Reversible Reactions, Equilibrium and the Haber Process
Reversible Reactions
Some chemical reactions can proceed in both directions. In a reversible reaction, the products can react together to re-form the original reactants. These reactions are represented using a double arrow (\(\rightleftharpoons\)) rather than a single arrow. A straightforward example is the thermal decomposition of ammonium chloride:
$$NH_4Cl \rightleftharpoons NH_3 + HCl$$
Heating drives the forward reaction; cooling drives the reverse. The energy is conserved: if the forward reaction is exothermic, the reverse reaction is endothermic by exactly the same amount.
In a reversible reaction, if the forward reaction is exothermic, the reverse reaction is endothermic by the same amount of energy — and vice versa.
Dynamic Equilibrium
When a reversible reaction takes place in a closed system (nothing enters or leaves), it eventually reaches dynamic equilibrium. At equilibrium:
- The rate of the forward reaction equals the rate of the reverse reaction.
- The concentrations of reactants and products remain constant (they are not necessarily equal).
- Both reactions are still happening — equilibrium is dynamic, not static. The reaction has not stopped.
At dynamic equilibrium, both forward and reverse reactions continue at the same rate. Concentrations stay constant because the rate of product formation equals the rate of product being converted back to reactants. The reaction has NOT stopped.
Le Chatelier’s Principle
Le Chatelier’s principle states: if conditions are changed in a system at equilibrium, the equilibrium shifts in the direction that opposes the change.
Temperature: increasing temperature shifts equilibrium in the endothermic direction (absorbs the extra heat). Decreasing temperature shifts it in the exothermic direction (releases heat to compensate). Temperature is the only factor that changes the value of the equilibrium constant.
Concentration: increasing the concentration of a reactant shifts equilibrium to the right (forward reaction) to use up the extra reactant. Increasing the concentration of a product shifts equilibrium to the left.
Pressure (gases only): increasing pressure shifts equilibrium towards the side with fewer moles of gas, to reduce the pressure. Decreasing pressure shifts it towards the side with more moles of gas.
The Haber Process
The Haber process manufactures ammonia from nitrogen and hydrogen:
$$N_2 + 3H_2 \rightleftharpoons 2NH_3 \quad (\Delta H = -92 \text{ kJ/mol})$$
The forward reaction is exothermic. The nitrogen comes from the air; the hydrogen is obtained mainly from natural gas (methane). The reaction is carried out at approximately 450°C and 200 atmospheres pressure, using a finely divided iron catalyst.
Temperature — rate vs yield: the forward reaction is exothermic. Lowering temperature would favour the forward reaction (better yield, around 98% at room temperature), but the rate would be too slow to be economical. Higher temperatures speed up the reaction but reduce yield. 450°C is a compromise — an acceptable yield (around 15–25%) at a commercially viable rate.
Pressure — more moles on the left: on the left there are 4 moles of gas (1 N2 + 3 H2); on the right there are 2 moles (2 NH3). By Le Chatelier’s principle, high pressure shifts equilibrium to the right (fewer moles of gas). Very high pressures give better yields but are extremely expensive and dangerous to maintain. 200 atmospheres is again a compromise: a high enough pressure to be practical without prohibitive engineering costs.
Iron catalyst: lowers the activation energy of both the forward and reverse reactions equally, so the system reaches equilibrium faster without changing the position of equilibrium. It does not improve yield, only rate.
Because the yield per pass is only around 15–25%, the unreacted nitrogen and hydrogen are separated from the ammonia and recycled back into the reactor. The ammonia is cooled and liquefied (it has a higher boiling point than N2 and H2), allowing the unreacted gases to remain as gas and be pumped back. Overall efficiency is very high despite the low single-pass yield.
Ammonia and Fertilisers (Triple-only)
Ammonia manufactured by the Haber process is used to make nitrogen-based fertilisers. Ammonia can be oxidised to form nitric acid, which then reacts with more ammonia to make ammonium nitrate, a widely used fertiliser. NPK fertilisers contain nitrogen, phosphorus, and potassium in ratios designed to optimise crop growth. They are formulated by combining ammonium salts with phosphate and potassium compounds.