Properties of nitric acid

Physical properties. Pure nitric acid (HNO₃) is a colourless, highly corrosive, fuming liquid with a sharp, choking smell. Bottles of it in the lab often look pale yellow rather than colourless — that tinge comes from a little of the acid slowly decomposing in light or on warming into dissolved nitrogen dioxide (NO₂). It is miscible with water in all proportions, so dilute and concentrated nitric acid are simply the same acid at different water content.

The oxidising personality. The single fact that organises everything else here is that nitric acid is a strong oxidising acid. Dilute HCl and dilute H₂SO₄ react with a reactive metal by simple displacement: the metal gives up electrons to H⁺ ions, and hydrogen gas bubbles off. Nitric acid, whether dilute or concentrated, reacts with most metals without liberating hydrogen (very dilute acid with magnesium or manganese is sometimes quoted as a rare textbook exception). Instead of H⁺ being reduced to H₂, the nitrate ion itself is reduced, and the exact mix of nitrogen-oxide products that forms depends on the metal's reactivity and the acid's strength — detail that sits beyond this syllabus. What you must carry forward is simpler and sturdier: no H₂ evolution with typical reactive metals is the signature of an oxidising acid. The same oxidising streak reaches beyond metals. On heating, nitric acid oxidises certain non-metals — carbon and sulphur, for instance, are converted to their oxides while the acid is reduced, releasing brown fumes of NO₂. You are not expected to balance these non-metal reactions in detail, only to recognise the pattern: nitric acid oxidises what it touches, and brown NO₂ fumes are the visible fingerprint of that reduction.

Action of heat on metal nitrates. Heat is where the oxidising nature becomes a predictable rulebook. Every metal nitrate breaks down on strong heating, but what it breaks down into depends on where the metal sits in the activity series — three tiers, from most reactive to least: Most reactive metals (potassium, sodium) give the metal nitrite plus oxygen:

where: $\Delta$ over the arrow means "heat strongly," and the upward arrow marks a gas escaping the test tube. Moderately reactive metals (lead, copper, zinc, calcium, magnesium) give the metal oxide plus nitrogen dioxide plus oxygen:

where: notice this tier goes one step further than the first — the nitrite is not stable enough to survive, so the metal ends up as an oxide, and the nitrogen escapes as NO₂ rather than staying combined. Least reactive metals (silver, mercury) decompose all the way to the free metal plus nitrogen dioxide plus oxygen:

where: here even the oxide is too weakly held together, so heating strips the metal all the way down to its element — the least reactive metals give up their oxygen bonds most easily of all. Reading the three equations side by side, the trend is that the less reactive the metal, the further its nitrate collapses on heating: nitrite, then oxide, then bare metal.

There is no single quantitative formula to derive here — thermal stability of a nitrate is a pattern, not a numerical law — so the reasoning is worth spelling out plainly. A metal nitrate is held together by the metal ion's grip on the nitrate group. The more reactive a metal is, the more strongly it holds onto oxygen-containing groups, so its nitrate resists breaking down completely: it loses only one oxygen atom per nitrogen and settles at the nitrite stage. A moderately reactive metal cannot hold its oxide together as an intact nitrite; the nitrogen is expelled as NO₂ and what remains is the simpler metal oxide. A weakly reactive metal cannot even hold onto oxygen as an oxide under these conditions, so decomposition runs all the way to the free metal. This is exactly the same activity-series logic used elsewhere in the "Study of Compounds" unit — the more reactive the metal, the more stable its compounds tend to be against heat — applied here specifically to nitrates. The brown ring test works for a related reason: nitric acid's oxidising nature reappears in reverse. Freshly prepared iron(II) sulphate solution provides Fe²⁺ ions, which are themselves easily oxidised. When concentrated sulphuric acid is poured carefully down the side of a tilted test tube, it sinks below the nitrate solution without mixing, forming two separate layers. At that narrow interface, the concentrated acid generates nascent nitric oxide from the nitrate, and this nitric oxide combines with unreacted Fe²⁺ ions to form a dark brown complex, visible as a brown ring. The ring stays confined to the interface because the reaction only happens where the layers meet — which is exactly why careful, unmixed layering (not shaking) is essential to seeing it.

Problem. Silver nitrate is heated strongly in a dry test tube. Write the equation, and describe what you would observe, including the residue left behind.

Step 1. Identify which tier silver belongs to. Silver sits near the bottom of the activity series, well below hydrogen, so it is one of the least reactive metals — its nitrate decomposes all the way to the free metal on heating, not merely to an oxide or nitrite.

Step 2. Write the balanced equation for that tier:

Step 3. Translate the equation into observations. Two gases are evolved: brown fumes of nitrogen dioxide, and a colourless gas (oxygen) that can be confirmed by testing with a glowing splint, which relights in its presence. The residue left in the test tube is a grey or white shiny solid of metallic silver.

Competitive stretch

If instead lead nitrate were heated, the same reasoning but one tier up gives $2\text{Pb(NO}_3)_2 \xrightarrow{\Delta} 2\text{PbO} + 4\text{NO}_2\uparrow + \text{O}_2\uparrow$ — brown fumes and oxygen are evolved exactly as with silver, but the residue left behind is yellow lead(II) oxide, not a shiny metal, because lead is moderately reactive rather than weakly reactive.

• Expecting nitric acid to release hydrogen gas with reactive metals the way HCl or dilute H₂SO₄ do — it is an oxidising acid and typically evolves no H₂ with metals.
• Writing the same decomposition products for every metal nitrate, regardless of the metal — the products (nitrite+O₂, oxide+NO₂+O₂, or metal+NO₂+O₂) depend strictly on the metal's tier in the activity series.
• Confusing the brown ring test with other confirmatory tests, such as the white-fumes test for HCl/NH₃ or the lime-water test for CO₂.
• Pouring concentrated sulphuric acid carelessly (mixing rather than layering) during the brown ring test — the ring only forms at a sharp interface, so the acid must be added slowly down the side of a tilted tube.

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