UNIT 2 CHEMISTRY · PEARSON AAQ · C1.5
Acid–Base Behaviour of Period 3 Compounds
FREE CHEMISTRY REVISION · TOPIC 27
Explain chemical behaviour through reactions and evidence.
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · C1.5
What you will learn
Classify the specified oxides Na₂O, MgO, Al₂O₃, P₄O₁₀ and SO₂; hydroxides NaOH, Mg(OH)₂ and Al(OH)₃; and SiO₂, H₃PO₄, H₂SO₄, HCl, HClO and HClO₄. Explain their reactions with acids or bases using balanced equations. Conditions and the amount of reagent matter.
1. Acids, bases and amphoteric substances
A Brønsted–Lowry acid donates a proton; a base accepts one. In aqueous reactions, acids increase the concentration of H₃O⁺ (often shortened to H⁺). An alkali is a soluble base that produces OH⁻ in water. A base can be insoluble: MgO neutralises acid even though it does not form a highly soluble alkaline solution.
H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l)
A basic oxide reacts with acid to form salt and water. An acidic oxide reacts with a base to form a salt, usually with water. An amphoteric oxide or hydroxide reacts with both acids and strong bases. Amphoteric does not mean neutral, and it does not mean equal acidity and basicity in every condition.
Acid–base reactions need not involve a change in oxidation number. In neutralisation, proton transfer or reaction of oxide with H⁺ can occur without electron transfer.
2. Basic metal oxides and hydroxides
Na₂O and MgO are basic oxides. Oxide ions accept protons: O²⁻ + 2H⁺ → H₂O. Sodium oxide reacts readily with water, producing a strongly alkaline sodium hydroxide solution. Magnesium oxide reacts more slowly with water; the magnesium hydroxide formed is sparingly soluble. Its saturated solution is alkaline, but there is much less dissolved OH⁻ than in a comparable concentrated NaOH solution.
Na₂O(s) + H₂O(l) → 2NaOH(aq)
MgO(s) + H₂O(l) → Mg(OH)₂(s)
NaOH is a strong, soluble base. Mg(OH)₂ is a base with low solubility; do not explain its limited dissolved OH⁻ simply by calling it a weak base. Both neutralise acids. A white MgO or Mg(OH)₂ solid dissolves as sufficient acid reacts with it; the exact rate depends on conditions.
Na₂O(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l)
MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)
Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)
3. Aluminium: amphoteric behaviour
Al₂O₃ and Al(OH)₃ react with acids and strong alkalis. They are not appreciably soluble in pure water. Aluminium oxide is not classified by a hypothetical pH of the dry solid: examine its chemical reactions. Acid dissolves the oxide or hydroxide to give hydrated aluminium ions, represented as Al³⁺(aq) in these simplified equations.
Al₂O₃(s) + 6H⁺(aq) → 2Al³⁺(aq) + 3H₂O(l)
Al(OH)₃(s) + 3H⁺(aq) → Al³⁺(aq) + 3H₂O(l)
In excess aqueous strong alkali, soluble tetrahydroxoaluminate ions form. White Al(OH)₃ dissolves in excess NaOH. Aluminium oxide also reacts with strong alkali; heating may be needed for a noticeable rate. The aqueous representation below is appropriate for these conditions.
Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq)
Al₂O₃(s) + 2OH⁻(aq) + 3H₂O(l) → 2[Al(OH)₄]⁻(aq)
Aqueous [Al(OH)₄]⁻ and aluminate products written for hot concentrated or molten conditions are alternative representations for different conditions. Do not mix them within one unbalanced equation.
4. Acidic non-metal oxides, including SiO₂
P₄O₁₀ and SO₂ are acidic oxides. Phosphorus(V) oxide reacts vigorously with water to produce phosphoric acid when water is in excess. Sulfur dioxide dissolves and establishes equilibria that make the water acidic. H₂SO₃ is a useful conventional overall representation for sulfurous acid in solution; most dissolved sulfur dioxide is not present as isolated H₂SO₃ molecules.
P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)
SO₂(aq) + 2H₂O(l) ⇌ H₃O⁺(aq) + HSO₃⁻(aq)
With sufficient NaOH, these acidic oxides form salts. The degree of neutralisation depends on the amount of base: the equations below use enough base for phosphate and sulfite products.
P₄O₁₀(s) + 12NaOH(aq) → 4Na₃PO₄(aq) + 6H₂O(l)
SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l)
SiO₂ is an acidic oxide with a giant covalent structure. It does not react appreciably with water and does not make an acidic solution simply by being stirred into it. It reacts with hot concentrated NaOH or molten alkali to form silicate; the equation shown uses a simplified metasilicate product. Its slow reaction and insolubility do not make it amphoteric or basic.
SiO₂(s) + 2NaOH(aq) → Na₂SiO₃(aq) + H₂O(l)
Hot concentrated aqueous alkali; silicate speciation depends on conditions.
