Thermochemistry and Energy Changes
Enthalpy basics
Enthalpy change ΔH is the heat exchanged at constant pressure. Exothermic reactions release heat (ΔH negative); endothermic absorb heat (ΔH positive). Enthalpy of combustion, formation and neutralisation are defined under standard conditions. Enthalpy of formation of an element in its standard state is zero.
Calculating ΔH
Use Hess's law: ΔH is independent of route. Combine equations algebraically, reversing signs when equations are reversed and scaling when multiplied. Alternatively use ΔH = ΣΔHf(products) − ΣΔHf(reactants), or bond energies: ΔH = Σ bonds broken − Σ bonds formed. Bond breaking is endothermic (positive), bond forming exothermic (negative).
Calorimetry uses q = mcΔT, where m is the mass of solution, c is 4.18 J g⁻¹ K⁻¹ and ΔT is the temperature change. Convert q to kJ and divide by moles to obtain molar enthalpy. Heat loss to surroundings makes measured values smaller in magnitude than literature values.
Traps
- Sign errors when reversing equations.
- Using the mass of solute instead of the total solution in q = mcΔT.
- Forgetting to convert J to kJ.
- Mixing bond energy and enthalpy of formation methods.
- Ignoring heat loss and heat capacity of the calorimeter.
Rule of thumb
Draw an energy cycle with the elements at the base. Arrows down are formation enthalpies; the difference between routes gives ΔH.
Worked example
Burning 0.50 g of methanol (M = 32 g/mol) raises 100 g of water by 10.0 °C. q = 100 × 4.18 × 10.0 = 4180 J = 4.18 kJ. n = 0.0156 mol. Molar enthalpy ≈ −268 kJ/mol (exothermic).