Nitric Acid

Nitric acid (HNO3) is a strong acid that produces nitrate salts. It is monoprotic, fully ionised in water, and — unlike hydrochloric and dilute sulfuric acid — an oxidising acid. Its main industrial importance is fertiliser manufacture.

Key facts

Property	Value
Formula	HNO3
Mr	1 + 14 + 3(16) = 63
Acid strength	Strong (fully dissociated: HNO3 -> H+ + NO3-)
Protons donated	1 (monoprotic)
Salt produced	Nitrate (NO3-) — always soluble
Main use	Making ammonium nitrate fertiliser

Concentrated nitric acid is a colourless, fuming liquid; the dilute laboratory acid is a colourless solution. Because it is an oxidising acid it can attack copper (giving brown NO2), a reaction shown but not required for 0620.

Where nitric acid appears in 0620

• Acids, bases and indicators: a typical strong acid, pH 1 when dilute.
• Preparation of salts: nitrate salts are made by titration (Group 1/ammonium) or the excess-solid method — never by precipitation.
• Fertilisers: the acid half of the NH3 + HNO3 fertiliser route.
• Tests for ions, gases and water: dilute HNO3 acidifies the silver nitrate (halide) and barium nitrate (sulfate) tests.

Key reactions

Nitric acid shows every reaction of a typical acid, always forming a nitrate salt.

With metals

Mg(s) + 2HNO3(aq) -> Mg(NO3)2(aq) + H2(g)

Zn(s) + 2HNO3(aq) -> Zn(NO3)2(aq) + H2(g)

Very dilute nitric acid with a reactive metal follows the general acid + metal pattern giving hydrogen. Copper and other unreactive metals do not react with the dilute acid.

With bases and carbonates

CuO(s) + 2HNO3(aq) -> Cu(NO3)2(aq) + H2O(l)

NaOH(aq) + HNO3(aq) -> NaNO3(aq) + H2O(l)

CaCO3(s) + 2HNO3(aq) -> Ca(NO3)2(aq) + H2O(l) + CO2(g)

With ammonia (the fertiliser reaction)

NH3(aq) + HNO3(aq) -> NH4NO3(aq)

This is the key industrial neutralisation. Ammonium nitrate is a nitrogen-rich fertiliser; the ammonia comes from the Haber process.

Testing for the nitrate ion

To confirm NO3-, add aqueous sodium hydroxide, then a little aluminium foil, and warm carefully. The aluminium reduces nitrate ions, releasing ammonia gas that turns damp red litmus paper blue. This distinguishes a nitrate from other soluble salts — see tests for common anions.

All nitrates are soluble

This solubility rule has direct exam consequences:

• No insoluble nitrate exists, so nitrates cannot be made by precipitation.
• Group 1 and ammonium nitrates are prepared by titration (acid + alkali).
• Nitrates of insoluble bases (e.g. copper) are made by adding excess metal oxide or carbonate to the warm acid, then filtering.

A student is given potassium hydroxide solution and dilute nitric acid. Describe how to prepare a pure, dry sample of potassium nitrate crystals. (4 marks)

Mark scheme

• Titrate: add the acid to a measured volume of alkali with an indicator until neutral, and record the exact volume of acid needed [1]
• Repeat using the same volumes of acid and alkali but with no indicator (so the salt is not contaminated) [1]
• Evaporate the solution slowly to the point of crystallisation (until crystals just start to form) [1]
• Leave to cool so crystals form, then filter and dry between filter papers [1]

Examiner note: both reactants are soluble, so precipitation is impossible and a titration is the only route. Candidates lose the final mark by writing “evaporate to dryness” — that decomposes the crystals; you evaporate only to the point of crystallisation.

Common exam mistakes

• Writing hydrogen as a product of copper with nitric acid. Copper is below hydrogen in the reactivity series, so it does not react with the dilute acid at all.
• Giving the nitrate ion the wrong charge. NO3- is 1-, so magnesium nitrate is Mg(NO3)2, not Mg(NO3).
• Using hydrochloric or sulfuric acid to acidify the silver nitrate or barium nitrate tests. These introduce chloride or sulfate ions and give a false positive; only dilute nitric acid is safe.
• Trying to make a nitrate by precipitation. Every nitrate is soluble, so you must use titration or the excess-solid method.