Managing Chemical Processes: Rate, Extent and the Compromise Between Them
Choosing conditions deliberately
Managing a chemical process means selecting temperature, pressure, concentration and catalyst so a plant is fast enough to be profitable and selective enough to be worth running. You must understand collision theory and the Maxwell–Boltzmann distribution (why a modest temperature rise produces a large rate increase), activation energy, homogeneous and heterogeneous catalysts and their surfaces, and the measurement of rate from a changing physical property. On the equilibrium side you need the dynamic model, the Kc expression, the reaction quotient, Le Chatelier's principle and the thermodynamic framework: ΔH, ΔS and ΔG = ΔH − TΔS. Industrial cases - Haber, Contact, methanol synthesis - are examined as compromises, never as single correct conditions.
Traps
- Claiming a catalyst increases yield. It increases the rate of approach to the same equilibrium position; Kc is unchanged.
- Forgetting that only temperature changes the value of Kc.
- Applying pressure effects to solids, liquids or aqueous species.
- Confusing rate with extent: a fast reaction may still give a poor yield.
- Getting the sign of ΔG wrong, or assuming a spontaneous reaction must be a fast one.
Worked example
For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ mol⁻¹, predict the effect of raising temperature and of raising pressure.
Raising temperature shifts an exothermic equilibrium left, lowering Kc and the equilibrium yield of ammonia - but it raises the rate, so the industrial compromise sits near 450 °C with an iron catalyst. Raising pressure favours the side with fewer gas molecules (four moles of gas become two), shifting the position right and increasing yield; the limit is plant cost and safety. Rule of thumb: count only gas moles when applying pressure.