Chemical changes: reactions of acids with metals, bases and carbonates

What you'll learn

This revision guide covers the chemical reactions of acids with metals, bases and carbonates — a fundamental topic in the AQA GCSE Chemistry specification. You'll learn to identify products, write balanced equations and understand the patterns that govern these important reactions. These concepts appear frequently in both Paper 1 and Paper 2, making them essential for exam success.

Key terms and definitions

Acid — a substance that produces hydrogen ions (H⁺) when dissolved in water; has a pH less than 7.

Base — a substance that neutralises acids to form a salt and water; metal oxides and metal hydroxides are common bases.

Alkali — a soluble base that produces hydroxide ions (OH⁻) when dissolved in water.

Salt — an ionic compound formed when the hydrogen ion in an acid is replaced by a metal ion or ammonium ion.

Neutralisation — a chemical reaction between an acid and a base that produces a salt and water.

Carbonate — a compound containing the carbonate ion (CO₃²⁻); reacts with acids to produce carbon dioxide, water and a salt.

Word equation — a chemical equation written using the names of reactants and products.

Balanced symbol equation — a chemical equation using formulae that shows equal numbers of each type of atom on both sides.

Core concepts

Reactions of acids with metals

When a reactive metal reacts with an acid, two products form: a salt and hydrogen gas. The general equation is:

acid + metal → salt + hydrogen

Only metals above hydrogen in the reactivity series react with dilute acids. These include:

Magnesium
Aluminium
Zinc
Iron

Unreactive metals such as copper, silver and gold do not react with dilute acids.

Example reactions:

hydrochloric acid + magnesium → magnesium chloride + hydrogen

2HCl(aq) + Mg(s) → MgCl₂(aq) + H₂(g)

sulfuric acid + zinc → zinc sulfate + hydrogen

H₂SO₄(aq) + Zn(s) → ZnSO₄(aq) + H₂(g)

nitric acid + iron → iron(II) nitrate + hydrogen

2HNO₃(aq) + Fe(s) → Fe(NO₃)₂(aq) + H₂(g)

Testing for hydrogen gas:

The hydrogen produced can be identified using the squeaky pop test. Insert a lit splint into a test tube containing the gas — hydrogen burns with a distinctive 'pop' sound.

Salt formation patterns:

The salt name depends on both the metal and the acid:

Hydrochloric acid → chloride salts
Sulfuric acid → sulfate salts
Nitric acid → nitrate salts

The first part of the salt name comes from the metal used.

Reactions of acids with bases (metal oxides and hydroxides)

Metal oxides and metal hydroxides are bases. When they react with acids, neutralisation occurs:

acid + base → salt + water

Metal oxide reactions:

sulfuric acid + copper oxide → copper sulfate + water

H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)

hydrochloric acid + zinc oxide → zinc chloride + water

2HCl(aq) + ZnO(s) → ZnCl₂(aq) + H₂O(l)

nitric acid + magnesium oxide → magnesium nitrate + water

2HNO₃(aq) + MgO(s) → Mg(NO₃)₂(aq) + H₂O(l)

Metal hydroxide reactions:

Metal hydroxides follow the same pattern. Soluble hydroxides (alkalis) can be reacted in solution:

hydrochloric acid + sodium hydroxide → sodium chloride + water

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

sulfuric acid + potassium hydroxide → potassium sulfate + water

H₂SO₄(aq) + 2KOH(aq) → K₂SO₄(aq) + 2H₂O(l)

nitric acid + calcium hydroxide → calcium nitrate + water

2HNO₃(aq) + Ca(OH)₂(aq) → Ca(NO₃)₂(aq) + 2H₂O(l)

Practical observations:

In metal oxide reactions, the solid oxide dissolves as neutralisation proceeds. The solution warms due to the exothermic nature of neutralisation. With insoluble bases, excess solid remains after the reaction is complete — this must be filtered off to obtain pure salt solution.

