Equilibrium and Rate in the Haber Process Revision Notes | GCSE AQA Higher Chemistry
450°C is a compromise temperature, while high pressure increases both rate and yield because 4 reactant molecules make 2. Free notes and 13 questions answered.
Updated 19 Mar 2026 · Revise Right Team
Equilibrium and Rate in the Haber Process is part of The Haber process and the use of NPK fertilisers (chemistry only) in GCSE AQA Higher Chemistry.
4 min
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GCSE AQA Higher
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What this lesson covers
450°C is a compromise temperature, while high pressure increases both rate and yield because 4 reactant molecules make 2. Free notes and 13 questions answered.
Explain how reaction conditions affect the rate and equilibrium position in the Haber process
Interpret graphs showing the relationship between temperature, pressure and ammonia yield
Evaluate the commercial compromise between rate and equilibrium position
Apply Le Chatelier's principle to predict how changing conditions affects ammonia production
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Use this lesson to stay inside The Haber process and the use of NPK fertilisers (chemistry only) while connecting the detail back to the wider Topic 10: Using resources topic.
Topic 10: Using resources
The Haber process and the use of NPK fertilisers (chemistry only)
GCSE AQA Higher Chemistry
Revision notes
Read the core explanations from Equilibrium and Rate in the Haber Process before testing yourself from memory.
1. The Industrial Challenge
Making ammonia sounds simple: combine nitrogen and hydrogen. But here's the problem — nitrogen molecules are incredibly stubborn. They're held together by one of the strongest bonds in chemistry.
The Haber process solved this by using high temperature and pressure with an iron catalyst. But there's a catch: the conditions that make the reaction fast aren't the same as the conditions that make lots of product.
2. The Haber Process Equation
Notice the double arrow — this reaction is reversible. Ammonia can break down back into nitrogen and hydrogen. The reaction is also exothermic, releasing 92 kJ for every mole of nitrogen that reacts.
At equilibrium, ammonia forms at exactly the same rate it breaks down. The concentrations stop changing, but the reaction never stops. Think of it like a busy roundabout — cars enter and leave at the same rate, so the number on the roundabout stays constant.
3. Le Chatelier's Principle in Action
Think of equilibrium like a stubborn teenager — whatever you try to do, it does the opposite. If you increase pressure, the equilibrium shifts to reduce pressure. If you heat it up, it shifts to cool down.
For the Haber process, this creates some interesting effects. Look at the equation again: 4 molecules of reactants make 2 molecules of product. Fewer molecules means less pressure.
Worked example: Predicting Equilibrium Shifts
The Haber process is at equilibrium. Predict what happens when we increase the pressure.
Left side: $N_2 + 3H_2 = 4$ molecules
Right side: $2NH_3 = 2$ molecules — The right side has fewer gas molecules
System opposes the pressure increase — Equilibrium shifts to reduce pressure
Shifts right (towards ammonia) — Fewer molecules = lower pressure
More ammonia is produced — Higher pressure favours ammonia formation
Answer: Equilibrium shifts right, producing more ammonia
Temperature is trickier. The forward reaction is exothermic (releases heat). So increasing temperature favours the reverse reaction — the one that absorbs heat to cool things down.
4. The Rate vs Yield Dilemma
Here's where it gets complicated. The conditions that give you the highest yield aren't the same as the conditions that give you the fastest rate. This creates a real headache for chemical engineers.
High pressure increases both rate and yield — that's good news. But temperature is the problem child. High temperature increases rate but decreases yield. Low temperature increases yield but makes the reaction painfully slow.
Industrial chemists can't wait forever for a reaction, even if it gives perfect yield. Time is money. They need a compromise that gives reasonable yield in reasonable time.
5. Reading Equilibrium Graphs
Exam questions love to show you graphs of temperature or pressure against percentage yield. The key is remembering that exothermic reactions hate heat and all reactions love pressure when they make fewer molecules.
For the Haber process graphs, you'll see yield decreasing as temperature increases (because it's exothermic) and yield increasing as pressure increases (because 4 molecules become 2).
6. Commercial Reality
The Haber process doesn't run at the theoretical optimum conditions. Real factories use around 450°C and 200 atmospheres pressure. Why not higher pressure for better yield?
