Energetics and Chemical Equilibrium
Core ideas
Exothermic reactions release energy to the surroundings (temperature rises); endothermic reactions absorb energy (temperature falls). Enthalpy change is the energy transferred at constant pressure; activation energy is the minimum energy needed to start a reaction. Bond breaking absorbs energy; bond making releases energy. Energy level diagrams show reactants, products and activation energy.
Reversible reactions can reach dynamic equilibrium in a closed system, where forward and reverse rates are equal and concentrations remain constant. Le Chatelier's principle predicts the effect of changing conditions: increasing temperature favours the endothermic direction; increasing pressure favours the side with fewer gas molecules; removing a product shifts equilibrium to the right. Catalysts do not shift equilibrium.
Traps to watch
- Equilibrium does not mean equal amounts; it means equal rates.
- A catalyst speeds up both directions equally and does not change the equilibrium position.
- Temperature changes alter the value of the equilibrium constant, but pressure changes do not.
- Exothermic reactions can still need activation energy to start.
Rule of thumb
For Le Chatelier: the system opposes the change. Heat it → favour the endothermic side; compress it → favour fewer gas molecules.
Worked example
N₂ + 3H₂ ⇌ 2NH₃ is exothermic. Increasing pressure favours the forward reaction because 4 gas molecules become 2, increasing ammonia yield.