Using resources: the Haber process and NPK fertilisers

What you'll learn

This topic explores how chemistry meets industrial-scale food production. You'll study the Haber process for manufacturing ammonia, including how chemists balance reaction conditions to maximise yield and profit. You'll also learn about NPK fertilisers, their formulations, and laboratory preparation methods that mirror industrial techniques.

Key terms and definitions

Haber process — the industrial manufacture of ammonia from nitrogen and hydrogen using an iron catalyst at high temperature and pressure

NPK fertiliser — a formulated fertiliser containing compounds of nitrogen (N), phosphorus (P) and potassium (K), the three essential elements for healthy plant growth

Reversible reaction — a chemical reaction that can proceed in both forward and backward directions, reaching a state of dynamic equilibrium

Compromise conditions — reaction conditions chosen to balance yield, rate and cost in industrial processes

Formulation — a mixture designed as a useful product, with each component present in a measured quantity and contributing to its properties

Eutrophication — excessive enrichment of water bodies with nutrients (often from fertiliser run-off), causing algal blooms and oxygen depletion

Precipitation reaction — a reaction in which an insoluble solid forms when two solutions are mixed

Ammonium salts — compounds formed when ammonia reacts with acids, commonly used as nitrogen sources in fertilisers

Core concepts

The Haber process: industrial ammonia production

The Haber process synthesises ammonia on a massive scale, primarily for fertiliser manufacture. The equation represents a reversible reaction:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) (reaction is exothermic in the forward direction)

Raw materials and their sources:

• Nitrogen — obtained from fractional distillation of liquid air (78% of the atmosphere is nitrogen)
• Hydrogen — extracted from natural gas (methane) or by reacting methane with steam

Industrial conditions:

• Temperature: approximately 450°C
• Pressure: approximately 200 atmospheres (20 MPa)
• Catalyst: iron

The forward reaction is exothermic (releases energy), so according to Le Chatelier's principle, lower temperatures favour ammonia formation. However, low temperatures produce ammonia too slowly for economic viability. The chosen temperature of 450°C represents a compromise between yield and rate.

Higher pressures favour the forward reaction because four gas molecules (1 N₂ + 3 H₂) form two molecules (2 NH₃). Pressures around 200 atmospheres provide reasonable yields without requiring prohibitively expensive high-pressure equipment.

The iron catalyst speeds up both forward and reverse reactions equally, allowing equilibrium to be reached faster without affecting the position of equilibrium or the yield.

Understanding compromise conditions

Industrial chemists don't simply maximise yield—they optimise profitability. This requires balancing multiple factors:

Why not use very high pressure?

• Equipment becomes extremely expensive to build and maintain
• Safety risks increase substantially
• Energy costs escalate
• The yield improvement beyond 200 atmospheres doesn't justify the additional expense

Why not use lower temperature?

• Reaction rate becomes too slow
• Less ammonia produced per day
• Economic losses outweigh the benefit of slightly higher yield

Economic considerations:

• Unreacted nitrogen and hydrogen are recycled through the reactor
• Ammonia is cooled and liquefied for easy separation and removal
• The process runs continuously to maximise efficiency
• Energy from the exothermic reaction is often recovered and reused

Approximately 10-15% of the reactant gases convert to ammonia per pass through the reactor under these conditions. While this seems low, recycling makes the overall process highly efficient.

NPK fertilisers: essential plant nutrients

Plants require three elements in large quantities for healthy growth:

Nitrogen (N) — essential for making proteins and chlorophyll; promotes leafy green growth

Phosphorus (P) — needed for respiration and photosynthesis reactions; supports root development and seed formation

Potassium (K) — activates enzymes and regulates water balance; improves disease resistance and fruit quality

NPK fertilisers are formulations containing compounds that provide all three elements. The NPK ratio (e.g., 15:15:15 or 20:10:10) indicates the percentage composition of each element.

Common compounds in NPK fertilisers

Nitrogen sources:

• Ammonium nitrate (NH₄NO₃) — highly soluble, provides nitrogen in two forms
• Ammonium sulfate ((NH₄)₂SO₄) — also supplies sulfur
• Ammonium phosphate ((NH₄)₃PO₄) — provides both nitrogen and phosphorus
• Urea (CO(NH₂)₂) — organic compound with high nitrogen content

Phosphorus sources:

• Ammonium phosphate ((NH₄)₃PO₄)
• Calcium phosphate (Ca₃(PO₄)₂)
• Single or triple superphosphate (treated phosphate rock)

Potassium sources:

• Potassium chloride (KCl)
• Potassium sulfate (K₂SO₄)
• Potassium nitrate (KNO₃) — also provides nitrogen

Laboratory preparation of fertilisers

You may need to describe how to prepare ammonium salts in the laboratory using neutralisation reactions or precipitation reactions.

