Fertilisers and the Haber Process GCSE Chemistry

Without artificial fertilisers, the world could not feed its eight billion people. The Haber process — developed by Fritz Haber and Carl Bosch in the early 1900s — is how we make the ammonia that fertilisers are built from. It is one of the most consequential chemical processes ever invented, and it sits at the core of GCSE chemistry.

Why do plants need fertilisers?

Plants need three key mineral nutrients for healthy growth, often abbreviated as NPK:

In natural ecosystems, these minerals are returned to the soil when organisms die and decompose. In agriculture, they are repeatedly removed in harvested crops and must be replaced — either through natural processes (crop rotation with nitrogen-fixing legumes, animal manure) or by applying artificial fertilisers containing soluble forms of N, P, and K.

Common artificial fertilisers include ammonium nitrate (NH₄NO₃, nitrogen fertiliser), ammonium phosphate ((NH₄)₃PO₄), and potassium chloride (KCl). All must be soluble so that plant roots can absorb them from soil water.

What is the Haber process?

The Haber process is the industrial synthesis of ammonia (NH₃) from its elements — nitrogen gas (N₂) from the air and hydrogen gas (H₂) from natural gas (methane).

Overall equation:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol

Note the reversible arrow (⇌) — this is an equilibrium reaction, meaning both the forward reaction (making ammonia) and the reverse reaction (ammonia breaking down) occur simultaneously. Industrially, we never reach 100% conversion.

The reaction is:

• Exothermic (ΔH negative) — it releases heat energy as ammonia is formed.
• A decrease in moles of gas — 4 moles of gas reactants (1 + 3) produce 2 moles of gas product.

These two facts — exothermic and decreasing gas moles — determine why the industrial conditions are what they are.

Why are the industrial conditions chosen?

The Haber process operates at:

• Temperature: approximately 450 °C
• Pressure: approximately 200 atmospheres (atm)
• Catalyst: iron (with potassium hydroxide and aluminium oxide as promoters)

Each condition involves a compromise using Le Chatelier's Principle (the idea that a system at equilibrium will shift to oppose a change):

Temperature — why 450 °C?

The reaction is exothermic. Le Chatelier's principle tells us that lower temperatures favour the forward reaction (the system shifts to produce heat when cooled, producing more NH₃). So for maximum yield, a very low temperature would be ideal.

However: at low temperatures, the rate of reaction is too slow to be economically viable. Even with a catalyst, reactions at 100 °C or 200 °C would produce ammonia so slowly that the process would not be commercially useful.

Compromise: 450 °C gives a reasonable yield (~15%) at a rate fast enough to be profitable. The unreacted gases are recycled, so the low yield per pass does not waste raw materials.

Pressure — why 200 atm?

The equation shows 4 moles of gas reactants and 2 moles of gas products. Le Chatelier's principle tells us that high pressure favours the side with fewer gas molecules — so higher pressure pushes the equilibrium to the right, producing more NH₃.

However: very high pressures are extremely expensive to generate and maintain (requiring heavy-duty compressors and thick-walled vessels) and create serious safety risks.

Compromise: 200 atm gives a good yield whilst remaining economically and safely viable. Some modern plants operate at 100–300 atm depending on their design.

Catalyst — iron with promoters

The iron catalyst does not change the position of equilibrium (it does not change the yield). It speeds up both the forward and reverse reactions equally, allowing equilibrium to be reached more quickly. Without the catalyst, even at 450 °C the reaction would be too slow.

Potassium hydroxide (KOH) and aluminium oxide (Al₂O₃) act as promoters — they keep the iron surface from sintering (clumping together) and maintain its catalytic activity over time.

How is ammonia converted into fertilisers?

Liquid ammonia itself can be injected directly into soil as a fertiliser, but more commonly it is converted into solid, stable compounds:

1. Ammonium nitrate (NH₄NO₃): ammonia reacted with nitric acid (itself made from ammonia via the Ostwald process): NH₃ + HNO₃ → NH₄NO₃. A very effective nitrogen fertiliser because it contains two forms of nitrogen.

