Laboratory preparation of carbon dioxide

Drop a handful of marble chips into a flask, pour in some dilute hydrochloric acid, and within seconds the mixture starts fizzing — bubbles of carbon dioxide streaming out through a delivery tube. This is one of the simplest and most reliable gas preparations done in a school laboratory, and it is also a small case study in chemical reasoning: why marble and not some other carbonate, why hydrochloric acid and not sulphuric acid, and why the collecting jar sits mouth-upward rather than inverted. Every choice in the setup traces back to a property of the substances involved — solubility, density, and reaction rate. This article rebuilds the full procedure, explains the balanced equation behind it, works through the standard exam questions on the topic, and collects the traps students most often fall into when writing this preparation from memory.

The core insight

Carbon dioxide is prepared in the laboratory by reacting marble chips (calcium carbonate) with dilute hydrochloric acid, and the gas — being denser than air — is collected in a mouth-upward jar by the upward displacement of air.

The core idea

The reactants. Carbon dioxide is generated by treating a carbonate with a dilute acid. The standard laboratory choice is marble chips, a naturally occurring, easily available form of calcium carbonate ( $\text{CaCO}_3$ ), combined with dilute hydrochloric acid ( $\text{HCl}$ ). Sodium carbonate ( $\text{Na}_2\text{CO}_3$ ) works as an alternative solid reactant in exactly the same way, though marble chips remain the standard choice because they are cheap, inert to store, and react at a controllable rate.

The reaction. Adding dilute hydrochloric acid to marble chips produces effervescence almost immediately:

\text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)

where: $\text{CaCO}_3(s)$ is solid calcium carbonate (the marble chips), $\text{HCl}(aq)$ is hydrochloric acid in aqueous (dissolved) form, $\text{CaCl}_2(aq)$ is aqueous calcium chloride, and $\text{CO}_2(g)$ is the gaseous product collected. The coefficient 2 in front of $\text{HCl}$ is essential: one calcium atom, two chlorine atoms, and two hydrogen atoms must balance on both sides. This is a double displacement reaction — calcium and hydrogen exchange partners between the carbonate ion and the chloride ion. A carbonic acid molecule ( $\text{H}_2\text{CO}_3$ ) forms only momentarily as an intermediate; it is unstable and decomposes at once into water and carbon dioxide gas, which escapes as the bubbles you see.

Collection method. Carbon dioxide is collected by the upward displacement of air. Because $\text{CO}_2$ is denser than air, gas entering a collecting jar through a delivery tube sinks toward the bottom and steadily pushes the lighter air upward and out through the mouth of the jar, which is kept facing upward throughout. The jar fills from the bottom up, exactly the opposite of how a gas lighter than air (such as hydrogen) would be collected.

Figure

Apparatus for preparing carbon dioxide from marble chips and dilute hydrochloric acid showing the acid dripping onto the chips in a flask and the gas travelling through a delivery tube into a mouth-upward collecting jar

Why it works

The setup is not arbitrary — every part of the procedure follows from a property of the chemicals involved, and this reasoning is exactly what exam questions probe.

Why marble chips. Marble is a convenient, inexpensive, and widely available solid form of calcium carbonate. It reacts with dilute acid at a steady, moderate pace rather than violently, which allows a controlled and continuous stream of gas to be collected over the course of an experiment — a powder would react too fast and too unevenly for a smooth, sustained supply.

Why dilute hydrochloric acid, not concentrated or sulphuric acid. Dilute hydrochloric acid reacts with calcium carbonate at a manageable, observable rate; using it in concentrated form would make the reaction too fast and hazardous to control safely in a school setting. Dilute sulphuric acid, meanwhile, is specifically avoided for this preparation: the calcium sulphate ( $\text{CaSO}_4$ ) it produces is only slightly soluble in water, and it coats the surface of the remaining marble chips. That coating blocks fresh acid from reaching the unreacted carbonate underneath, so the reaction slows drastically and effectively stops well before all the marble has reacted. Hydrochloric acid avoids this problem entirely because the calcium chloride it forms is freely soluble and simply dissolves into the solution, leaving the marble surface exposed for the reaction to continue.

Why upward displacement of air. The choice of collection method always depends on two properties of the gas: whether it is denser or lighter than air, and whether it is soluble in water. Carbon dioxide is denser than air, which rules out downward displacement (used for gases like hydrogen) and also rules out collection over water for many purposes, since $\text{CO}_2$ is somewhat soluble in water. Collecting it by upward displacement of air — gas entering a jar held mouth-upward, sinking to the bottom, and pushing the lighter air out — is therefore the natural method for this particular gas.

