The Haber Process and NPK Fertilisers

(Separate Chemistry only) The whole of 4.10.4 (the Haber process and NPK fertilisers) is assessed on AQA GCSE Chemistry, not on Combined Science.

Crops take nitrogen out of the soil to grow, and it needs replacing. The Haber process is the industrial method that manufactures ammonia ($N{H}_3$), which is then used to make nitrogen-based fertilisers. Without it, world food production could not feed the current population.

Ammonia is built from two raw materials:

• Nitrogen, obtained from the air (air is about 78% nitrogen).
• Hydrogen, obtained from natural gas (mainly methane).

The purified nitrogen and hydrogen are combined in a reversible reaction:

$N_2(g)+3H_2(g)\rightleftharpoons 2N{H}_3(g)$

The $\rightleftharpoons$ symbol means the reaction is reversible: ammonia can break back down into nitrogen and hydrogen at the same time as it is being formed. This is the central fact that shapes the whole process.

Nitrogen comes from the air; hydrogen comes from natural gas. The reaction is reversible, so not all the reactants turn into ammonia in one pass.

The purified nitrogen and hydrogen are passed over a catalyst under set conditions. You must recall these three conditions.

Condition	Value
Temperature	about 450°C
Pressure	about 200 atmospheres
Catalyst	iron

The steps inside the plant are:

1. Nitrogen (from the air) and hydrogen (from natural gas) are purified and mixed.
2. The gases are passed over an iron catalyst at about 450°C and about 200 atmospheres, where some react to form ammonia.
3. The mixture leaving the reactor contains ammonia together with unreacted nitrogen and hydrogen. It is cooled so that the ammonia liquefies (ammonia has a higher boiling point than nitrogen or hydrogen) and is removed as a liquid.
4. The unreacted nitrogen and hydrogen are recycled back into the reactor so they are not wasted.

The catalyst does not increase the amount of ammonia made. It speeds up the reaction so equilibrium is reached faster, letting the plant produce ammonia more quickly.

Cooling to remove ammonia also removes it from the reaction mixture, which keeps the process producing more, and recycling the leftover gases makes the process efficient.

(Higher Tier only) The rest of this lesson, from here to the end of the trade-off discussion, applies the ideas of dynamic equilibrium and Le Chatelier's principle. These are assessed at Higher tier only.

In a closed system, a reversible reaction reaches dynamic equilibrium: the forward reaction ($N_2+3H_2\to 2N{H}_3$) and the reverse reaction ($2N{H}_3\to N_2+3H_2$) happen at the same rate, so the amounts of each substance stay constant. Equilibrium does not mean the reaction has stopped; both reactions continue, just at equal rates.

Le Chatelier's principle predicts the direction of change: if a condition of a system at equilibrium is changed, the position of equilibrium shifts to oppose that change. Two facts about the Haber reaction let us apply it:

• The forward reaction is exothermic (it releases energy).
• The forward reaction goes from 4 molecules of gas ($1N_2+3H_2$) to 2 molecules of gas ($2N{H}_3$), so it produces fewer gas molecules.

These two facts determine how temperature and pressure affect the yield of ammonia, which is the key to explaining why the industrial conditions are chosen as they are.

(Higher Tier only) Applying Le Chatelier and rate to justify the temperature.

The forward reaction is exothermic, so temperature pulls yield and rate in opposite directions.

Effect on yield. Raising the temperature shifts the equilibrium in the endothermic (reverse) direction to oppose the added heat. That means a higher temperature gives a lower yield of ammonia. On yield alone, a low temperature would be best.

Effect on rate. Raising the temperature makes particles collide more often and with more energy, so the reaction is faster and equilibrium is reached sooner. On rate alone, a high temperature would be best.

These two effects conflict. A very low temperature would give a high yield but so slowly it is uneconomic; a very high temperature would be fast but give little ammonia. The chosen value of about 450°C is a compromise: a moderate yield achieved at a fast enough rate.

A common exam answer: 450°C is a compromise. A lower temperature would increase the yield (the forward reaction is exothermic) but the rate would be too slow, so a moderate temperature balances a reasonable yield against a reasonable rate.

The iron catalyst helps here too. It speeds up the reaction, so a good rate is possible without needing an even higher temperature that would lower the yield further.