Reactions of acids with carbonates

Carbonates and hydrogencarbonates react with acids to produce three products: a salt, water and carbon dioxide gas.

acid + carbonate → salt + water + carbon dioxide

This reaction pattern applies to both metal carbonates and hydrogencarbonates.

Example reactions:

hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide

2HCl(aq) + CaCO₃(s) → CaCl₂(aq) + H₂O(l) + CO₂(g)

sulfuric acid + sodium carbonate → sodium sulfate + water + carbon dioxide

H₂SO₄(aq) + Na₂CO₃(s) → Na₂SO₄(aq) + H₂O(l) + CO₂(g)

nitric acid + copper carbonate → copper nitrate + water + carbon dioxide

2HNO₃(aq) + CuCO₃(s) → Cu(NO₃)₂(aq) + H₂O(l) + CO₂(g)

Testing for carbon dioxide:

Carbon dioxide turns limewater (calcium hydroxide solution) cloudy or milky. This is the standard test for CO₂ gas at GCSE level. The cloudiness is caused by the formation of insoluble calcium carbonate:

CO₂(g) + Ca(OH)₂(aq) → CaCO₃(s) + H₂O(l)

Observable signs:

Carbonate reactions produce effervescence (fizzing) as carbon dioxide gas is released. The solid carbonate gradually disappears as it reacts. These reactions are commonly used in Caribbean households when lime juice (citric acid) reacts with calcium carbonate in water, producing temporary hardness removal.

Ionic equations for neutralisation

The ionic equation for neutralisation reactions between acids and alkalis shows only the reacting ions:

H⁺(aq) + OH⁻(aq) → H₂O(l)

This equation represents the fundamental change occurring in any acid-alkali neutralisation. The spectator ions (those not involved in the reaction) are omitted.

For a complete ionic equation involving hydrochloric acid and sodium hydroxide:

H⁺(aq) + Cl⁻(aq) + Na⁺(aq) + OH⁻(aq) → Na⁺(aq) + Cl⁻(aq) + H₂O(l)

The sodium and chloride ions are spectator ions, leaving the simplified ionic equation above.

Making soluble salts

Acid reactions provide methods for preparing soluble salts in the laboratory:

Method 1: Excess insoluble base or carbonate

Add excess metal oxide, hydroxide or carbonate to warm dilute acid
Continue adding until no more dissolves and effervescence stops
Filter to remove excess solid
Heat the filtrate gently to evaporate some water
Leave to crystallise
Filter, wash and dry the crystals

Method 2: Titration (for soluble bases)

Use a pipette to measure a fixed volume of alkali into a conical flask
Add indicator
Add acid from a burette until the indicator changes colour
Record the volume of acid needed
Repeat without indicator using the recorded volumes
Evaporate the solution to obtain salt crystals

This method ensures no excess reactant contaminates the salt.

pH changes during neutralisation

During neutralisation of an acid with an alkali:

The pH starts low (acidic, around pH 1-2)
As alkali is added, pH gradually increases
At the equivalence point (complete neutralisation), pH = 7
Excess alkali raises pH above 7 (alkaline)

Universal indicator or pH meters can monitor these changes. The rapid pH change near the equivalence point makes this the most important region for titration endpoints.

Worked examples

Example 1:

Write a balanced symbol equation for the reaction between magnesium and sulfuric acid.

(3 marks)

Answer:

Word equation: sulfuric acid + magnesium → magnesium sulfate + hydrogen (1 mark)

Symbol equation: H₂SO₄ + Mg → MgSO₄ + H₂ (1 mark)

This equation is already balanced (1 mark for stating this OR for correctly balancing if unbalanced)

Mark scheme notes: Award marks for correct formulae and balancing. Common errors include incorrect salt formula.

Example 2:

A student adds copper carbonate powder to dilute hydrochloric acid.