Testing for a full jar. Because carbon dioxide does not support combustion, a simple test confirms when the jar is full: a lit splint held at the mouth of the jar is extinguished once enough $\text{CO}_2$ has accumulated to fill the space and displace the air the flame needs.

Worked example

Question: Describe, step by step, how you would prepare and collect a jar of carbon dioxide gas in the laboratory starting from marble chips and dilute hydrochloric acid, explain why the jar is kept mouth-upward during collection, and state how you would confirm the jar is full.

Solution:

Step 1. Take marble chips (calcium carbonate) in a flask fitted with a thistle funnel and a delivery tube, so that acid can be added without disturbing the gas-tight setup.

Step 2. Pour dilute hydrochloric acid onto the marble chips through the thistle funnel. Effervescence begins immediately as the reaction proceeds:

\text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)

Step 3. Lead the gas produced through the delivery tube into a collecting jar kept mouth-upward. Because carbon dioxide is denser than air, it sinks to the bottom of the jar as it enters and gradually displaces the lighter air upward and out — this is the method of upward displacement of air, and it is the reason the jar must face upward rather than be inverted.

Step 4. Confirm the jar is full by holding a lit splint at its mouth: the flame is extinguished once carbon dioxide has accumulated enough to fill the jar, since $\text{CO}_2$ does not support combustion.

Final answer: marble chips reacted with dilute HCl generate CO₂ gas, which is collected in a mouth-upward jar by upward displacement of air (because CO₂ is denser than air) and confirmed complete when a lit splint held at the mouth is extinguished.

Competitive stretch

Question: A student reacts marble chips with dilute sulphuric acid instead of dilute hydrochloric acid and finds gas production slows and almost stops after a short time. Explain why, and state the correct acid to use.

Step 1. The attempted reaction is $\text{CaCO}_3(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{CaSO}_4 + \text{H}_2\text{O}(l) + \text{CO}_2(g)$ .

Step 2. The calcium sulphate formed is only slightly soluble in water, so it coats the surface of the remaining marble chips rather than dissolving away.

Step 3. This coating blocks acid from reaching the unreacted calcium carbonate beneath it, so the reaction slows and effectively halts.

Final answer: the insoluble calcium sulphate coating stops the reaction; dilute hydrochloric acid should be used instead, since the calcium chloride it forms is freely soluble and does not coat the chips.

Common traps

Writing the equation unbalanced — the acid needs a coefficient of 2 in front of $\text{HCl}$ so that calcium, chlorine, and hydrogen all balance on both sides.
Trying to collect carbon dioxide by downward displacement of air (as used for hydrogen). CO₂ is denser than air, so it must be collected by upward displacement of air with the jar mouth kept upward.
Assuming any dilute acid works equally well. Dilute sulphuric acid is a poor choice because the calcium sulphate produced coats the marble chips and stops the reaction.
Forgetting that sodium carbonate is a valid alternative solid reactant to marble chips, reacting with dilute hydrochloric acid in the same manner.

Ask the tutor: if you would like to walk through labelling this apparatus diagram yourself or practise balancing the equation from scratch, ask your AI tutor to guide you through it.

Why is marble used instead of just any calcium carbonate, like chalk powder?

Marble chips are used because they are cheap, easy to obtain, and react with dilute acid at a steady, controllable rate, giving a smooth and continuous supply of gas. A fine powder such as chalk has a much larger surface area, so it would react far too fast and vigorously, making the gas evolution difficult to control and the reaction harder to manage safely in a school lab. The lump form of marble chips is exactly what keeps the reaction rate manageable.

Could I collect carbon dioxide over water instead of by upward displacement of air?

It is not the preferred method because carbon dioxide is somewhat soluble in water — some of the gas would dissolve into the water rather than being collected, giving an inaccurate or reduced yield. Since CO₂ is also denser than air, upward displacement of air is both more accurate and simpler to set up, which is why it is the standard laboratory method taught for this preparation.

Why does the carbonic acid mentioned in the mechanism not appear in the final equation?

When calcium carbonate reacts with hydrochloric acid, carbonic acid ( $\text{H}_2\text{CO}_3$ ) does form as an intermediate step, but it is highly unstable and decomposes almost instantly into water and carbon dioxide gas. Because this decomposition happens essentially immediately, the overall balanced equation skips straight from reactants to the final stable products — water and carbon dioxide — rather than showing carbonic acid as a separate stage.