Method 1: Neutralisation (for soluble salts)

To prepare ammonium sulfate:

1. Measure dilute sulfuric acid into a beaker using a measuring cylinder
2. Add ammonia solution gradually using a burette or pipette
3. Use an indicator (e.g., methyl orange) or pH meter to detect the end point
4. Alternatively, add ammonia in excess, then add acid drop by drop until neutral
5. Heat the solution gently to evaporate some water
6. Leave to crystallise, then filter and dry the crystals

Equation: 2NH₃(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq)

Method 2: Precipitation (for insoluble salts)

To prepare calcium phosphate:

1. Mix solutions of calcium nitrate and sodium phosphate
2. A white precipitate of calcium phosphate forms immediately
3. Filter to separate the precipitate from the solution
4. Wash the residue with distilled water to remove soluble impurities
5. Dry the product

Equation: 3Ca(NO₃)₂(aq) + 2Na₃PO₄(aq) → Ca₃(PO₄)₂(s) + 6NaNO₃(aq)

Environmental and sustainability issues

Fertiliser overuse consequences:

• Eutrophication — excess fertiliser washes into rivers and lakes, causing algal blooms that block light and deplete oxygen, killing aquatic life
• Groundwater contamination with nitrates poses health risks
• Energy-intensive production contributes to carbon emissions
• Mining phosphate rock depletes non-renewable resources

Sustainable practices:

• Precision farming applies fertiliser only where and when needed
• Controlled-release formulations reduce waste and run-off
• Crop rotation with nitrogen-fixing legumes reduces artificial fertiliser requirements
• Organic fertilisers (manure, compost) recycle nutrients
• Buffer strips of vegetation prevent fertiliser reaching waterways

The Haber process itself requires substantial energy input, mostly from fossil fuels. Research into more sustainable ammonia production methods, including renewable energy sources and improved catalysts, continues worldwide.

Worked examples

Example 1: Explaining the Haber process conditions

Question: The Haber process uses a temperature of 450°C and a pressure of 200 atmospheres. Explain why these conditions are used rather than higher pressure and lower temperature, which would give a better yield of ammonia. [4 marks]

Mark scheme answer:

Higher pressure (1 mark):

• Very high pressure requires very expensive equipment / increases costs significantly
• OR presents greater safety hazards / technical difficulties

Lower temperature (1 mark):

• Lower temperature gives a slower rate of reaction
• OR would produce less ammonia per unit time / uneconomical

Compromise (1 mark):

• The conditions chosen are a compromise between rate and yield
• OR balance economic factors with chemical yield

Additional relevant point (1 mark):

• Energy costs would increase substantially
• OR unreacted gases are recycled, improving overall efficiency

Example 2: Calculating fertiliser composition

Question: A farmer uses an NPK fertiliser with the ratio 20:10:15. Calculate the mass of nitrogen in a 50 kg bag of this fertiliser. [2 marks]

Solution:

The ratio 20:10:15 means 20% nitrogen (1 mark)

Mass of nitrogen = 20/100 × 50 kg = 10 kg (1 mark)

Example 3: Preparing a fertiliser

Question: Describe how you would prepare a pure, dry sample of ammonium nitrate crystals in the laboratory, starting from ammonia solution and dilute nitric acid. [6 marks]

Mark scheme answer:

Method points (any 5 for 5 marks):

• Measure a volume of dilute nitric acid into a beaker
• Add ammonia solution gradually / use a burette or pipette
• Use an indicator (e.g., methyl orange or universal indicator)
• Continue adding until the mixture is neutral / pH 7
• OR add excess ammonia then add acid drop by drop until neutral
• Heat the solution gently to evaporate some water / concentrate the solution
• Leave to crystallise / cool to allow crystals to form
• Filter to separate crystals from solution
• Wash crystals with cold distilled water
• Dry the crystals (e.g., between filter papers or in a warm oven)

Equation (1 mark): NH₃(aq) + HNO₃(aq) → NH₄NO₃(aq)

Common mistakes and how to avoid them

• Confusing "rate" and "yield" — higher temperature increases the rate but decreases the yield for the Haber process. Yield refers to how much product forms at equilibrium; rate refers to how quickly equilibrium is reached.

• Stating the catalyst increases yield — catalysts speed up reactions but don't change the position of equilibrium or the final yield. In the Haber process, iron increases the rate at which equilibrium is achieved.

• Forgetting the Haber process is reversible — always use ⇌ not → in the equation. Ammonia constantly decomposes back to nitrogen and hydrogen even as it forms.

• Misidentifying the limiting factor in plant growth — nitrogen, phosphorus and potassium are all essential; if any one is deficient, it limits growth regardless of how much of the others is present (Liebig's Law of the Minimum).

• Incorrect preparation methods — ammonium salts are soluble, so prepare them by neutralisation followed by crystallisation, NOT by precipitation and filtration.

• Vague explanations about compromise conditions — be specific. Don't just say "compromise between rate and yield"—explain exactly why higher pressure is impractical (cost, safety) and why lower temperature is uneconomical (too slow).

Exam technique for "Using resources: the Haber process and NPK fertilisers"

• "Explain" questions worth 4-6 marks require you to link ideas together. For compromise conditions, connect the chemical principle (Le Chatelier's principle) to practical factors (cost, rate, safety). State the principle, then explain the consequence, then give the industrial consideration.

• Preparation methods — structure your answer chronologically. Use numbered steps or clear sequencing words (first, then, next, finally). Include the key practical details: how you detect neutralisation, how you separate the product, how you ensure purity.

• Mathematical questions on NPK ratios are usually straightforward percentage calculations—show your working clearly for method marks even if your final answer is wrong.

• "Evaluate" or "assess" questions on sustainability require balanced arguments. Discuss both benefits (increased food production) and problems (eutrophication, energy use), then reach a reasoned judgment.

Quick revision summary

The Haber process manufactures ammonia from nitrogen (from air) and hydrogen (from natural gas) using an iron catalyst at 450°C and 200 atmospheres—compromise conditions balancing rate, yield and cost. Ammonia produces nitrogen-based fertilisers. NPK fertilisers are formulations providing nitrogen, phosphorus and potassium for plant growth. Prepare soluble ammonium salts by neutralisation and crystallisation. Excess fertiliser causes environmental damage through eutrophication and groundwater contamination, requiring sustainable farming practices.