(a) Write a word equation for this reaction. (1 mark)

(b) Describe what the student would observe. (2 marks)

Answers:

(a) hydrochloric acid + copper carbonate → copper chloride + water + carbon dioxide (1 mark)

OR any correct word equation with all five substances named correctly (1 mark)

(b) Effervescence/fizzing/bubbles produced (1 mark)

Green/blue solution forms as the black/green powder dissolves (1 mark)

OR any two valid observations from: gas produced, solid disappears, solution warms, colour change

Alternative marking: Award 1 mark for partially correct formulae, 2 marks for all correct formulae, 1 mark for correct balancing.

Example 3:

A student neutralises 25.0 cm³ of sulfuric acid with sodium hydroxide solution.

The equation for the reaction is:

H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)

(a) Name the salt produced. (1 mark)

(b) Write the ionic equation for this neutralisation reaction. (1 mark)

Answers:

(a) Sodium sulfate (1 mark)

Accept: Na₂SO₄

Do not accept: sodium sulfide, sulfate

(b) H⁺(aq) + OH⁻(aq) → H₂O(l) (1 mark)

Accept: 2H⁺(aq) + 2OH⁻(aq) → 2H₂O(l) OR H⁺ + OH⁻ → H₂O without state symbols

Do not accept: equations including spectator ions

Common mistakes and how to avoid them

Writing incorrect salt names: Remember the salt name combines the metal (or metal from the base) with the acid ending. Hydrochloric acid always gives chlorides, never chlorines or chlorates. Create a revision table matching acids to their salts.
Forgetting carbon dioxide in carbonate reactions: Carbonate reactions produce three products, not two. Always write salt + water + carbon dioxide. The mnemonic "Carbonates Create Carbon dioxide" helps.
Unbalanced equations: Count atoms carefully on both sides. Start by balancing metals, then non-metals, then hydrogen and oxygen last. Check your work by recounting. In 2HCl + Mg → MgCl₂ + H₂, there are 2H and 2Cl on each side.
Confusing bases and alkalis: All alkalis are bases, but not all bases are alkalis. Alkalis are soluble bases that dissolve in water. Sodium hydroxide is an alkali; copper oxide is a base but not an alkali (it's insoluble).
Incorrect formulae for polyatomic ions: Learn common ion formulae: sulfate is SO₄²⁻ (not SO₃²⁻), carbonate is CO₃²⁻, nitrate is NO₃⁻, hydroxide is OH⁻. Use brackets when needed: Ca(OH)₂ not CaOH₂.
Wrong tests for gases: Hydrogen gives a squeaky pop with a lit splint. Carbon dioxide turns limewater milky/cloudy. Don't confuse these tests or use non-standard descriptions like "white" for limewater.

Exam technique for "Chemical changes: reactions of acids with metals, bases and carbonates"

Command words matter: "State" requires a simple answer without explanation. "Describe" needs observable changes. "Explain" requires scientific reasoning linking cause and effect. "Write a balanced equation" requires correct formulae AND equal atoms on both sides — both earn separate marks.
Observations must be specific: Write "colourless gas" not just "gas", "blue solution forms" not "colour change", "effervescence" or "fizzing" not "reaction". Vague answers like "it reacts" or "bubbles appear" may not earn marks without specifying what type of bubbles or products.
Equation marks are split: Typically 2-3 marks for a balanced symbol equation: 1-2 marks for correct formulae, 1 mark for balancing. You can score partial marks even if not fully correct. Always attempt equations even if unsure — write the word equation first to guide your symbol equation.
Show your working in calculations: Even for simple tasks like balancing equations, write intermediate steps. If asked to name a salt, identify the metal and acid separately, then combine. This maximises partial credit opportunities and demonstrates your reasoning clearly to examiners.

Quick revision summary

Acids react with metals above hydrogen in the reactivity series to produce salts and hydrogen gas. Acids neutralise metal oxides and hydroxides (bases) to form salts and water only. Carbonates react with acids producing salts, water and carbon dioxide, identified by the limewater test. Salt names combine the metal with the acid type: chlorides from hydrochloric acid, sulfates from sulfuric acid, nitrates from nitric acid. Master these three reaction patterns, practice balancing equations, and learn the observable signs of each reaction type for exam success.