Preparation of HCl gas
Open a bottle of concentrated hydrochloric acid and it fumes in air — a sharp, choking smell that tells you at once this is a gas dissolved in water, not a liquid on its own. That gas is hydrogen chloride, HCl, and the ICSE Class 10 laboratory method for making it uses nothing more exotic than common salt and concentrated sulphuric acid, heated gently in a flask. This article walks through the two-stage reaction that produces HCl gas, explains in plain reasoning why each choice in the setup — concentrated acid, gentle then strong heating, downward delivery, a sulphuric-acid drying bottle — is forced by the properties of the gas itself, and finishes with the confirmatory white-fumes test examiners love to ask about. Every equation and number here matches the syllabus scope: no reaction mechanism, only observation-based chemistry you can reproduce on paper in an exam.
Heating common salt with concentrated (never dilute) sulphuric acid displaces HCl gas in two temperature-dependent steps, and the gas's high density and extreme water-solubility force downward delivery collection instead of collection over water.
The core idea
The reagents. Solid sodium chloride (common salt) is heated with concentrated sulphuric acid in a round-bottomed flask fitted with a thistle funnel and a delivery tube. Concentrated acid is essential — it behaves as a strong, non-volatile acid capable of displacing the more volatile HCl from the salt on heating. Dilute sulphuric acid is mostly water and simply cannot do this displacement.
Stage 1 — below about 200°C, gentle heating gives:
where: NaCl is sodium chloride (the salt), H₂SO₄ is concentrated sulphuric acid, NaHSO₄ is sodium bisulphate (sodium hydrogen sulphate), and HCl↑ is hydrogen chloride gas escaping the mixture.
Stage 2 — above about 200°C, once gas evolution from stage 1 slows, stronger heating drives the bisulphate to react with more salt:
where: Na₂SO₄ is sodium sulphate, the final solid residue left in the flask once all the salt is consumed.
Collection. HCl is denser than air and dissolves in water almost explosively (the classic "fountain experiment"), so it is collected dry by downward delivery — upward displacement of air — in an upright gas jar, never over a water trough.
Drying. The moist gas is first bubbled through concentrated sulphuric acid in a Woulfe's bottle, which soaks up water vapour without reacting with HCl, so a dry gas reaches the collecting jar.
Why it works
Every step in this preparation is forced by two facts about HCl gas: it is denser than air, and it is exceptionally soluble in water. Nothing here is arbitrary.
Why concentrated acid, not dilute. The reaction only proceeds if the acid used is a stronger, less volatile acid than HCl itself, so that it can push HCl out of the salt. Concentrated H₂SO₄ satisfies this; diluting it with water removes that advantage, so no useful amount of gas is produced.
Why heat gently first, then strongly. Starting with strong heat makes stage 1 violent, and the escaping gas can carry droplets of sulphuric acid with it — this spray contaminates the collected gas and is a safety hazard. Gentle heat lets stage 1 proceed steadily. Only once that first stage slows down is the heat increased, to push the leftover sodium bisulphate into stage 2 and use up the remaining salt.
Why downward delivery, not collection over water. Gases like hydrogen are collected over a water trough because they barely dissolve in water. HCl is the opposite case — it is so soluble that a jar of the gas inverted over water causes water to rush up and fill the jar almost instantly (the fountain experiment). Collecting HCl over water would mean the gas dissolves into the trough rather than displacing it, so essentially nothing is collected. Because HCl is also denser than air, it can instead be trapped by simply letting it sink to the bottom of an upright jar and push the lighter air out through the top — downward delivery, or upward displacement of air.
Why concentrated H₂SO₄ again, this time as a drying agent. The gas leaving the reaction flask carries moisture. Passing it through a second, separate portion of concentrated sulphuric acid removes this water vapour without reacting with the HCl itself, because concentrated H₂SO₄ has a strong affinity for water but does not react with dry HCl gas. Note this is a different role from the first use of concentrated H₂SO₄ — one is the reagent, the other is the desiccant.
