Laboratory preparation

The reaction. Hydrogen is prepared in the school laboratory by treating a moderately reactive metal — zinc is the standard choice — with a dilute acid. Dilute sulphuric acid is the usual partner:

where: $\text{Zn}(s)$ is solid zinc granules, $\text{dil. H}_2\text{SO}_4(aq)$ is dilute sulphuric acid in water, $\text{ZnSO}_4(aq)$ is the dissolved zinc sulphate salt left behind, and $\text{H}_2(g)$ is the hydrogen gas that escapes as bubbles. Dilute hydrochloric acid works just as well and gives the same style of reaction, only the salt changes:

where: the metal and the gas produced are identical to the sulphuric-acid case; only the acid and the resulting chloride salt differ. Note the word "dilute" in both equations — it is not decorative. Concentrated sulphuric acid is a strong oxidiser and reacts with zinc quite differently (producing sulphur dioxide instead of clean hydrogen), so this specific preparation depends on the acid being dilute.

Why zinc, specifically. Three practical reasons make zinc the metal of choice for this preparation, not any of the dozens of other metals that could in principle react with an acid:

• It reacts with dilute acids at a moderate, steady, controllable rate — brisk enough to produce a usable stream of gas, slow enough to be safe to watch and time.
• It is cheap, reasonably pure, and available in granulated form, which gives a large reacting surface without the metal dissolving away too fast.
• It sits at a "just right" point in the reactivity order for classroom purposes: markedly more reactive metals would react too violently to control, and markedly less reactive metals would barely react at all.

Collecting the gas. Hydrogen is collected by downward displacement of water — an inverted, water-filled jar sitting in a trough, into which the delivery tube carries gas that bubbles up and collects at the closed top, pushing water down and out as it accumulates. This method works precisely because hydrogen is:

• Insoluble in water, so none of the gas is lost by dissolving into the water it bubbles through on its way up.
• Lighter than air, which is why downward displacement of air is also possible for hydrogen — but collection over water is the standard classroom method here, since it lets you watch the gas volume grow as a visible rising bubble of water displaced.

This preparation is entirely descriptive chemistry — there is no numeric formula to derive — but the choices behind it follow a clear chain of reasoning that examiners like to test directly.

Start with the reactivity constraint. Any metal-plus-acid preparation needs a reaction that is neither dangerously fast nor uselessly slow. Very reactive metals (sodium, potassium) react with dilute acids so violently that the reaction is unsafe to run in an open flask — heat and gas are released far too quickly to control or collect cleanly. Very unreactive metals (copper, silver) do not react with dilute acids at all, so no hydrogen is produced. Zinc occupies the narrow, useful middle ground: reactive enough to give a steady stream of bubbles within seconds, controlled enough that the reaction can run safely for minutes on a bench.

Then add the acid constraint. The acid must also be dilute. A dilute acid supplies hydrogen ions ($\text{H}^+$) at a manageable concentration, giving the same steady, moderate reaction rate as the metal choice. Concentrated acids change the chemistry of the reaction itself (concentrated sulphuric acid, for instance, acts as an oxidising agent rather than simply supplying $\text{H}^+$ ions), so "dilute" is a required condition of the preparation, not an incidental detail.

Finally, reason about collection. Once hydrogen gas is being produced, it must be trapped without loss. An open container fails immediately: hydrogen is lighter than air and would simply rise and mix into the room. Passing the gas through a delivery tube into an inverted, water-filled jar solves this because hydrogen does not dissolve in the water it bubbles through — every molecule that enters the jar stays there, displacing water downward and out, so the jar fills with pure, uncontaminated hydrogen that can be sealed, tested, or measured by volume.

Question. A student sets up the laboratory preparation of hydrogen using zinc granules and dilute sulphuric acid, then repeats the experiment with dilute hydrochloric acid instead. Write both balanced equations with state symbols, and explain why the gas is collected by downward displacement of water rather than in an open container, and why a very reactive metal such as sodium would not be substituted for zinc.

Step 1. Identify the reactants and products for the sulphuric-acid case: zinc metal plus dilute sulphuric acid gives a soluble salt (zinc sulphate) plus hydrogen gas. The balanced equation, with state symbols, is $\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)$. Each side has one zinc atom, two hydrogen atoms and one sulphate group, so it is already balanced.

Step 2. Repeat for dilute hydrochloric acid: zinc plus dilute hydrochloric acid gives zinc chloride plus hydrogen. Balancing hydrogen and chlorine atoms requires two $\text{HCl}$ for every one $\text{Zn}$, giving $\text{Zn}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)$ — one zinc, two chlorine, and two hydrogen atoms on each side.

Step 3. Explain the collection method: since hydrogen is lighter than air, an open container would let the gas escape upward and mix into the room air rather than being trapped for study. Since hydrogen does not dissolve in water, bubbling it into an inverted, water-filled jar traps it cleanly — the gas accumulates at the closed top of the jar and pushes water down and out, with no gas lost to solution.

Step 4. Explain why sodium is unsuitable: sodium is far more reactive than zinc and reacts with dilute acids extremely fast and violently, releasing hydrogen too quickly and with too much heat to collect safely in an open lab setup. Zinc's moderate reactivity is precisely what makes it, rather than sodium or copper, the correct choice for a controlled classroom preparation.

Answer: $\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)$ and $\text{Zn}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)$; the gas is collected over water because it is insoluble and lighter than air, and sodium is rejected because it is too reactive to control.

• Forgetting the word "dilute" for the acid — concentrated acids do not behave the same way and are not used for this preparation.
• Assuming any metal can stand in for zinc — the choice is deliberate, based on zinc's moderate, controllable reaction rate, not just on it being "a metal that reacts with acid."
• Mixing up "downward displacement of water" (used here, because hydrogen is insoluble) with "downward displacement of air" (a different valid method, used because hydrogen is light, but not the standard one for this preparation).
• Writing an unbalanced equation — especially forgetting the coefficient 2 in front of HCl, or dropping state symbols such as $(s)$, $(aq)$ and $(g)$.

Ask the tutor: Want to test yourself with a diagram-labelling question on the apparatus, or compare this method side by side with the electrolysis of water?