic, meaning $ΔH_c^\circ$ values are negative because energy is released to the surroundings.
- Standard States: All reactants and products must be in their standard states.
- Applications: $ΔH_c^\circ$ values are essential for comparing the energy efficiency of different fuels, such as coal, gasoline, and hydrogen.
Oxygen is a gas (O₂(g)), and water is a liquid (H₂O(l)) at standard conditions.
ExampleCombustion of Butane
The combustion of butane (C₄H₁₀), a fuel commonly used in lighters, is represented as:
$$\text{C}4\text{H}{10}(g) + \frac{13}{2}\text{O}_2(g) \to 4\text{CO}_2(g) + 5\text{H}_2O(l) \quad \Delta H_c^\circ = -2878 \, \text{kJ mol}^{-1}$$
This means burning 1 mole of butane releases 2878 kJ of energy.
Tip- $ΔH_c$ values for common substances are listed in Section 14 of the IB Chemistry Data Booklet.
- Use these values for accurate calculations.