2. Ammonium sulphate ((NH₄)₂SO₄): ammonia reacted with sulphuric acid. Also used as a soil acidifier.

3. Urea (CO(NH₂)₂): the most widely used nitrogen fertiliser globally. Made from ammonia and carbon dioxide.

What are the environmental problems with fertiliser use?

Fertilisers are extremely beneficial for food production, but their misuse or overuse causes environmental damage:

Eutrophication:

1. Excess fertiliser (particularly nitrates and phosphates) is washed from fields into rivers and lakes by rain — a process called leaching.
2. The elevated mineral levels cause algae to grow explosively (an algal bloom), covering the water surface.
3. The algal bloom blocks sunlight from reaching underwater plants, which die.
4. Bacteria decompose the dead plant material. The bacteria reproduce rapidly and consume most of the dissolved oxygen in the water.
5. The drop in oxygen levels (deoxygenation) causes fish and other aquatic organisms to suffocate and die.
6. The waterway becomes biologically dead — a process known as eutrophication.

Nitrates from fertilisers can also leach into groundwater and drinking water, where high concentrations (above ~50 mg/L nitrate in the UK) can be harmful, particularly for infants.

Mitigation strategies include buffer strips of grass beside waterways, precision application of fertilisers (only when and where plants need them), and the use of slow-release fertiliser formulations.

How do you answer Haber process questions in the exam?

For a question asking you to justify the choice of conditions, always:

1. State the effect using Le Chatelier's principle.
2. State the opposing practical reason.
3. State the chosen compromise.

Example: "Why is 450 °C used rather than a lower temperature? Lower temperatures would produce a higher yield of ammonia (Le Chatelier's principle — the exothermic reaction is favoured at lower temperatures). However, the rate of reaction would be too slow to be economically viable. 450 °C is a compromise between yield and rate."

Frequently asked questions

Where does the hydrogen come from in the Haber process?

The hydrogen comes from steam reforming of natural gas (methane). Methane reacts with steam over a nickel catalyst at high temperature:

CH₄(g) + H₂O(g) → CO(g) + 3H₂(g)

The carbon monoxide (CO) produced is then reacted with more steam (the water-gas shift reaction) to produce more hydrogen and carbon dioxide, which is removed. This source of hydrogen makes the Haber process dependent on fossil fuels — both as a fuel to heat the reactor and as a raw material. Research into green hydrogen (made by electrolysis of water using renewable electricity) aims to make fertiliser production carbon-neutral in future.

Why is the nitrogen fixed — what does that mean?

The atmosphere is about 78% nitrogen gas (N₂), but most organisms cannot use N₂ directly because the N≡N triple bond is extremely strong (bond energy ~945 kJ/mol) and very difficult to break. "Fixing" nitrogen means converting N₂ into a compound that plants (and other organisms) can use, such as ammonium or nitrate. The Haber process is one form of nitrogen fixation; bacteria in the root nodules of legumes (clover, peas, beans) also fix nitrogen biologically using the enzyme nitrogenase. Before the Haber process, agriculture depended entirely on biological fixation and was fundamentally limited by nitrogen availability.

What happened before the Haber process existed?

Before 1913, the primary source of nitrogen fertiliser was guano (bird and bat droppings) mined from islands off the coast of South America (particularly Peru and Chile), and sodium nitrate from the Atacama Desert. Demand for these natural sources vastly exceeded supply, and famines were predicted as the global population grew. The Haber-Bosch process (industrialised by Carl Bosch at BASF) is estimated to have enabled the agricultural productivity that supports roughly half the current global population — it is sometimes described as the most important chemical process in human history.

Does the Haber process produce pollution?

The process itself does not produce significant pollution at the point of manufacture (nitrogen, hydrogen, and ammonia are not toxic in controlled industrial quantities). The carbon footprint arises from the steam reforming step, which uses and produces carbon dioxide. The environmental impact is primarily from how the fertilisers are used (eutrophication, nitrous oxide emissions from soils), not from the synthesis itself. Modern Haber-Bosch plants are highly energy-efficient compared to earlier designs, recovering heat from the exothermic reaction to pre-heat incoming reactants.

Let Professor Curie at aitutors.me guide you through Le Chatelier's principle in the Haber process, help you structure the equilibrium-conditions question, and quiz you on the eutrophication chain — step by step.