Reaction Pathways and Energy
What you must understand
Energy changes arise from bond breaking (endothermic) and bond forming (exothermic). Use ΔH = Σ(bonds broken) − Σ(bonds formed) and Hess's law with standard enthalpies of formation or combustion. Draw and interpret energy profile diagrams: activation energy, transition state, and the effect of catalysts in lowering Eₐ without changing ΔH. Understand that ΔH is a state function, so the pathway does not matter.
Reaction rate depends on collision frequency and energy; explain the effects of concentration, temperature, surface area, pressure and catalyst using collision theory. Know that temperature increases the proportion of collisions exceeding Eₐ, not just the collision count.
For reaction pathways, understand substitution, addition, elimination, oxidation and reduction, and the conditions and reagents that favour each. Recognise that reaction pathways are chosen for yield, selectivity, safety and atom economy. Entropy and Gibbs free energy (ΔG = ΔH − TΔS) determine spontaneity; a reaction with positive ΔH can still be spontaneous if ΔS is sufficiently positive at high temperature.
Traps to watch
- Catalysts do not change ΔH or equilibrium position - only the rate of approach.
- Bond energy signs: breaking is positive, forming is negative; a common sign error.
- Equilibrium vs rate: a large K does not mean a fast reaction.
- Temperature changes shift equilibrium (K changes); concentration changes do not change K.
Worked example
Using average bond energies: H–H 436, Cl–Cl 243, H–Cl 431 kJ mol⁻¹. For H₂ + Cl₂ → 2HCl: broken = 436 + 243 = 679; formed = 2 × 431 = 862. ΔH = 679 − 862 = −183 kJ mol⁻¹ (exothermic).
Rule of thumb: if ΔH is negative and ΔS is positive, the reaction is spontaneous at all temperatures.