Dilute sulphuric acid
Pour concentrated sulphuric acid into a large excess of water and something changes completely. The oily, charring liquid that eats through sugar and paper becomes an ordinary, safe-to-handle acid — dilute sulphuric acid, H₂SO₄(aq) — that fizzes with metals, neutralises bases, and releases carbon dioxide from carbonates just like dilute hydrochloric acid does. The chemistry does not get more exotic with sulphuric acid; it gets more typical, because dilution lets the acid ionise freely. This article covers the general acid reactions of the dilute acid — with metals, bases, and carbonates or bicarbonates — and the confirmatory test for the sulphate ion using barium chloride. It deliberately stops short of the concentrated acid's special oxidising and dehydrating behaviour, which belongs to a separate topic. By the end you should be able to write balanced equations for every reaction type here, and describe the sulphate test precisely enough to defend it against a false positive.
Dilute sulphuric acid ionises freely in water and behaves as a typical strong dibasic acid — reacting with metals, bases, and carbonates exactly as dilute HCl does — while the sulphate ion it carries is confirmed by a barium chloride precipitate that survives acid washing.
The core idea
It is a typical strong acid, not a special one. Concentrated H₂SO₄ is mostly molecular and poorly ionised; adding a large excess of water breaks it apart into ions that move and react freely:
where: H₃O⁺ is the hydronium ion responsible for acidic character (turning blue litmus red), and SO₄²⁻ is the sulphate ion carried into every salt the acid forms.
It reacts with reactive metals to liberate hydrogen gas. Metals above hydrogen in the activity series (zinc, magnesium, iron) displace H₂:
It reacts with bases in neutralisation, and with carbonates or bicarbonates to release carbon dioxide. Both follow the ordinary pattern of any dibasic acid:
It is dibasic, so it can give two families of salt. With two replaceable hydrogens, the amount of base decides whether the product is a normal salt (Na₂SO₄) or an acidic salt (NaHSO₄) — a distinction unique to di- and tri-basic acids, absent for a monobasic acid like HCl.
Why it works
Dilution is the whole story. Concentrated H₂SO₄ has very little free water, so it stays largely as intact molecules and behaves as a powerful dehydrating and oxidising agent rather than an ordinary acid — that stronger, separate behaviour is deliberately excluded from this article. Adding a large excess of water breaks the H₂SO₄ apart into H₃O⁺ and SO₄²⁻ ions, and it is these free-moving, hydrated ions that carry out the familiar acid chemistry: H₃O⁺ is what turns litmus red, gets reduced to H₂ at a reactive metal surface, and is neutralised by OH⁻ from a base or by the carbonate/bicarbonate ion (which breaks down its own weak acid, carbonic acid, into water and CO₂). Because both replaceable hydrogens of H₂SO₄ can ionise, the acid is dibasic — it can hand over one H⁺ or two, which is exactly why it can build both a normal salt and an acidic salt from the same starting materials, depending on how much base is present to accept the protons.
Worked example
Problem: Dilute sulphuric acid is added to magnesium ribbon in a test tube. Write the balanced equation, describe what you would observe, and state how you would confirm the gas evolved. Then say what salt forms if instead excess NaOH, rather than a limited amount, is reacted with the same acid.
Step 1. Magnesium sits above hydrogen in the activity series, so it displaces hydrogen gas from the dilute acid:
Step 2. Observation: brisk effervescence of a colourless, odourless gas; the magnesium ribbon dissolves gradually as the reaction proceeds, and the test tube becomes noticeably warm because the reaction is exothermic.
Step 3. Confirming the gas: bring a lit splint near the mouth of the test tube. The gas ignites with a characteristic "pop," confirming it is hydrogen.
Step 4. For the second part, reacting H₂SO₄ with excess NaOH drives the neutralisation all the way, using both replaceable hydrogens:
With excess base, the product is the normal salt sodium sulphate, Na₂SO₄ — not the acidic salt NaHSO₄, which would only form if the base were limited to roughly half that amount.
Competitive stretch
If a student is instead given a solution and asked to confirm it contains sulphate ions rather than carbonate ions, the test is: add dilute HCl first to destroy any carbonate present, then add barium chloride solution. A dense white precipitate that persists after the HCl step confirms sulphate:
Barium sulphate is insoluble even in dilute acids, which is exactly why it survives the HCl wash while barium carbonate (the false-positive risk) would have dissolved.
Common traps
- Testing for sulphate with barium chloride alone, skipping the dilute HCl step — this risks a false positive, since barium carbonate is also a white precipitate but dissolves in dilute acid.
- Assuming dilute H₂SO₄ shows the concentrated acid's oxidising or dehydrating behaviour — dilution suppresses that entirely; the dilute acid is a typical acid only.
- Forgetting that H₂SO₄ is dibasic and can give an acidic salt (NaHSO₄) as well as a normal salt (Na₂SO₄), depending on how much base reacts.
- Writing an unbalanced equation with NaOH — remember 2 mol NaOH are needed per mol H₂SO₄ to reach the normal salt.
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Why does dilute sulphuric acid behave so differently from the concentrated acid if it is the same compound?
Concentration decides how ionised the acid is. Concentrated H₂SO₄ has very little free water, so it stays mostly as intact molecules and shows its own special behaviour — strong dehydration (charring sugar or paper) and oxidation of metals like copper, which do not react with the dilute acid at all. Dilution supplies a large excess of water, letting the acid ionise fully into H₃O⁺ and SO₄²⁻, and it is these free ions that give the ordinary, predictable acid reactions covered here. Same compound, very different behaviour, purely because of how much water is present.
Can I use barium chloride to test for sulphate without adding HCl first, just to save a step?
You can, but the result becomes unreliable. Barium chloride reacts with both sulphate ions and carbonate ions to give a white precipitate — BaSO4 and BaCO3 look identical to the eye at first. If the sample happens to contain carbonate as well as, or instead of, sulphate, you would wrongly conclude sulphate is present. Adding dilute HCl first destroys any carbonate (releasing CO2 gas), so any white precipitate that appears afterward with barium chloride can only be barium sulphate, which does not dissolve in dilute acid. Skipping that step trades a reliable test for a faster but unsafe one.
Since H2SO4 has two hydrogens, does it always form the acidic salt NaHSO4 instead of Na2SO4?
No — which salt forms depends entirely on how much base is supplied, not on the acid's basicity alone. If NaOH is present in excess (or in the exact 2:1 mole ratio to H2SO4), both hydrogens are neutralised and the product is the normal salt Na2SO4. If NaOH is limited to roughly half that amount, only one hydrogen is neutralised per H2SO4 molecule, leaving the acidic salt NaHSO4. Dibasic simply means the acid can give up two protons; whether it actually does depends on the proportions used in the reaction.