Fertilisers

Fertilisers are chemical substances added to soil to supply essential nutrients that plants require for healthy growth. The three most critical macronutrients are nitrogen (N), phosphorus (P), and potassium (K) — collectively referred to as NPK. Each element plays a distinct biochemical role: nitrogen is a key component of amino acids, proteins, and chlorophyll; phosphorus is essential for ATP synthesis and DNA structure; and potassium regulates osmotic balance and enzyme activation. Without adequate supplies of these elements, crop yields fall dramatically, which is why the industrial manufacture of synthetic fertilisers has become one of the most consequential areas of applied chemistry in the modern world. The global demand for food has driven the development of large-scale chemical processes to produce NPK fertilisers efficiently and economically. Nitrogen-based fertilisers, including ammonium nitrate (NH₄NO₃) and ammonium sulfate ((NH₄)₂SO₄), are synthesised from ammonia produced via the Haber process. Phosphate fertilisers are derived from phosphate rock — primarily fluorapatite, Ca₅(PO₄)₃F — which is treated with strong acids to release plant-available phosphate ions. Potassium fertilisers typically come from mined potash deposits such as potassium chloride (KCl) and require relatively little chemical processing. Understanding the chemical reactions underpinning these manufacturing routes is central to industrial inorganic chemistry. NPK fertilisers are sold in formulations identified by three numbers representing the percentage by mass of N, P₂O₅, and K₂O respectively. For example, a 20-10-10 fertiliser contains 20% nitrogen, 10% phosphorus (expressed as P₂O₅), and 10% potassium (expressed as K₂O) by mass. This standardised labelling system allows farmers to tailor nutrient delivery to specific soil conditions and crop needs. The industrial chemistry behind producing and blending these compounds requires precise stoichiometric control, knowledge of acid-base reactions, and careful management of exothermic processes — all core topics in general inorganic chemistry.

Key takeaways
  • Understanding NPK Formulations
  • Treatment of Phosphate Rock with Sulfuric Acid — Making Superphosphate
  • Treatment of Phosphate Rock with Nitric Acid — The Odda Process
  • Production of Ammonium Nitrate
  • Production of Ammonium Sulfate

Understanding NPK Formulations

NPK formulations are expressed as three numbers: %N : %P₂O₅ : %K₂O by mass. Nitrogen content is reported directly as elemental N. Phosphorus content is expressed as the oxide equivalent P₂O₅ (even though P₂O₅ itself is not present in the fertiliser). Potassium content is expressed as K₂O. To convert: %P = %P₂O₅ × (2×31)/(2×31 + 5×16) = %P₂O₅ × 0.436. Similarly, %K = %K₂O × 0.830. Compound fertilisers blend multiple nutrient sources to achieve desired ratios.

Treatment of Phosphate Rock with Sulfuric Acid — Making Superphosphate

Phosphate rock (fluorapatite, Ca₅(PO₄)₃F) is insoluble in water and not directly available to plants. It is treated with concentrated sulfuric acid (H₂SO₄) to produce single superphosphate (SSP), a mixture of calcium dihydrogen phosphate Ca(H₂PO₄)₂ and calcium sulfate CaSO₄. The reaction is: Ca₅(PO₄)₃F + 5H₂SO₄ → 3H₃PO₄ + 5CaSO₄ + HF. In practice the H₃PO₄ produced reacts with further apatite: Ca₅(PO₄)₃F + 7H₃PO₄ + 5H₂O → 5Ca(H₂PO₄)₂ + HF. Single superphosphate contains approximately 16–21% P₂O₅. Triple superphosphate (TSP, ~46% P₂O₅) is made by treating phosphate rock with phosphoric acid instead, avoiding CaSO₄ as a by-product.

Treatment of Phosphate Rock with Nitric Acid — The Odda Process

In the Odda (or nitrophosphate) process, phosphate rock is dissolved in nitric acid: Ca₅(PO₄)₃F + 10HNO₃ → 5Ca(NO₃)₂ + 3H₃PO₄ + HF. The solution is cooled to crystallise out calcium nitrate tetrahydrate, Ca(NO₃)₂·4H₂O, which is separated and later converted to ammonium nitrate by reaction with ammonia: Ca(NO₃)₂ + 2NH₃ + CO₂ + H₂O → 2NH₄NO₃ + CaCO₃. The remaining phosphoric acid solution is neutralised with ammonia to give ammonium phosphates. This process elegantly produces both N and P fertiliser components from a single starting acid treatment.

Production of Ammonium Nitrate

Ammonium nitrate (NH₄NO₃, 34% N) is the world's most widely used nitrogen fertiliser. It is manufactured by neutralising nitric acid with ammonia gas in a highly exothermic acid-base reaction. The reaction is carried out in a neutralisation tower where liquid ammonia or ammonia gas is added to concentrated nitric acid (55–65%). The product solution (~80–95% NH₄NO₃) is concentrated by evaporation, then prilled or granulated for safe handling. Safety is paramount: NH₄NO₃ is a powerful oxidiser and can detonate under confinement and heat, so industrial production requires strict safety controls.

