Preparation of nitric acid
Pour concentrated sulphuric acid onto potassium nitrate crystals and warm the flask gently, and something odd happens: instead of the usual flask-and-delivery-tube setup you have used for every other gas preparation, the teacher reaches for a strange, curved-necked glass vessel called a retort — no rubber cork in sight. That single change in apparatus is the whole story of why nitric acid is hard to make safely. Nitric acid (HNO₃) is prepared in the school laboratory by heating a nitrate salt with concentrated sulphuric acid, but the product is a volatile, corrosive, light-sensitive liquid that attacks rubber and decomposes if handled carelessly. This article walks through the reaction, explains why every piece of apparatus and every condition — glass instead of rubber, gentle heat instead of strong heat, a cooled receiver — is chosen specifically to protect the acid from decomposing before it can be collected, and works through the reasoning ICSE examiners expect for both the preparation and its common variations.
Heating a nitrate salt with concentrated sulphuric acid displaces the volatile nitric acid, which must then be collected gently, in glass, and kept cool, because HNO₃ decomposes readily under heat or light.
The core idea
Nitric acid is prepared by heating a nitrate salt — usually potassium nitrate (KNO₃) or sodium nitrate (NaNO₃) — with concentrated sulphuric acid in an all-glass retort, applying only gentle heat:
where: KNO₃ is the solid nitrate salt, H₂SO₄ (concentrated) is the displacing acid, KHSO₄ is potassium hydrogensulphate left behind in the retort, and HNO₃ is the nitric acid vapour that distils off and is condensed. Three design choices govern this preparation, and each answers a "why" the syllabus expects a student to explain.
Why concentrated sulphuric acid, and not dilute: the reaction works only because concentrated H₂SO₄ is a strong, non-volatile acid. Being non-volatile, it stays behind in the retort while the more volatile nitric acid is driven off as vapour — the same displacement logic used when concentrated H₂SO₄ liberates HCl gas from a chloride salt. Dilute sulphuric acid, being mostly water, cannot displace HNO₃ this way.
Why an all-glass retort, not a flask with rubber tubing: nitric acid vapour is highly corrosive and attacks rubber and cork almost immediately. A retort is a single curved glass vessel whose long, downward-sloping neck itself carries the vapour to a receiver, so no rubber joint is ever needed. This is the detail examiners probe most often, because it is the one clear difference from every other gas-preparation setup a student has seen (ammonia, hydrogen chloride, and so on all tolerate rubber delivery tubes).
Why gentle heating and a cooled receiver: nitric acid is thermally and photochemically unstable — strong heat, or even light, decomposes it as fast as it forms:
where: the four decomposition products are brown nitrogen dioxide gas, oxygen gas, and water; the reaction is favoured by heat or light and reversed nowhere under lab conditions, so once decomposition starts, product is permanently lost. Gentle heating limits how much HNO₃ decomposes in the retort, and cooling the receiver (commonly in an ice bath) condenses the vapour quickly into liquid, removing it from further heat exposure and minimising decomposition after collection.
Why it works
The choice of every apparatus detail traces back to one fact: nitric acid is both volatile and unstable, so the preparation is a race to displace it, condense it, and cool it before it decomposes or escapes. Displacement needs a non-volatile, strong acid — concentrated H₂SO₄. Collection needs a corrosion-resistant path with no rubber — the all-glass retort. Preservation of the yield needs low heat and a cold receiver, because both heat and light push the reaction forward. The pale yellow tinge often seen in freshly prepared nitric acid is simply dissolved NO₂ from a small amount of this decomposition — evidence of the same instability the apparatus is designed to fight, not a sign of failure. This is also why concentrated nitric acid, once made, is stored in dark (amber) bottles: the storage condition is a direct extension of the preparation's core reasoning.
Worked example
Problem: A student sets up the laboratory preparation of nitric acid but, to save time, heats the retort strongly and leaves the collected acid in an open, uncovered receiver flask in bright sunlight for twenty minutes. (a) Write the equation for the intended preparation. (b) Explain, with an equation, what happens to the collected acid, and describe its likely appearance afterwards.
Solution:
Step 1. The intended preparation reaction, from a nitrate salt and concentrated sulphuric acid, is
Step 2. Strong heating and prolonged sunlight both supply exactly the conditions that trigger nitric acid's decomposition, since HNO₃ is both thermally and photochemically unstable. The relevant equation is
Step 3. Some fraction of the collected HNO₃ decomposes into nitrogen dioxide gas, oxygen gas, and water. The NO₂ produced is brown and partly redissolves in the remaining liquid acid rather than escaping completely, so the acid in the flask darkens.
The collected acid will appear pale yellow to brown instead of colourless, and its effective yield of usable HNO₃ is reduced — exactly the outcome the gentle-heating and cooled-receiver instructions are meant to prevent.
Common traps
- Using dilute sulphuric acid instead of concentrated — only concentrated H₂SO₄ is non-volatile and strong enough to displace nitric acid from the nitrate salt.
- Assuming a flask with rubber cork and delivery tube (as used for other gases) works here — nitric acid vapour corrodes rubber, so an all-glass retort with no rubber joints is essential.
- Heating strongly to "speed up" the preparation — this decomposes HNO₃ as fast as it forms, lowering yield and darkening the product.
- Treating a pale yellow tinge in the collected acid as a mistake — small amounts of dissolved NO₂ from minor decomposition are normal and expected, not a sign the preparation failed.
- Confusing this laboratory retort method with the industrial Ostwald process, which uses ammonia and platinum catalysts under completely different conditions.
Ask the tutor: want to compare, side by side, why HCl and HNO₃ preparations both start with concentrated H₂SO₄ yet need completely different apparatus?
Why can't we just use a normal flask and delivery tube like we do for hydrogen chloride or ammonia
Those gases don't attack rubber the way nitric acid vapour does. HCl and NH₃ preparations rely on a rubber cork and delivery tube to guide the gas into a collection vessel, and rubber tolerates them fine. Nitric acid vapour is far more corrosive; it degrades rubber and cork almost on contact, which would contaminate the product and eventually cause leaks. The retort's own curved glass neck replaces the delivery tube entirely, so there is nothing rubber for the vapour to touch.
Is the pale yellow colour of the nitric acid we collect in the lab actually a problem
Not really — it's expected. Even with gentle heating, a small amount of HNO₃ always decomposes into NO₂, which dissolves back into the liquid acid and tints it pale yellow to light brown. Pure HNO₃ is colourless, but lab-prepared acid is rarely perfectly pure. As long as the heating was gentle and the receiver was kept cool, a faint yellow tinge is normal and doesn't indicate the preparation went wrong; it just reflects a small, unavoidable amount of decomposition.
How is this different from how nitric acid is manufactured industrially
Completely different route. The lab method here starts from a solid nitrate salt and concentrated sulphuric acid, heated gently in a retort — a small-batch, one-step displacement reaction. Industrial manufacture uses the Ostwald process, which starts from ammonia gas, oxidises it catalytically over hot platinum, and proceeds through several gas-phase steps to make nitric acid on a large, continuous scale. The two share no reagents or apparatus and are examined as separate topics.