Introduction

Reactions between metals and acids are fundamental redox processes that produce characteristic products. Understanding these reactions helps explain salt formation, hydrogen evolution, and redox behaviour.

Understanding Metal–Acid Reactions

When a metal reacts with an acid, a salt and hydrogen gas are formed. These reactions are essential illustrations of redox chemistry, where electrons are transferred between species. Students must write full equations rather than ionic equations, focusing on accurate product prediction and correct formula construction.

The General Reaction Pattern

Most metals above hydrogen in the reactivity series react with dilute acids. The products formed depend on the acid used and the charge of the metal ion produced.

Typical reaction pattern:

• Metal + Acid → Salt + Hydrogen

• The metal is oxidised, and the hydrogen ions in the acid are reduced.

This diagram summarises the general reaction of a metal with a dilute acid to form a salt and hydrogen gas. The example illustrates magnesium reacting with hydrochloric acid to produce magnesium chloride and hydrogen. This directly reflects the metal–acid equation patterns required in OCR A-Level Chemistry. Source

To support these ideas, the following definition clarifies oxidation.

In metal–acid reactions, metals lose electrons to form positive ions, while hydrogen ions gain electrons to form hydrogen gas.

Why Hydrogen Gas Is Produced

Acids contain hydrogen ions (H⁺), which act as oxidising agents. Because metals donate electrons readily, these ions are reduced to hydrogen gas. The salt formed depends entirely on the anion present in the acid.

Identifying the Salt Formed

When naming the salt:

• The first part of the salt’s name comes from the metal.

• The second part comes from the acid’s anion.

Common patterns include:

• Hydrochloric acid → chloride

• Sulfuric acid → sulfate

• Nitric acid → nitrate

These naming rules ensure consistency when constructing equation products.

Full Equations for Metal–Acid Reactions

The OCR specification emphasises full equations only, meaning that all reactants and products must appear as complete species.

Common Acids Used

Metals generally react with:

• Hydrochloric acid (HCl)

• Sulfuric acid (H₂SO₄)

• Nitric acid (HNO₃)

Students must account for acid concentration and formulae when determining stoichiometric relationships.

General Equation Format

The general full equation appears below.

These equations demonstrate how oxidation and reduction occur simultaneously, even though ionic equations are not required.

A short descriptive explanation helps distinguish situations where deviations may occur. For instance, some metals such as copper do not react with dilute acids because they are less reactive than hydrogen and cannot displace it from solution.

Stoichiometric Considerations

Although the specification does not require ionic equations here, it is essential to write balanced equations. This involves balancing:

• The atoms of each element

• The charge where applicable (if metal forms multiple oxidation states)

• The diatomic nature of hydrogen gas (H₂)

Balanced equations communicate the correct mole ratios needed for reaction understanding.

Redox Aspects of Metal–Acid Reactions

These reactions clearly illustrate redox behaviour:

• Metal atoms are oxidised to metal cations.

• Hydrogen ions are reduced to hydrogen gas.

Oxidation numbers are a useful tool for identifying these changes, though not explicitly required in equation writing for this subsubtopic.

Because these reactions always involve a transfer of electrons, they sit at the intersection of redox chemistry and acid reactions within the specification.

Factors Affecting Metal–Acid Reactivity

Several factors influence whether a reaction occurs and how vigorously it proceeds.

Position in the Reactivity Series

Metals higher in the series react more readily. For OCR A-Level, the key idea is that only metals more reactive than hydrogen displace hydrogen gas from acids.

Nature of the Acid

Different acids produce different salts:

• HCl → chlorides

• H₂SO₄ → sulfates

• HNO₃ → nitrates

Dilute nitric acid behaves similarly to the others at this level, though concentrated nitric acid can exhibit oxidising behaviour beyond the scope of this subsubtopic.

Metal Charge and Ion Formation

Some metals form multiple oxidation states. It is important to write the correct salt formula that reflects the oxidation state adopted in the reaction.

Salts produced in metal–acid reactions are typically soluble and remain in solution after hydrogen gas escapes.

A metal forming a +2 ion will produce salts such as magnesium sulfate (MgSO₄), while a metal forming a +1 ion will form salts such as sodium chloride (NaCl). Students should confirm oxidation states when constructing final equations.

Observations During the Reaction

Metal–acid reactions typically display easily observable features:

• Effervescence from hydrogen gas production

• Temperature increase due to exothermic behaviour

• Metal dissolution as the solid metal gradually disappears

• Formation of an aqueous salt solution

This photograph shows magnesium ribbon reacting with dilute hydrochloric acid in a test tube. The bubbles represent hydrogen gas released as magnesium forms magnesium chloride. The page also mentions collecting and igniting hydrogen, which exceeds syllabus requirements, but the image itself remains directly relevant. Source

These observations provide useful qualitative evidence for the reaction and align with practical chemistry encountered in early laboratory experience.

Writing Accurate Full Equations

To meet OCR expectations, students must ensure:

• Correct formulas for all reactants and products

• Balanced equations with respect to atoms and hydrogen gas

• Accurate salt naming corresponding to the acid used

• Inclusion of hydrogen gas as H₂, not simply “H”

These principles ensure precision when writing full equations for reactions between metals and acids.