Production of Ammonium Sulfate

Ammonium sulfate ((NH₄)₂SO₄, 21% N) is produced by reacting ammonia with sulfuric acid. It is also recovered as a by-product from the coking industry (where ammonia scrubbing with H₂SO₄ captures NH₃ from coal gas). Although lower in nitrogen content than ammonium nitrate, it provides sulfur (24% S), making it valuable for sulfur-deficient soils. The product is less hazardous than NH₄NO₃ and is slightly acidifying, which can be beneficial for alkaline soils.

The Key Equation

\ceCa5(PO4)3F+5H2SO4−>3H3PO4+5CaSO4+HF\ce{Ca5(PO4)3F + 5H2SO4 -> 3H3PO4 + 5CaSO4 + HF}
  • \ce{Ca5(PO4)3F} — Fluorapatite (mol): The primary mineral in phosphate rock; insoluble in water and not bioavailable without chemical treatment.
  • \ce{H2SO4} — Sulfuric acid (mol): Concentrated sulfuric acid used as the acidifying agent to break down the apatite lattice.
  • \ce{H3PO4} — Phosphoric acid (mol): Intermediate product that reacts further with apatite to give plant-available dihydrogen phosphate ions.
  • \ce{CaSO4} — Calcium sulfate (gypsum) (mol): By-product of single superphosphate production; present in SSP and reduces the overall P₂O₅ content.
  • \ce{HF} — Hydrogen fluoride (mol): Toxic gas by-product that must be scrubbed from plant emissions to prevent environmental and health hazards.

The Key Equation

\ceNH3(g)+HNO3(aq)−>NH4NO3(aq)ΔH=−144 kJ mol−1\ce{NH3(g) + HNO3(aq) -> NH4NO3(aq)} \quad \Delta H = -144\,\text{kJ mol}^{-1}
  • \ce{NH3} — Ammonia (mol): Produced industrially by the Haber process from N₂ and H₂; acts as the base in this neutralisation.
  • \ce{HNO3} — Nitric acid (mol): Produced by the Ostwald process from ammonia; acts as the acid in the neutralisation.
  • \ce{NH4NO3} — Ammonium nitrate (mol): Contains 34% N by mass; both ammonium (NH₄⁺) and nitrate (NO₃⁻) provide plant-available nitrogen.
  • \Delta H — Enthalpy change (kJ mol⁻¹): The strongly exothermic nature of this reaction requires careful temperature control in industrial reactors.

Worked Example

Problem. A compound fertiliser is labelled 15-15-15 (NPK by mass). Calculate the minimum mass of ammonium nitrate (NH₄NO₃, M = 80.0 g/mol) needed to provide the nitrogen in 500 kg of this fertiliser. Assume all nitrogen comes from NH₄NO₃.

m(\ceNH4NO3)=75,000 g0.350≈214,000 g=214 kgm(\ce{NH4NO3}) = \frac{75{,}000\,\text{g}}{0.350} \approx 214{,}000\,\text{g} = 214\,\text{kg}

Answer. Approximately 214 kg of ammonium nitrate is needed to supply the nitrogen content of 500 kg of 15-15-15 fertiliser.

Figure
Diagram
Industrial Fertiliser Production: OverviewHaber ProcessN₂ + 3H₂ → 2NH₃Ostwald ProcessNH₃ → HNO₃Phosphate RockCa₅(PO₄)₃FAcid PlantH₂SO₄ / HNO₃NeutralisationNH₃ + HNO₃Neutralisation2NH₃ + H₂SO₄Acid Treatment+ H₂SO₄ or HNO₃NH₄NO₃34% N fertiliser(NH₄)₂SO₄21% N fertiliserSuperphosphateCa(H₂PO₄)₂ + CaSO₄NPK Blende.g. 20-10-10Raw Materials & ProcessesReaction StepsProducts
Origins · 1 minFritz Haber and how this was discovered
FH
Fritz Haber
1868–1934 · German

Fritz Haber developed the Haber process (1909) for synthesising ammonia from nitrogen and hydrogen gases under high pressure (150–300 atm) and temperature (400–500°C) using an iron catalyst: N₂ + 3H₂ ⇌ 2NH₃. This reaction is the foundational step for producing virtually all synthetic nitrogen fertilisers, enabling the dramatic expansion of global food production throughout the 20th century. It is estimated that Haber's process feeds approximately half the world's population today. He was awarded the Nobel Prize in Chemistry in 1918. Haber's legacy is deeply ambivalent: the same chemical mastery that enabled mass food production also led him to develop chlorine gas weapons used in World War I — making him one of history's most morally complex scientific figures.

  • The Haber Process — The Haber process produces the ammonia that is the essential feedstock for manufacturing ammonium nitrate and ammonium sulfate fertilisers.
  • The Ostwald Process — The Ostwald process converts ammonia to nitric acid, which is directly used to produce ammonium nitrate and to dissolve phosphate rock in the Odda process.
  • Acid-Base Neutralisation Reactions — Fertiliser production relies heavily on neutralisation reactions — ammonia (base) reacting with nitric acid or sulfuric acid (acids) to form ammonium salts.
  • Stoichiometry and Percentage Yield — Calculating the mass of fertiliser produced from given quantities of reactants, or finding percentage composition, requires precise stoichiometric reasoning.
  • Eutrophication and Environmental Impact of Fertilisers — Excess NPK fertiliser runoff into waterways causes eutrophication — algal blooms and oxygen depletion — a major environmental consequence of industrial fertiliser use.