Worked example
Problem: A student prepares HCl gas by heating sodium chloride with concentrated sulphuric acid, keeping the temperature below 200°C throughout. State the equation for the reaction, describe how the gas should be collected and dried, and give the confirmatory test that proves the gas collected really is HCl.
Step 1. Below 200°C only the first stage of the reaction occurs, so the equation is:
Step 2. Collection: since HCl is denser than air and extremely soluble in water, it is collected by downward delivery (upward displacement of air) into a dry, upright gas jar — never over a water trough.
Step 3. Drying: before reaching the jar, the moist gas is passed through concentrated sulphuric acid in a Woulfe's bottle, which removes water vapour without reacting with the HCl.
Step 4. Confirmatory test: a glass rod dipped in ammonium hydroxide is brought near the mouth of the jar. Dense white fumes of ammonium chloride form:
Answer: the equation is NaCl + H₂SO₄ → NaHSO₄ + HCl↑ (below 200°C); the gas is collected dry by downward delivery and dried over concentrated H₂SO₄; and it is confirmed by dense white fumes of NH₄Cl with an ammonium hydroxide-dipped rod.
Competitive stretch
If the same flask were heated strongly from the very start instead of gently, expect two consequences an examiner may probe: first, the vigorous reaction risks spraying droplets of sulphuric acid into the delivery tube along with the gas, contaminating the sample; second, because both stages can now proceed together, the equation actually representing the gas evolved is a mixture, and a well-prepared answer should still quote both equations (stage 1 and stage 2) and explain that gentle heating is preferred specifically to keep the process controlled and stage-wise.
Common traps
- Writing dilute sulphuric acid as the reagent — only concentrated H₂SO₄ works, since it must act as a strong, non-volatile acid to displace HCl.
- Forgetting the two-stage nature of the reaction; many students only give the first equation and miss the above-200°C step with excess NaCl.
- Trying to collect HCl over water by analogy with hydrogen or oxygen — HCl's extreme solubility rules this out completely.
- Mixing up the two roles of concentrated H₂SO₄ — one portion is the reagent that reacts with NaCl, a separate portion is the drying agent that only removes moisture.
Ask the tutor: if you want a labelled walkthrough of why the fountain experiment happens or a comparison with how other gases like ammonia are dried and collected, just ask the AI tutor about "Preparation of HCl gas."
Why can we not just collect HCl the same way we collect hydrogen, over a trough of water?
Hydrogen is collected over water because it barely dissolves in it, so the gas simply pushes water out of an inverted jar. HCl behaves in exactly the opposite way — it is so soluble that a jar of the gas placed over water causes water to rush up and fill the jar almost instantly, an effect demonstrated by the fountain experiment. If you tried the water-trough method with HCl, nearly all of the gas would dissolve into the trough water instead of being collected, leaving you with essentially nothing in the jar.
Both steps use concentrated sulphuric acid — is that not the same acid doing two things at once?
It is the same chemical, concentrated H₂SO₄, but used in two separate, physically distinct portions for two different purposes. The first portion sits in the reaction flask with the sodium chloride and takes part in the actual reaction, being consumed as it forms sodium bisulphate or sodium sulphate. The second portion sits in a separate Woulfe's bottle downstream, purely to soak moisture out of the gas stream — it does not react with the dry HCl gas passing through it at all. Keep the two roles, reagent and desiccant, clearly separate in an answer.
Why does the reaction need two different temperature stages instead of just going straight to completion
At a lower temperature, only enough energy is available to break NaCl and H₂SO₄ apart into NaHSO₄ and HCl — the first, easier step. The sodium bisulphate formed is itself capable of reacting further with leftover salt, but this second reaction needs a higher temperature to proceed. Heating gently at first lets the first stage go steadily and safely; only once that supply of gas starts slowing down does raising the temperature make sense, since by then most of the low-temperature pathway has been used up and pushing harder recovers the remaining HCl locked in the leftover salt and bisulphate.