Money. Higher pressure needs thicker pipes, stronger pumps, and more energy. The extra cost outweighs the extra ammonia produced. There's also a clever trick — unreacted nitrogen and hydrogen get recycled back through the reactor.
The iron catalyst is crucial too. It doesn't change the equilibrium position, but it massively increases the rate. Without it, the reaction would be too slow to be commercially viable, even at high temperature.
7. Exam Technique and Common Mistakes
Exam questions on the Haber process often ask you to explain why specific conditions are used. The key is mentioning both rate AND equilibrium effects, plus economic factors.
Watch out for questions asking about graphs. If yield decreases with temperature, the reaction must be exothermic. If yield increases with pressure, there must be fewer molecules on the product side.
Remember that industrial conditions aren't about getting perfect yield — they're about making profit. The Haber process uses moderate conditions because they give the best balance of rate, yield, and cost.
Sources and content standards
Published by Revise Right. Use the official specification reference to check the requirements for your course. Content updated: 2026-03-19.
The forward reaction in the Haber process is exothermic, releasing 92 kJ for every mole of nitrogen that reacts. This is why yield decreases as temperature increases, and the reverse reaction, which breaks ammonia down, absorbs heat.
Why is 450°C used in the Haber process?
450°C is a compromise temperature for the Haber process. The forward reaction is exothermic, so a high temperature decreases the yield of ammonia, but a low temperature makes the reaction painfully slow, so around 450°C gives a reasonable yield in reasonable time.
Why is a pressure of 200 atmospheres used in the Haber process?
High pressure increases both rate and yield in the Haber process, because 4 molecules of reactants make 2 molecules of product. Factories use around 200 atmospheres rather than higher because higher pressure needs thicker pipes, stronger pumps and more energy, and the extra cost outweighs the extra ammonia.
What does the iron catalyst do in the Haber process?
The iron catalyst massively increases the rate of the Haber process but does not change the equilibrium position. Without it, the reaction would be too slow to be commercially viable, even at high temperature.
What is Le Chatelier's principle?
Le Chatelier's principle states that when conditions change in an equilibrium system, the position shifts to oppose that change. For example, increasing the pressure in the Haber process shifts the equilibrium right, towards ammonia, because the right side has fewer gas molecules.
What happens to the Haber process equilibrium if pressure is increased?
Increasing pressure shifts the Haber process equilibrium to the right, producing more ammonia. There are 4 gas molecules on the left (N₂ + 3H₂) and 2 on the right (2NH₃), so by Le Chatelier's principle the system shifts towards fewer molecules to reduce the pressure.
How does temperature affect the yield of ammonia?
Increasing temperature decreases the yield of ammonia because the forward reaction is exothermic, so higher temperature favours the reverse reaction, which absorbs heat. However, high temperature increases the rate, while low temperature increases yield but makes the reaction painfully slow.
Why isn't a higher pressure used in the Haber process?
A higher pressure is not used in the Haber process because it costs too much: it needs thicker pipes, stronger pumps and more energy. The extra cost outweighs the extra ammonia produced, so factories use around 200 atmospheres.
How do you read a Haber process yield graph?
Haber process yield graphs show yield decreasing as temperature increases, because the reaction is exothermic, and yield increasing as pressure increases, because 4 molecules become 2. The same logic also works the other way: if yield falls with temperature the reaction must be exothermic.
How can a graph show whether a reaction is exothermic?
A graph shows a reaction is exothermic if the percentage yield decreases as temperature increases. Also, if yield increases with pressure, there must be fewer gas molecules on the product side, as in the Haber process.
Why is nitrogen so unreactive?
Nitrogen is stubborn and hard to react because its molecules are held together by one of the strongest bonds in chemistry. This is why the Haber process needs high temperature and pressure with an iron catalyst to make nitrogen react with hydrogen.
How do you explain the conditions used in the Haber process?
The Haber process conditions are explained by mentioning rate, equilibrium and economic factors together. Around 450°C and 200 atmospheres give the best balance of rate, yield and cost, because industrial conditions are about making profit rather than getting perfect yield.
What is the rate versus yield dilemma in the Haber process?
The rate versus yield dilemma is that the conditions giving the highest yield are not the same as those giving the fastest rate. High pressure increases both, but high temperature increases rate while decreasing yield, so chemists choose a compromise that gives reasonable yield in reasonable time.