Chemical reactions: rates, energetics and equilibrium
Chemical reactions: rates, energetics and equilibrium tests collision theory, enthalpy changes and reversible reactions.
Rates and collision theory
Reaction rate increases with higher concentration, higher temperature, greater surface area and a catalyst. Collision theory explains this: particles must collide with at least the activation energy and in the correct orientation. A catalyst provides an alternative pathway with lower activation energy. Increasing temperature increases both collision frequency and the proportion of collisions with enough energy, so the effect is usually large.
Energetics
Exothermic reactions release energy to the surroundings and have negative ΔH; endothermic reactions absorb energy and have positive ΔH. Bond breaking requires energy (endothermic); bond forming releases energy (exothermic). The overall enthalpy change depends on the balance between bonds broken and bonds formed.
Equilibrium
A reversible reaction reaches dynamic equilibrium in a closed system when forward and reverse rates are equal and concentrations remain constant. Le Chatelier's principle predicts the shift when conditions change: increasing pressure favours the side with fewer gas molecules; increasing temperature favours the endothermic direction; removing a product shifts the reaction forward. A catalyst does not change the position of equilibrium.
Worked example
For N2 + 3H2 ⇌ 2NH3, ΔH is negative. If the temperature is increased, the equilibrium shifts in the endothermic direction, which is to the left, so the yield of ammonia decreases. The rate of both reactions increases, so equilibrium is reached faster but with a lower yield. This is why the Haber process uses a compromise temperature.
Traps
A catalyst does not increase yield. Equilibrium does not mean equal concentrations; it means constant concentrations with equal rates. Activation energy is not the same as ΔH. An endothermic reaction in solution cools the mixture; an exothermic one warms it.