Manufacture of nitric acid
Nitric acid is needed in enormous quantities — for fertilisers, explosives and dyes — far more than any laboratory retort could ever supply. The laboratory method heats a nitrate salt with concentrated sulphuric acid, which works fine for a test-tube of acid but is hopelessly slow and expensive at industrial scale, since nitrate salts are themselves costly to source. Industry instead starts from ammonia, which Haber's Process already supplies cheaply and continuously, and runs it through a three-stage catalytic sequence called the Ostwald Process. This article walks through the three stages — catalytic oxidation of ammonia, oxidation of the nitric oxide formed, and absorption in water — with their equations, the reasoning behind the conditions used (catalyst, temperature, excess air), and a worked stoichiometry calculation linking mass of ammonia feedstock to mass of nitric acid produced.
The Ostwald Process manufactures nitric acid industrially by catalytically oxidising ammonia to nitric oxide over hot platinum–rhodium, oxidising that nitric oxide further to nitrogen dioxide, and absorbing the nitrogen dioxide in water with excess air to give nitric acid.
The core idea
The Ostwald Process converts ammonia into nitric acid in three linked stages, each pushed toward completion by the conditions chosen.
Stage 1 — catalytic oxidation of ammonia. Ammonia gas mixed with excess air is passed over a heated platinum–rhodium gauze catalyst at about 800 degrees Celsius:
where: NH₃ is ammonia gas, O₂ is oxygen from the excess air, and the platinum–rhodium gauze is a catalyst that speeds up the reaction without being consumed; the arrow's Δ marks the high temperature needed.
Stage 2 — oxidation of nitric oxide. The hot gases are cooled, and the nitric oxide (NO) formed reacts with more oxygen from the excess air to give nitrogen dioxide:
where: NO is nitric oxide (colourless), and NO₂ is nitrogen dioxide (a brown gas); cooling favours this step, unlike Stage 1 which needs heat.
Stage 3 — absorption in water. Nitrogen dioxide is absorbed into water in the presence of excess air/oxygen to form nitric acid:
where: HNO₃ is nitric acid; the acid collected this way is dilute, typically around 60% concentration, and can be concentrated by distillation afterwards if needed. Adding all three stages together and cancelling the intermediates (NO and NO₂ appear on both sides) gives one overall equation linking feedstock to product directly:
where: this overall equation is not a separate reaction step — it is bookkeeping obtained by summing Stages 1 to 3 and dividing through — but it is exactly what a stoichiometry calculation needs, since the intermediate NO and NO₂ never have to be tracked explicitly.
Why it works
The whole process is designed around one constraint: nitric acid is needed in bulk, so every stage must be a fast, continuous, driven-to-completion reaction rather than a batch preparation.
Why ammonia, not a nitrate salt. The laboratory route (heating a nitrate salt with concentrated sulphuric acid) is a small-batch reaction limited by how much nitrate salt is available and how fast heat can be supplied. Ammonia, by contrast, is manufactured continuously and cheaply by Haber's Process, and its catalytic oxidation over a fixed platinum–rhodium gauze can run as an unbroken industrial stream — feeding gas in one end and drawing nitric acid out the other, around the clock.
Why excess air at every stage. In Stage 1, excess oxygen ensures nearly all the ammonia is oxidised to nitric oxide rather than left unreacted or oxidised to unwanted products. In Stage 2, without ample oxygen the conversion of NO to NO₂ would stay incomplete and nitric oxide would simply escape the system unconverted. In Stage 3, excess air again drives the absorption reaction so that essentially all the nitrogen dioxide ends up as nitric acid instead of forming a nitric–nitrous acid mixture or venting as unreacted NO₂ gas. At each stage, more oxygen than the bare equation demands means the reaction is pushed further toward completion, and yield goes up.
Why cooling between Stage 1 and Stage 2. The hot gases leaving the catalyst chamber are deliberately cooled before Stage 2, because the oxidation of NO to NO₂ proceeds more readily at lower temperature; cooling also condenses out some of the water vapour carried over from Stage 1, keeping the gas stream cleaner for absorption.
Worked example
Problem: Using the overall equation of the Ostwald process, calculate the mass of nitric acid that can theoretically be obtained from 34 kg of ammonia, assuming complete conversion at every stage. (Molar masses: N = 14, H = 1, O = 16, so NH₃ = 17 g/mol and HNO₃ = 63 g/mol.)
Step 1. Write the overall equation obtained by summing the three Ostwald stages: . This shows 1 mole of ammonia gives 1 mole of nitric acid — a direct 1:1 mole ratio, so the intermediates NO and NO₂ do not need to be tracked separately for this calculation.
Step 2. Convert the given mass of ammonia to moles. , and moles of NH₃ .
Step 3. Apply the 1:1 mole ratio from the overall equation. Moles of HNO₃ produced (same as moles of NH₃, since the ratio is 1:1).
Step 4. Convert moles of HNO₃ back to mass. Mass of HNO₃ . The theoretical yield of nitric acid from 34 kg of ammonia is 126 kg. In an actual plant the real yield would be somewhat lower than this theoretical value, since no industrial process reaches 100% conversion at every stage.
Common traps
- Confusing the order of the intermediates — NO (nitric oxide) forms first in Stage 1, and only then is oxidised to NO₂ (nitrogen dioxide) in Stage 2; do not swap this sequence.
- Naming the catalyst as iron — that is Haber's Process for ammonia synthesis. The Ostwald process uses platinum, usually alloyed with rhodium.
- Explaining the process with equilibrium or Le Chatelier reasoning — the syllabus wants the descriptive steps and conditions only, not equilibrium justification.
- Skipping straight to the overall equation in a mole calculation without first knowing it is derived by summing and cancelling the three stage equations — write it out or state where it comes from if asked to show working.
Ask the tutor: want a side-by-side comparison of the conditions used in Haber's Process versus the Ostwald process, or a quick drill on stoichiometry calculations using the overall Ostwald equation?
Why is ammonia used as the starting material instead of a nitrate salt like the laboratory method uses?
Because scale changes what is practical. The laboratory method (heating a nitrate salt with concentrated sulphuric acid) is a small batch reaction limited by how much nitrate salt you have on hand, and nitrate salts are relatively costly. Ammonia, on the other hand, is already being produced cheaply and continuously by Haber's Process, and its catalytic oxidation can run as an unbroken industrial stream. That continuous, low-cost feedstock is what makes the Ostwald process suitable for manufacturing nitric acid in the very large quantities needed for fertilisers and explosives.
Do I need to memorise the overall equation NH3 + 2O2 gives HNO3 + H2O separately from the three stage equations?
It helps to know where it comes from rather than memorise it as a fourth, unrelated equation. It is simply the three Ostwald stage equations added together, with the intermediates NO and NO₂ cancelling out because they appear on both sides of the combined sum. For descriptive questions, examiners usually want the three individual stage equations in order. The overall equation becomes genuinely useful only when a question asks for a mole or mass calculation linking ammonia feedstock directly to nitric acid product, since it skips the need to track the intermediates.
Why does the gas mixture need to be cooled between the first and second stages?
The first stage runs hot, at about 800 degrees Celsius, because that temperature is needed for the platinum-rhodium catalyst to oxidise ammonia efficiently. But the second stage, oxidation of nitric oxide to nitrogen dioxide, actually proceeds better at a lower temperature. So the hot gas stream leaving the catalyst chamber is cooled before it is led into the second stage; this also has the side benefit of condensing out some of the water vapour formed in Stage 1, leaving a cleaner gas stream for the oxidation and absorption steps that follow.