The Haber Process

The Haber process makes ammonia, a key chemical used to produce fertilisers that help crops grow. It is a reversible reaction, so the system can go forwards and backwards.

The reaction and sources

Balanced reversible equation:

$$\mathrm{N_2(g)} + 3\,\mathrm{H_2(g)} \rightleftharpoons 2\,\mathrm{NH_3(g)} \quad (\text{exothermic})$$

• Nitrogen (N₂): from air (about 78% nitrogen).
• Hydrogen (H₂): from natural gas (methane) by reacting it with steam.

Reversible reactions and equilibrium

In a closed system, equilibrium happens when the forward and reverse reactions occur at the same rate, so the concentrations stay constant. It is like a busy two-way street with equal traffic in both directions.

Typical industrial conditions

• Temperature: 450 °C
• Pressure: 20 000 kPa ($200$ atm)
• Catalyst: iron

Why these conditions?

• Temperature: The forward reaction releases heat (exothermic). Lower temperature gives more ammonia but a slow rate. 450 °C is a compromise for a fast enough rate with reasonable yield.
• Pressure: On the left there are 4 gas molecules (1 N₂ + 3 H₂) and on the right 2 (NH₃). Higher pressure pushes equilibrium to fewer molecules, so more ammonia. 200 atm is chosen for good yield and safe, economical equipment.
• Catalyst: Iron speeds up both forward and reverse reactions equally. It does not change the position of equilibrium; it helps the plant reach equilibrium faster so a lower temperature can be used.
• Removal and recycling: Ammonia is cooled and liquefied, removed from the mixture. Unreacted N₂ and H₂ are recycled, which drives the overall yield higher.

Common misconceptions

• “Higher temperature always makes more product.” Not here: because the forward reaction is exothermic, higher temperature reduces ammonia yield (but increases rate).
• “Catalyst increases equilibrium yield.” It does not; it only helps reach equilibrium faster.