The heat of formation, more precisely called the standard enthalpy of formation, tells how much heat is absorbed or released when 1 mole of a compound forms from its elements in their standard states. It matters because chemists can use tabulated formation enthalpies to predict the heat change of many reactions without measuring each one directly. A negative value means the formation process releases heat, while a positive value means it absorbs heat.
These values are central to thermochemistry, fuel analysis, materials chemistry, and reaction design.
Standard formation enthalpies are measured relative to elements in their most stable forms at standard conditions, so those elements are assigned ΔHf° = 0 kJ/mol. For a reaction, the overall enthalpy change is found by subtracting the total formation enthalpy of the reactants from that of the products, using the balanced chemical equation as the mole ratio. This works because enthalpy is a state function, so only the starting and ending chemical states matter.
A worked example often uses ΔHrxn° = ΣnΔHf°(products) - ΣnΔHf°(reactants), where n is the coefficient from the balanced equation.
Key Facts
- Standard enthalpy of formation is written ΔHf° and has units of kJ/mol.
- ΔHf° is the enthalpy change when 1 mol of a compound forms from its elements in their standard states.
- Elements in their standard states have ΔHf° = 0 kJ/mol, such as O2(g), H2(g), N2(g), C(s, graphite), and Na(s).
- Reaction enthalpy from formation data: ΔHrxn° = ΣnΔHf°(products) - ΣnΔHf°(reactants).
- Use coefficients from the balanced equation to multiply each substance's ΔHf° value.
- If ΔHrxn° < 0, the reaction is exothermic; if ΔHrxn° > 0, the reaction is endothermic.
Vocabulary
- Enthalpy
- Enthalpy is a thermodynamic quantity related to the heat absorbed or released by a system at constant pressure.
- Standard enthalpy of formation
- Standard enthalpy of formation is the heat change when 1 mole of a compound forms from its elements in their standard states.
- Standard state
- The standard state is the reference form of a substance, usually at 1 bar pressure and a specified temperature such as 25°C.
- State function
- A state function depends only on the initial and final states of a system, not on the path taken between them.
- Exothermic reaction
- An exothermic reaction releases heat to the surroundings and has a negative enthalpy change.
Common Mistakes to Avoid
- Forgetting to balance the equation first is wrong because formation enthalpies must be multiplied by the correct mole coefficients.
- Using elements with nonzero ΔHf° values is wrong when the element is in its standard state, since standard-state elements are assigned ΔHf° = 0 kJ/mol.
- Reversing products and reactants in the formula is wrong because ΔHrxn° equals products minus reactants, not reactants minus products.
- Ignoring physical states is wrong because the same chemical formula can have different ΔHf° values in different states, such as H2O(l) and H2O(g).
Practice Questions
- 1 Calculate ΔHrxn° for CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l), using ΔHf° values: CH4(g) = -74.8 kJ/mol, CO2(g) = -393.5 kJ/mol, H2O(l) = -285.8 kJ/mol, O2(g) = 0 kJ/mol.
- 2 Calculate ΔHrxn° for 2CO(g) + O2(g) -> 2CO2(g), using ΔHf° values: CO(g) = -110.5 kJ/mol, CO2(g) = -393.5 kJ/mol, O2(g) = 0 kJ/mol.
- 3 Explain why graphite has ΔHf° = 0 kJ/mol but diamond does not, even though both are forms of pure carbon.