(Higher Tier only) Applying Le Chatelier to pressure.

Pressure affects the yield because the forward reaction reduces the number of gas molecules (4 → 2).

Effect on yield. Increasing the pressure shifts the equilibrium towards the side with fewer gas molecules, which is the ammonia side. So a higher pressure gives a higher yield of ammonia. It also increases the rate, because the particles are closer together and collide more often.

On yield and rate alone, the highest possible pressure would be best. So why stop at about 200 atmospheres?

The limit is cost and safety. Very high pressures need extremely thick, strong pipes and vessels and use a lot of energy to compress the gases, which is expensive and dangerous. About 200 atmospheres is chosen as a compromise between a good yield and the cost and safety of building and running the equipment.

Change	Effect on yield of ammonia	Why the compromise value is used
Higher temperature	Lower yield (forward is exothermic)	450°C balances lower yield against a fast rate
Higher pressure	Higher yield (fewer gas molecules)	200 atm balances high yield against cost/safety

Plants need compounds of three elements to grow well: nitrogen (N), phosphorus (P) and potassium (K). NPK fertilisers contain all three, and they are formulations: mixtures of salts made in measured proportions so that the crop gets the right balance of each element.

The salts are made industrially by several processes:

• Ammonia (from the Haber process) is used to make ammonium salts (such as ammonium nitrate and ammonium sulfate) and to make nitric acid.
• Potassium chloride and potassium sulfate are obtained by mining and can be used directly.
• Phosphate rock is mined, but it is insoluble, so plants cannot absorb it. It is treated with an acid to make soluble salts. Which salts form depends on the acid used:

Phosphate rock treated with	Products
Nitric acid	Phosphoric acid + calcium nitrate
Sulfuric acid	Single superphosphate (calcium phosphate + calcium sulfate)
Phosphoric acid	Triple superphosphate (calcium phosphate)

Phosphate rock cannot be used directly because it is insoluble. Treating it with an acid produces soluble phosphate salts that plant roots can take up.

The same compounds, such as an ammonium salt, can be made in a school laboratory or in an industrial plant, and the exam may ask you to compare the two given information.

In the lab, ammonium sulfate is made by a titration: an acid is neutralised with ammonia solution in small batches, one at a time, using simple glassware. In industry the same salt is made on a huge continuous scale using several integrated processes, with the ammonia coming directly from the Haber process.

Feature	Laboratory preparation	Industrial production
Scale	Small (grams), batch by batch	Very large, continuous
Apparatus	Simple glassware, one reaction	Large integrated plant, many processes
Ammonia source	Bought-in ammonia solution	Made on site by the Haber process
Purpose	Making a pure sample	Mass production at low cost per tonne

The lab method is fine for a pure sample but far too slow and small for the millions of tonnes of fertiliser farming needs, which is why the industrial route is a continuous, integrated process.

1. Saying the reaction goes to completion

The Haber reaction is reversible and reaches equilibrium, so not all the nitrogen and hydrogen turn into ammonia in one pass. That is exactly why the unreacted gases are recycled. Never treat it as a one-way reaction.

2. Claiming the catalyst increases the yield

The iron catalyst speeds up the reaction so equilibrium is reached faster, but it does not change the amount of ammonia made at equilibrium. If a question asks about yield, the catalyst is not the answer.

3. Getting the temperature argument backwards (Higher Tier)

The forward reaction is exothermic, so a higher temperature gives a lower yield of ammonia. Students often say a high temperature increases the yield because it speeds the reaction; it speeds the reaction but reduces the yield, which is the whole point of the compromise.

4. Forgetting why pressure is limited (Higher Tier)

Higher pressure does increase the yield, so the reason it is capped at about 200 atmospheres is cost and safety, not chemistry. Very high pressures need expensive, strong equipment and are hazardous.

5. Mixing up the raw material sources

Nitrogen comes from the air, hydrogen from natural gas. Do not write that nitrogen comes from natural gas or that hydrogen comes from the air.

6. Naming the wrong phosphate salt

Match the acid to the product: nitric acid gives calcium nitrate (plus phosphoric acid), sulfuric acid gives single superphosphate, and phosphoric acid gives triple superphosphate. Learn the three pairings rather than guessing.