Concentrated sulphuric acid

Pour dilute sulphuric acid onto zinc and you get the ordinary story — hydrogen gas bubbles off, exactly as any acid reacting with a reactive metal should. Pour hot concentrated sulphuric acid onto copper — a metal that dilute acids cannot touch at all — and something else happens entirely: the copper dissolves, the solution turns blue, and a choking gas comes off that is not hydrogen. Concentrated H₂SO₄ has stopped behaving like a normal acid and started behaving like an oxidising agent. Pour a few drops onto white sugar and a third personality appears: the sugar blackens, swells, and gives off heat, as the acid rips water out of a solid that already looks bone dry. This article ties both special powers — oxidising and dehydrating — to one underlying cause, works a full example on each, and lists the acid's major industrial uses.

The core insight

Concentrated (not dilute) sulphuric acid is a strong oxidising agent toward hot metals and non-metals, and a powerful dehydrating agent toward compounds containing combined hydrogen and oxygen, because as a hot, nearly water-free liquid it can both accept electrons (getting reduced to SO₂) and seize water molecules (or the H and O atoms that make them) far more aggressively than its dilute, heavily hydrated form can.

The core idea

Two personalities, one acid. Dilute sulphuric acid is H₂SO₄ diluted in a large excess of water; it behaves like a textbook acid — it reacts with reactive metals (above hydrogen in the activity series) by simple hydrogen displacement, releasing H₂ gas. Concentrated sulphuric acid is the acid with little or no water present, and hot concentrated H₂SO₄ does something dilute acid cannot: it oxidises the substance it touches, and it is itself reduced. With copper — a metal below hydrogen, so unreactive with any dilute acid — hot concentrated acid reacts on heating:

\text{Cu} + 2\text{H}_2\text{SO}_4(\text{conc.}) \xrightarrow{\Delta} \text{CuSO}_4 + \text{SO}_2\uparrow + 2\text{H}_2\text{O}

where: Cu is copper metal, H₂SO₄(conc.) is hot concentrated sulphuric acid, CuSO₄ is copper sulphate (the blue solution), and SO₂ is sulphur dioxide, the gas evolved instead of hydrogen. The acid also oxidises certain non-metals such as carbon, again releasing SO₂:

\text{C} + 2\text{H}_2\text{SO}_4(\text{conc.}) \xrightarrow{\Delta} \text{CO}_2\uparrow + 2\text{SO}_2\uparrow + 2\text{H}_2\text{O}

Dehydrating action. Concentrated H₂SO₄ has an intense affinity for water — strong enough to strip water not just from wet surfaces but from inside molecules that contain hydrogen and oxygen in the 2:1 ratio of water, even when no free water is present. The classic case is cane sugar, sucrose C₁₂H₂₂O₁₁:

\text{C}_{12}\text{H}_{22}\text{O}_{11} \xrightarrow{\text{conc. H}_2\text{SO}_4} 12\text{C} + 11\text{H}_2\text{O}

where: C₁₂H₂₂O₁₁ is sucrose, conc. H₂SO₄ is the dehydrating agent (it is not consumed in a simple 1:1 sense — it grabs the water and dilutes itself), 12 C is the black, spongy carbon residue left behind, and 11 H₂O is the water pulled out of the sugar's structure. The same dehydrating action chars paper and cloth, since both are largely cellulose, another carbohydrate.

Figure

Two beakers side by side showing copper with concentrated sulphuric acid releasing sulphur dioxide gas and turning blue on the left, and sugar with concentrated sulphuric acid turning into black spongy carbon on the right

Why it works

Both special behaviours trace back to the same fact: concentrated sulphuric acid is a hot, nearly anhydrous, highly polar liquid, not a dilute solution swimming in water. In dilute acid, every H₂SO₄ molecule is already surrounded and satisfied by water molecules, so the only chemistry available is the standard acid one — donating H⁺ ions to a reactive metal and being displaced as H₂ gas. Concentrated acid has almost no water of its own, so the sulphur atom (in its highest, +6 oxidation state) is free to act as an electron acceptor instead: it can pull electrons from a metal or non-metal, get reduced from S(+6) to S(+4), and leave as SO₂ gas. This is why copper, which cannot lose electrons to H⁺ in dilute acid (copper sits below hydrogen in the activity series), can still lose electrons to the sulphur atom itself in hot concentrated acid — no hydrogen displacement is needed at all. The dehydrating power comes from the same "water-starved" character. Concentrated H₂SO₄ mixing with water is strongly exothermic — the molecule genuinely "wants" to become hydrated — so when it meets a compound like sugar that contains hydrogen and oxygen in the same 2:1 ratio as water, it is thermodynamically favourable for the acid to tear those atoms out as H₂O, even though no free water was present to begin with. What is left behind, once the hydrogen and oxygen are stripped away, is just the carbon skeleton — hence the black, carbon-rich residue. Dilute acid, already saturated with water, has no comparable pull and cannot do this.

