Question
Apply Le Chatelier’s principle to the Haber process for ammonia synthesis. Explain the effect of temperature, pressure, and concentration on the equilibrium. Why are the actual industrial conditions (450°C, 200 atm) a compromise?
(NCERT Class 11, Equilibrium)
Solution — Step by Step
Key features: exothermic (negative ), moles of gas decrease (4 moles → 2 moles), reversible.
Le Chatelier’s principle: Increasing pressure shifts equilibrium towards the side with fewer gas moles.
Reactant side: 1 + 3 = 4 moles of gas. Product side: 2 moles of gas.
High pressure favours ammonia production. At 200 atm, the yield of NH₃ is significantly higher than at 1 atm. Even higher pressures would give better yields, but they’re impractical due to equipment costs and safety risks.
The forward reaction is exothermic (). By Le Chatelier’s principle:
- Low temperature shifts equilibrium towards products (favours the exothermic direction)
- High temperature shifts equilibrium towards reactants
So thermodynamically, low temperature gives higher yield. But kinetically, low temperature means an extremely slow reaction rate — you’d wait years for equilibrium.
450°C is the compromise: fast enough rate with a reasonable yield (~15-20% NH₃ per pass). The unreacted gases are recycled.
Removing NH₃ as it forms (by condensation) shifts equilibrium forward → more NH₃ produced. This is why the product is continuously removed in the industrial process.
Adding excess N₂ or H₂ also shifts equilibrium forward.
Catalyst (finely divided iron with Al₂O₃ and K₂O promoters): Does NOT change the equilibrium position or yield. It only increases the rate at which equilibrium is reached — making it possible to get useful rates at 450°C instead of needing even higher temperatures.
Why This Works
Le Chatelier’s principle states that if a system at equilibrium is disturbed, it shifts in the direction that partially counteracts the disturbance. For the Haber process:
- Increase pressure → system shifts to reduce pressure → moves toward fewer gas moles → forward reaction
- Increase temperature → system shifts to absorb heat → moves toward endothermic direction → reverse reaction
- Remove product → system shifts to replace removed substance → forward reaction
The industrial conditions are a balance between thermodynamic favourability (low T, high P for maximum yield) and kinetic practicality (high T for faster rate, moderate P for equipment safety). The catalyst helps bridge this gap by allowing reasonable rates at a “moderate” temperature.
Alternative Method
You can also analyse this quantitatively using the equilibrium constant. for an exothermic reaction decreases with increasing temperature (van’t Hoff equation). At 200°C, is much larger than at 500°C — but the reaction at 200°C is so slow it’s useless without an extraordinary catalyst.
This is a guaranteed question in CBSE board exams (3-5 marks). The expected structure: write the equation, apply Le Chatelier’s principle for each factor (T, P, concentration, catalyst), state the actual industrial conditions (450°C, 200 atm, Fe catalyst), and explain why they’re a compromise. NEET asks this as an MCQ — “Increasing pressure in the Haber process will ___?”
Common Mistake
Students often write “catalyst shifts equilibrium towards products.” This is wrong. A catalyst speeds up both the forward AND reverse reactions equally. It reduces the time to reach equilibrium but does not change the equilibrium position or the value of . A catalyst increases the rate, not the yield. If the question asks “what increases the yield of NH₃?” — the answer includes high pressure, low temperature, and removing product, but NOT the catalyst.