Le Chatelier's Principle and the Haber Process (HT)AQA GCSE Chemistry: Revision notes
Section 1
What Does It Mean for a Reaction to Be Reversible?
A reversible reaction is one whose products can react to re-form the original reactants, shown using the symbol ⇌. The direction favoured by the reaction can be changed by changing conditions such as temperature, pressure or concentration.
When a reversible reaction happens in a closed system, it can reach dynamic equilibrium: the forward and reverse reactions occur at exactly the same rate, so the overall amounts of reactants and products stay constant, even though both reactions are still taking place.
Section 2
How Does Le Chatelier's Principle Predict Equilibrium Shifts?
Le Chatelier's Principle states that if a system at equilibrium experiences a change in conditions, the system will respond in a way that partially counteracts that change.
This lets us make qualitative predictions — not exact numerical ones — about how changing concentration, temperature or pressure will shift the position of equilibrium.
Section 3
What Happens to Equilibrium When We Change Concentration, Temperature or Pressure?
| Change | Effect on equilibrium position |
|---|---|
| Increase concentration of a reactant | Shifts to favour the forward reaction, using up the extra reactant |
| Increase concentration of a product | Shifts to favour the reverse reaction |
| Increase temperature | Shifts to favour the endothermic direction (absorbs the extra heat) |
| Decrease temperature | Shifts to favour the exothermic direction |
| Increase pressure (gases only) | Shifts to favour the side with fewer gas molecules |
| Decrease pressure (gases only) | Shifts to favour the side with more gas molecules |
These predictions apply specifically to reactions at equilibrium, and pressure changes only affect equilibria involving gases.
For pressure, count the number of gas molecules on each side of the equation — the side with fewer molecules is favoured by an increase in pressure, regardless of which side is 'reactants' or 'products'.
Section 4
How Does This Apply to the Haber Process?
The Haber process manufactures ammonia by the reversible reaction:
nitrogen + hydrogen ⇌ ammonia (exothermic in the forward direction)
The reaction has 4 moles of gas on the reactant side (1 N₂ + 3 H₂) and 2 moles of gas on the product side (2 NH₃), so:
- High pressure (~200 atmospheres) favours the forward reaction (fewer gas molecules), giving a higher yield of ammonia
- Because the forward reaction is exothermic, low temperature would favour the forward reaction and give a higher yield — but low temperature also makes the reaction too slow to be commercially useful
- In practice, a compromise temperature of about 450 °C is used, together with an iron catalyst, to give a reasonably fast rate at an acceptable yield
- This shows the trade-off between the position of equilibrium (favouring high yield) and the rate of reaction (favouring a viable production speed) that industrial chemists must balance, alongside the cost of raw materials and energy
Although lower temperature would shift equilibrium further towards ammonia, the Haber process uses 450 °C rather than a much lower temperature, because reaction rate would otherwise be far too slow for industrial production.
Must Know
- Le Chatelier's Principle: a system at equilibrium shifts to counteract any imposed change
- Increasing reactant concentration shifts equilibrium towards products; increasing product concentration shifts it towards reactants
- Increasing temperature favours the endothermic direction; decreasing temperature favours the exothermic direction
- Increasing pressure favours the side of a gaseous equilibrium with fewer gas molecules
- The Haber process (N₂ + 3H₂ ⇌ 2NH₃, exothermic forward) uses high pressure (~200 atm) to favour ammonia
- A compromise temperature (~450 °C) with an iron catalyst balances yield against an acceptable reaction rate
That's the notes covered.
Carry on to the next subtopic.