R3.2.4 Acid-metal reactions

Acids Reacting with Reactive Metals to Release Hydrogen

Imagine you’re holding a piece of zinc metal.
What would happen if you dropped it into a beaker of dilute hydrochloric acid?
Almost instantly, bubbles of gas begin to form on the zinc’s surface, and the metal starts to dissolve.

The Chemistry of Acids and Metals

When acids react with reactive metals, the metal atoms lose electrons to form positive ions, while hydrogen ions $H^{+}$ from the acid gain electrons to form hydrogen gas $H_{2}$ .
This process is a redox reaction—a reaction involving both oxidation (loss of electrons) and reduction (gain of electrons).

The general word equation for this reaction is:

$Metal + Acid → Salt + Hydrogen Gas$

Example

When zinc reacts with hydrochloric acid, the products are zinc chloride (a salt) and hydrogen gas:

$Z n (s) + 2 H C l (a q) \to Z n C l_{2} (a q) + H_{2} (g)$

Here’s what happens at the atomic level:

Oxidation: Zinc atoms lose two electrons to form zinc ions ( $Z n^{2 +}$ ):
$Z n (s) \to Z n^{2 +} (a q) + 2 e^{-}$
Reduction: Hydrogen ions ( $H^{+}$ ) from the acid gain electrons to form hydrogen gas:
$2 H^{+} (a q) + 2 e^{-} \to H_{2} (g)$

Together, these half-equations illustrate the electron transfer from zinc to hydrogen, which is the hallmark of a redox reaction.

Tip

When writing half-equations, always check that the number of electrons lost in oxidation matches the number gained in reduction.

Example

Reaction Between Zinc and Sulfuric Acid:

When zinc reacts with sulfuric acid, a salt (zinc sulfate) and hydrogen gas are produced:

$Z n (s) + H_{2} S O_{4} (a q) \to Z n S O_{4} (a q) + H_{2} (g)$

In this reaction:

Zinc displaces hydrogen from sulfuric acid due to its higher reactivity.
Hydrogen gas is released as bubbles.
Zinc sulfate is formed as an aqueous salt in the solution.

This is a typical example of an acid reacting with a reactive metal to produce hydrogen gas.

Why Do Only Certain Metals React with Acids?

Not all metals react with acids to release hydrogen gas.
The key lies in the reactivity series, which ranks metals based on their tendency to lose electrons and form positive ions.

Hint

Metals higher in the series (e.g., magnesium, zinc, iron) react readily with acids because they are easily oxidized.
Metals lower in the series (e.g., copper, silver, gold) do not react with acids under normal conditions because they resist oxidation.

Common Mistake

Students often assume that all metals react with acids. Remember, only metals above hydrogen in the reactivity series can displace hydrogen ions to form hydrogen gas.

Example

For example, when iron reacts with sulfuric acid, it forms iron(II) sulfate and hydrogen gas:
$F e (s) + H_{2} S O_{4} (a q) \to F e S O_{4} (a q) + H_{2} (g)$
This reaction is used in industrial processes to clean iron surfaces before further processing.

Note

Hydrochloric acid and sulfuric acid are commonly used in these reactions because they are strong acids that readily release hydrogen ions ( $H^{+}$ ).

Experimental Observations and Tests for Hydrogen Gas

In the laboratory, you can observe this reaction through the release of bubbles (effervescence) as hydrogen gas is produced.
To confirm the gas is hydrogen, perform the "pop test":
1. Collect the gas in an inverted test tube.
2. Bring a lit splint close to the mouth of the test tube.
3. If the gas is hydrogen, it will ignite with a characteristic "pop" sound.

Tip

Always ensure the reaction vessel is open or vented to prevent pressure buildup from hydrogen gas.

Predicting and Writing Balanced Equations

Let’s practice predicting the products of metal-acid reactions and writing balanced equations.

Example

Reaction of Magnesium with Hydrochloric Acid

Magnesium reacts vigorously with hydrochloric acid:
$M g (s) + 2 H C l (a q) \to M g C l_{2} (a q) + H_{2} (g)$

Oxidation: $M g (s) \to M g^{2 +} (a q) + 2 e^{-}$
Reduction: $2 H^{+} (a q) + 2 e^{-} \to H_{2} (g)$

Self review

Write the balanced equation for the reaction between aluminum and hydrochloric acid. Identify the oxidation and reduction half-equations.

Common Mistake

Not all acids react with metals to release hydrogen gas. For example:

Weak acids like acetic acid react slowly.
Concentrated oxidizing acids like nitric acid often produce nitrogen oxides instead of hydrogen gas.

Common Mistake

When writing equations, don’t forget to include the diatomic nature of hydrogen gas ( $H_{2}$ ) in the products.

Reflection

Theory of Knowledge

Consider how redox reactions apply to other fields, such as biology (e.g., respiration) or environmental science (e.g., corrosion).
How do these fields adapt the definitions of redox processes to fit their needs?

You've read 2/2 free chapters this week.

Upgrade to PLUS or PRO to unlock all notes, for every subject.