Worked example

Problem: 6.4 g of copper turnings are heated with excess concentrated sulphuric acid until the reaction is complete. Identify the products, state the colour and gas observed, and find the mass of copper sulphate formed. (Cu = 64, S = 32, O = 16.)

Step 1. Write the balanced equation for hot concentrated acid on copper — this is the oxidising reaction, not a hydrogen-displacement reaction, because copper lies below hydrogen in the activity series and cannot react with any dilute acid:

\text{Cu} + 2\text{H}_2\text{SO}_4(\text{conc.}) \xrightarrow{\Delta} \text{CuSO}_4 + \text{SO}_2\uparrow + 2\text{H}_2\text{O}

Step 2. Find moles of copper reacting. Molar mass of Cu is 64 g/mol, so

n(\text{Cu}) = \frac{6.4}{64} = 0.1\ \text{mol}

Step 3. Use the 1:1 mole ratio between Cu and CuSO₄ in the equation to get moles of copper sulphate formed: $n(\text{CuSO}_4) = 0.1\ \text{mol}$ . The molar mass of CuSO₄ is $64 + 32 + 4(16) = 160\ \text{g/mol}$ , so the mass formed is

m(\text{CuSO}_4) = 0.1 \times 160 = 16\ \text{g}

Step 4. State the observations to complete the answer: the copper dissolves and the solution turns blue (dissolved Cu²⁺ as copper sulphate), while a colourless, choking gas with a sulphurous smell (SO₂) is evolved — this gas turns moist blue litmus red and turns acidified potassium dichromate paper green, confirming it is sulphur dioxide and not hydrogen.

The mass of copper sulphate formed is 16 g, with SO₂ gas evolved and the solution turning blue — no hydrogen gas appears, unlike a reactive metal in dilute acid.

Competitive stretch

If the same 0.1 mol of copper were instead added to excess dilute sulphuric acid, nothing would happen at room temperature: copper is below hydrogen in the activity series, so dilute acid — which relies purely on hydrogen displacement — cannot touch it. The 16 g of CuSO₄ and the SO₂ gas are only obtained because the acid was concentrated and heated, letting sulphur itself act as the oxidising agent.

Common traps

Mixing up the two special properties: the oxidising reaction (with metals/non-metals like copper and carbon, evolving SO₂) and the dehydrating reaction (with carbohydrates like sugar, leaving carbon) are separate behaviours with separate examples — do not use one equation to explain the other's observations.
Expecting hydrogen gas from concentrated H₂SO₄ reacting with a metal — no H₂ is ever produced here; the gas is SO₂ because the acid itself is reduced.
Assuming these properties belong to sulphuric acid in general — oxidising and dehydrating behaviour are shown only by the hot, concentrated acid, never by the dilute form.
Describing the sugar reaction as "burning" — it is a chemical dehydration by the acid (water is removed as H and O in 2:1 ratio), not combustion, even though it looks charred and releases heat.

Ask the tutor: Ask the AI tutor for a side-by-side table comparing how dilute and concentrated sulphuric acid each react with the same metal.

Why does copper react with hot concentrated sulphuric acid but not with dilute sulphuric acid at all

Dilute sulphuric acid can only react with a metal by displacing hydrogen, which requires the metal to be above hydrogen in the activity series — copper is below it, so no reaction happens, hot or cold. Hot concentrated sulphuric acid does not rely on hydrogen displacement at all; instead the sulphur atom itself accepts electrons from copper and gets reduced to SO₂, so the reaction proceeds by oxidation-reduction rather than simple displacement, and copper's position in the activity series becomes irrelevant.

Is concentrated sulphuric acid dehydrating sugar the same thing as burning it

No. Burning (combustion) needs oxygen from the air and produces carbon dioxide and water as the carbon is oxidised away. Dehydration by concentrated sulphuric acid needs no external oxygen at all — the acid simply pulls hydrogen and oxygen atoms out of the sugar molecule in the same 2:1 ratio as water, leaving the carbon behind as a solid black residue. The strong heat given off during dehydration can make it look like burning, but the chemistry and the products are different.

If concentrated sulphuric acid oxidises carbon to carbon dioxide, why does it also leave carbon behind from sugar

These are two separate, different reactions with different substances. With sugar, the acid's dehydrating action removes water, leaving elemental carbon as a residue — the carbon itself is not further oxidised in that reaction. Separately, if concentrated sulphuric acid is heated directly with carbon (not sugar), its oxidising action converts that carbon into CO₂ while the acid is reduced to SO₂. So whether carbon appears as a product or is consumed depends on which of the acid's two properties — dehydrating or oxidising — is being demonstrated.