Industrial Processes & EquilibriumOxford AQA IGCSE Chemistry: Revision notes

In a reversible reaction, the products of the forward reaction can react to reform the original reactants. When a reversible reaction happens in a closed system (nothing can enter or leave), the reaction reaches equilibrium: the forward and reverse reactions occur at exactly the same rate, so the relative amounts of reactants and products stay constant, even though both reactions are still happening.

The relative amounts of reactants and products at equilibrium depend on the conditions used.

• Temperature: raising the temperature increases the yield from the endothermic direction of the reaction and decreases the yield from the exothermic direction; lowering temperature has the reverse effect
• Pressure (gaseous reactions): increasing pressure favours the side of the reaction that produces the fewer molecules of gas (as shown by the balanced symbol equation); decreasing pressure favours the side producing the greater number of molecules

In industry, the conditions used for a reversible reaction (temperature, pressure and the use of a catalyst) are chosen as a compromise between three factors: maximising the yield of useful product, achieving a fast enough rate of reaction, and keeping costs (energy, equipment) as low as possible. The conditions that give the highest possible yield are not always used, if they would make the rate too slow or the process too expensive.

The Haber process manufactures ammonia from nitrogen and hydrogen: N₂ + 3H₂ ⇌ 2NH₃. It is reversible and reaches equilibrium in a closed system. The conditions used (iron catalyst, about 450 °C, about 200 atmospheres pressure) are a compromise: high pressure favours ammonia (fewer gas molecules on the product side) but is costly to generate; moderate (not very high or very low) temperature balances a reasonable yield against an acceptable rate, since the forward reaction is exothermic.

The Contact process manufactures sulfuric acid. Its key reversible step is 2SO₂ + O₂ ⇌ 2SO₃, catalysed by vanadium(V) oxide at about 450 °C and atmospheric pressure. Atmospheric pressure is used (rather than high pressure) because the reaction already gives a good yield of SO₃ at normal pressure, so the extra cost of generating high pressure is not worthwhile; the moderate temperature and catalyst give a fast enough rate at reasonable cost.