About four-fifths of the air you breathe is nitrogen, yet most plants cannot use it directly and farmers cannot simply spray it on a field. The N≡N triple bond in a nitrogen molecule is exceptionally strong, which makes N₂ relatively unreactive under ordinary conditions. The Haber process, which won Fritz Haber the 1918 Nobel Prize in Chemistry[3], forces nitrogen to react with hydrogen to make ammonia, the starting point for most nitrogen fertiliser. Making it work economically depends on one big idea: equilibrium.
The chemistry chain
1. Reaction
N₂ + 3H₂ ⇌ 2NH₃: reversible and exothermic in the forward direction.
2. Equilibrium
Forward and reverse rates match; Le Chatelier predicts the shift.
3. Yield vs rate
Position of equilibrium decides yield; kinetics decides rate.
4. Compromise
Industry balances yield, speed and cost.
Some reactions do not run to completion. They can go both ways, and after a while they settle into a state where nothing seems to change. Understanding that state, and how to push it around, is one of the most useful ideas in chemistry, and the Haber process is where it earns its keep.
Reactions that run both ways
In a reversible reaction, products can turn back into reactants. In a closed container, the forward reaction slows as reactants are used up while the reverse reaction speeds up as products build. Eventually the two rates match. That is dynamic equilibrium: the concentrations stay constant, but the molecules are still reacting in both directions.
Both curves flatten once the forward and reverse rates match.
Concentration against time for N₂ + 3H₂ ⇌ 2NH₃ (schematic). Press the button to disturb the system and watch it re-settle.
Le Chatelier's principle
If you disturb a system at equilibrium, the position shifts to oppose the disturbance. It is a prediction tool, not a mystery.
- Add a reactant and the system uses some of it up, shifting towards products. Remove a product and it makes more.
- Increase pressure and, for reacting gases, the system shifts towards the side with fewer gas molecules.
- Increase temperature and the system shifts in the endothermic direction, absorbing the extra heat. Lowering the temperature favours the exothermic direction.
- Add a catalyst and the position does not change at all; only the time taken to reach equilibrium does.
Exam tipIn the exam
A full-marks Le Chatelier answer has three parts: (1) state the change, (2) say which way the position of equilibrium shifts (left or right), and (3) justify it with the principle — the system shifts to oppose the change. Never say a catalyst "shifts the equilibrium", and don't confuse rate with yield. Mark-scheme wording is similar across AQA, OCR and Edexcel, but always check your own specification and past-paper mark schemes.
Quick check
The pressure on the Haber equilibrium is increased. What happens to the position of equilibrium?
Try it yourself
Before the worked example, play with the Haber process. Flip pressure, temperature and the catalyst, and watch which way equilibrium moves and how fast it gets there.
N₂ + 3H₂ ⇌ 2NH₃ (exothermic forward)
Pressure
Temperature
Iron catalyst
Illustrative only: the bars show direction, not real Haber-process yields or rates.
- High pressure favours the side with fewer gas molecules (2 NH₃ instead of 4), so equilibrium shifts towards ammonia.
- High temperature favours the endothermic (reverse) direction, so less ammonia at equilibrium — but the reaction is faster.
- The iron catalyst speeds up both directions equally: rate rises, the equilibrium position does not move.
Conceptual model, not a quantitative simulation. It shows the direction of each shift for N₂ + 3H₂ ⇌ 2NH₃, not real industrial yields.
Worked example: making ammonia
The Haber process combines nitrogen and hydrogen: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ mol⁻¹ (the negative sign means the forward reaction is exothermic). Ammonia is the starting point for most nitrogen fertiliser, and it is widely credited with supporting a large share of the world's food supply.
Count the gas molecules: four on the left (one N₂ and three H₂), two on the right. High pressure therefore shifts equilibrium towards ammonia. The forward reaction is exothermic, so a low temperature would also favour ammonia, but it would make the reaction painfully slow. Industrial plants compromise: typically a high pressure of around 200 atmospheres, a moderate temperature of around 400–450 °C, and an iron catalyst to speed things up (exact figures vary between plants and textbooks)[2]. Removing the ammonia as a liquid keeps pulling the equilibrium to the right[1], and unreacted nitrogen and hydrogen are recycled, so little is wasted.