5. The specified acids
Strength describes the extent of ionisation in water; concentration describes how much acid is present per unit volume. These are different properties. pH depends on concentration, ionisation and conditions; a formula alone does not give one exact pH. Monoprotic acids can donate one proton per molecule, diprotic two and triprotic three in successive stages.
| Acid | Aqueous behaviour | Complete neutralisation with NaOH |
|---|---|---|
| H₃PO₄, phosphoric acid | Weak, triprotic; successive ionisations become less favourable | H₃PO₄ + 3NaOH → Na₃PO₄ + 3H₂O |
| H₂SO₄, sulfuric acid | Diprotic; first ionisation essentially complete in dilute solution, second is an equilibrium | H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O |
| HCl, hydrochloric acid | Strong, monoprotic in water | HCl + NaOH → NaCl + H₂O |
| HClO, chloric(I) / hypochlorous acid | Weak, monoprotic | HClO + NaOH → NaClO + H₂O |
| HClO₄, chloric(VII) / perchloric acid | Strong, monoprotic in water | HClO₄ + NaOH → NaClO₄ + H₂O |
The table lists five formulas: H₃PO₄, H₂SO₄, HCl, HClO and HClO₄. Together with the separately discussed acidic oxide SiO₂, these complete Pearson’s additional named compounds. All acids and soluble salts in the table are aqueous; water is liquid. Full neutralisation requires the stated molar ratio. Less base gives acid salts such as NaH₂PO₄ or NaHSO₄.
HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)
HClO(aq) + H₂O(l) ⇌ H₃O⁺(aq) + ClO⁻(aq)
H₃PO₄(aq) + H₂O(l) ⇌ H₃O⁺(aq) + H₂PO₄⁻(aq)
Use a reversible arrow for weak-acid ionisation. HClO and HClO₄ are different acids with different strengths and different salts; neither is interchangeable with HCl. The oxidation number of chlorine also differs, but neutralisation itself need not change it.
6. Compare the pattern across Period 3
| Substances | Classification | Evidence |
|---|---|---|
| Na₂O, MgO; NaOH, Mg(OH)₂ | Basic | React with acid to form salt and water |
| Al₂O₃, Al(OH)₃ | Amphoteric | React with acid and excess strong alkali |
| SiO₂, P₄O₁₀, SO₂ | Acidic oxides | React with bases; water reaction is not required |
| H₃PO₄, H₂SO₄, HCl, HClO, HClO₄ | Acids | Donate protons and neutralise bases |
The broad oxide pattern is basic → amphoteric → acidic as the bonding changes from predominantly ionic metal oxides to covalent non-metal oxides. It is a qualitative acid–base pattern, not a claim that water solubility or pH changes smoothly at every step.
7. Four worked examples
Example 1 — magnesium oxide and acid
Task: Write a reaction demonstrating that MgO is basic. Choose HCl. Mg²⁺ needs two Cl⁻, so the salt is MgCl₂. One oxide oxygen forms one water molecule and needs two H atoms.
MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
Check Mg 1, O 1, H 2 and Cl 2 on each side. Dissolution as acid is added is evidence of reaction; no hydrogen gas is expected from this oxide–acid neutralisation.
Example 2 — prove amphoteric behaviour
Task: Explain why a white hydroxide dissolving in both acid and excess NaOH may be Al(OH)₃. In acid it accepts H⁺ to form water; in excess OH⁻ it forms a soluble complex. Use two fresh portions.
Al(OH)₃ + 3H⁺ → Al³⁺ + 3H₂O
Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻
The first has charge +3 on both sides; the second has charge −1 on both sides. Both atoms and charge balance.
Example 3 — full versus partial neutralisation
Task: How much NaOH fully neutralises 1 mol H₃PO₄? Three acidic hydrogens can be replaced in successive stages, so 3 mol NaOH is required.
H₃PO₄ + 3NaOH → Na₃PO₄ + 3H₂O
With only 1 mol NaOH per mole of acid, use H₃PO₄ + NaOH → NaH₂PO₄ + H₂O instead. “Weak” does not mean only one proton can ever be neutralised.
Example 4 — insoluble but acidic
Task: A student says SiO₂ cannot be acidic because it does not lower the pH of water. Acidic oxide classification is established by reaction with base, not by assuming it dissolves in water. Hot concentrated alkali gives silicate.
SiO₂ + 2NaOH → Na₂SiO₃ + H₂O
Count Si 1, Na 2, O 4 and H 2 on each side. The conditions explain why no rapid reaction is seen in cold water.
Common misconceptions
- “Every basic substance is an alkali.” Alkalis are soluble bases; insoluble oxides can still neutralise acids.
- “Every acidic oxide reacts with water.” SiO₂ reacts with strong base under suitable conditions but not appreciably with water.
- “Amphoteric means pH 7.” It describes reactions with both acids and bases.
- “Weak acid means dilute acid.” Strength is extent of ionisation, concentration is amount per volume.
- “HClO and HClO₄ form the same salt.” Their anions are ClO⁻ and ClO₄⁻ respectively.
Exam tips
- Name the compound and provide a reaction that demonstrates the classification.
- State excess strong alkali for aluminium hydroxide and suitable heating for the SiO₂–alkali reaction.
- Balance atoms and charge; never alter formula subscripts to repair an equation.
- Identify whether neutralisation is partial or complete before choosing phosphate or sulfate salts.
- Avoid assigning one universal pH to a solid or to a solution of unspecified concentration.
Check your understanding
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