Thermochemistry and Energetics
What you must understand
Distinguish system from surroundings, and exothermic from endothermic processes in terms of sign conventions for q and enthalpy. Use calorimetry: q = mcΔT for solutions and q = CΔT for a calorimeter. Apply Hess's law, standard enthalpies of formation, and bond enthalpies to compute reaction enthalpy. Understand state functions, internal energy, and the first law of thermodynamics, including work of expansion. Know the meaning of entropy and Gibbs free energy, and how the sign of ΔG predicts spontaneity, with ΔG = ΔH - TΔS.
Traps to watch for
- Sign errors: exothermic reactions have negative ΔH, and heat leaving the system is negative q.
- Forgetting to reverse the sign of ΔH when reversing a reaction in Hess's law, or to scale it with coefficients.
- Using bond enthalpies of reactants minus products instead of products minus reactants for atomisation-based bond energy calculations.
- Assuming a spontaneous reaction must be fast; thermodynamics says nothing about rate.
- Mixing up ΔH of formation for elements in their standard states, which is zero.
Worked example
Estimate ΔH for the combustion of methane using bond enthalpies: CH4 + 2O2 gives CO2 + 2H2O. Bonds broken: four C-H (about 413 each) and two O=O (about 498), totalling about 2648 kJ. Bonds formed: two C=O (about 799 each) and four O-H (about 463 each), totalling about 3450 kJ. ΔH is bonds broken minus bonds formed, about -802 kJ, consistent with an exothermic reaction.
Rule of thumb
For Gibbs energy, compute ΔH and ΔS separately, then combine at the given temperature in kelvin.