GCSE coreYield is not the same as rate
Yield is decided by the position of equilibrium, which Le Chatelier's principle predicts. Rate is decided by kinetics: a higher temperature makes particles move faster and collide more often and with more energy, so more collisions have at least the activation energy. A catalyst gives the reaction a lower-energy route. Pressure and temperature affect both, so an exam answer should say which one it is talking about.
| Condition | Effect on yield | Effect on rate | What industry does |
|---|---|---|---|
| High pressure | Increases (4 → 2 gas molecules) | Increases | About 200 atm — higher costs more and is harder to make safe |
| Low temperature | Increases (forward is exothermic) | Decreases | Compromise around 400–450 °C |
| Iron catalyst | No change | Increases | Always used |
| Removing NH₃ | Increases | — | Ammonia is liquefied and removed; unreacted N₂ and H₂ are recycled |
A-Level extensionBeyond the GCSE rule
The pressure rule above assumes pressure is changed by changing the volume. If you raise the total pressure by adding an unreactive gas such as argon at constant volume, the partial pressures of the reacting gases do not change, so the position of equilibrium does not shift. Also, Le Chatelier's principle predicts the direction of a shift, while the equilibrium constant Kc (which depends only on temperature) tells you by how much.
Challenge (A-Level)
At constant volume, argon (an unreactive gas) is added to the Haber equilibrium mixture. What happens to the position of equilibrium?
Quick check
Why isn't the Haber process run at the lowest possible temperature?
Exam-style question
Explain why industry does not use the lowest possible temperature in the Haber process. [3 marks]
Show model answer and mark breakdown
- A lower temperature shifts the equilibrium towards ammonia, because the forward reaction is exothermic, giving a higher yield (1 mark)
- but at a low temperature the rate is too slow, because fewer collisions have at least the activation energy (1 mark)
- so a compromise temperature (about 400–450 °C) with an iron catalyst gives an acceptable yield in an acceptable time (1 mark)
Examiner focus: give a cause, the chemical principle, and the consequence for both yield and rate.
Common mistake: saying the catalyst 'stops working' at low temperature, or saying equilibrium 'takes longer' without mentioning yield. The mark scheme wants yield and rate kept separate.
That is the real lesson: chemistry in industry is a negotiation between yield, rate and cost, and equilibrium reasoning is how you make the argument.
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Sources and further reading
- University of Bristol (Paul May): The Haber Process — high pressure shifts the equilibrium right; removing liquid ammonia; iron catalyst
- LibreTexts Chemistry: The Haber Process — reaction, ΔH, typical industrial conditions
- NobelPrize.org: The Nobel Prize in Chemistry 1918 (Fritz Haber) — prize for the synthesis of ammonia from its elements
- Study Mind: Conditions of the Haber Process (GCSE Chemistry) — revision-level summary of the compromise
Last reviewed: 26 September 2026. Figures are typical published values and rules can change. Exam wording varies by specification, so check yours.
Frequently asked questions
What does dynamic equilibrium actually mean?
In a closed system, a reversible reaction reaches a point where the forward and reverse reactions happen at the same rate, so the amounts of reactants and products stop changing — even though both reactions are still going on all the time.
What does Le Chatelier's principle say?
If a system at equilibrium is disturbed by a change in concentration, pressure or temperature, the position of equilibrium shifts in the direction that opposes the change.
Does a catalyst change the position of equilibrium?
No. A catalyst speeds up the forward and reverse reactions equally, so equilibrium is reached faster, but the final proportions of reactants and products do not change.
Why not run the Haber process at the lowest possible temperature?
The forward reaction is exothermic, so lower temperature would favour more ammonia at equilibrium — but the reaction rate would become far too slow to be useful. Industry uses a compromise temperature that balances yield against speed.
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