Thermochemistry and Thermodynamics
Core concepts
Thermochemistry tracks energy changes in reactions. Exothermic reactions release energy (ΔH negative); endothermic reactions absorb it (ΔH positive). Candidates must apply Hess's law, use standard enthalpy of formation and combustion data, and calculate bond-energy changes: ΔH ≈ energy of bonds broken − energy of bonds formed.
Calorimetry uses q = mcΔT to find heat transferred, then divides by moles to obtain molar enthalpy. Thermodynamics adds entropy (ΔS) and Gibbs free energy: ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG is negative; this depends on the signs of ΔH and ΔS and on temperature.
Traps to watch
- Sign errors: exothermic means negative ΔH, not positive.
- Forgetting to convert mass to kilograms or temperature change to kelvin where required.
- Mixing bond energies (positive, energy to break) with enthalpy of formation (negative for stable compounds).
- Assuming a spontaneous reaction is fast - thermodynamics says nothing about rate.
- Using Celsius in the TΔS term.
Worked example
Calculate ΔH for a reaction where bonds broken total 1200 kJ and bonds formed total 1500 kJ: ΔH = 1200 − 1500 = −300 kJ, so the reaction is exothermic. Rule of thumb: breaking bonds costs energy (endothermic); forming bonds releases it.