AQA A-Level Chemistry: Properties of Period 3 elements and their oxides (7405)

This guide teaches section 3.2.4 Properties of Period 3 elements and their oxides of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards. The whole section is marked A-level only in the specification, so it is not part of the AS course. It sits in inorganic chemistry, which the specification places on A-level Paper 1 (all of section 3.2), and A-level Paper 3 can test any content.

For quick recall, use the Period 3 oxides revision notes. To test yourself, use the Period 3 oxides practice questions. The course hub is AQA A-Level Chemistry, and the printable checklist lists every specification point. To find your weak spots first, try one of the free 10-minute diagnostics.

Background: periodicity (the elements Na to Ar), bonding and acids and bases.

What this unit covers

Spec point	What you must be able to do	Status
3.2.4	Describe the reactions of Na and Mg with water	A-level only
3.2.4	Describe the trends in the reactions of Na, Mg, Al, Si, P and S with oxygen, forming Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₂ and SO₃	A-level only
3.2.4	Explain the trend in melting point of the highest oxides of Na to S in terms of structure and bonding	A-level only
3.2.4	Describe the reactions of Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₂ and SO₃ with water and the pH of the solutions; explain the trend in terms of bonding	A-level only
3.2.4	Draw the structures of the acids and anions formed when P₄O₁₀, SO₂ and SO₃ react with water	A-level only
3.2.4	Write equations for reactions of the oxides of Na to S with given acids and bases	A-level only

The specification lists no required practical in section 3.2.4. It suggests you could burn elements in oxygen and test the pH of the resulting oxides (apparatus and techniques AT a, c and k; practical skill PS 2.2).

Reactions of sodium and magnesium with water

Sodium reacts vigorously with cold water. It floats, melts into a ball, fizzes and moves about. The solution is strongly alkaline (pH about 13 to 14).

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Magnesium reacts very slowly with cold water, giving a few bubbles of hydrogen. Magnesium hydroxide is only slightly soluble, so the solution is weakly alkaline (pH about 9 to 10).

Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)

Magnesium reacts much faster with steam. It burns with a bright white flame and forms magnesium oxide, not the hydroxide.

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

Why sodium is more reactive. In both reactions the metal atoms lose their outer electrons to form positive ions. A sodium atom loses one electron. A magnesium atom must lose two, and the second is removed from a positive ion, so much more energy is needed in total. Magnesium therefore reacts more slowly.

Reactions of the elements with oxygen

All six elements Na to S burn when heated in oxygen. Each forms its highest oxide, except sulfur, which first forms SO₂.

Element	Observation on burning	Equation	Oxidation state
Na	yellow-orange flame, white solid	4Na + O₂ → 2Na₂O	+1
Mg	bright white flame, white solid	2Mg + O₂ → 2MgO	+2
Al	powder burns with a bright white flame, white solid	4Al + 3O₂ → 2Al₂O₃	+3
Si	needs strong heating, white solid	Si + O₂ → SiO₂	+4
P	burns with a bright flame, dense white fumes	P₄ + 5O₂ → P₄O₁₀	+5
S	blue flame, colourless choking gas	S + O₂ → SO₂	+4

Two points to know:

• A piece of aluminium seems unreactive because a thin, tough layer of Al₂O₃ covers its surface and stops further reaction. Powdered aluminium has a large surface area and burns well.
• SO₂ is only converted to SO₃ by further oxidation, which needs a catalyst (vanadium(V) oxide) and heat: 2SO₂ + O₂ ⇌ 2SO₃. In SO₃ sulfur is +6.

The maximum oxidation state rises by one each time from Na (+1) to S (+6), because it equals the number of outer electrons that the atom can use in bonding.

Worked example 1

What mass of sodium oxide forms when 2.30 g of sodium burns completely in oxygen? (Ar: Na = 23.0, O = 16.0)

moles of Na  = 2.30 / 23.0 = 0.100 mol
4Na → 2Na₂O, so moles of Na₂O = 0.100 / 2 = 0.0500 mol
Mr of Na₂O   = 2(23.0) + 16.0 = 62.0
mass of Na₂O = 0.0500 × 62.0 = 3.10 g

Melting points of the highest oxides

The highest oxides are Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀ and SO₃. Their melting points rise from Na₂O to MgO, stay high for Al₂O₃ and SiO₂, then fall sharply for P₄O₁₀ and SO₃. MgO has the highest melting point of the set. In order: MgO > Al₂O₃ > SiO₂ > Na₂O > P₄O₁₀ > SO₃.

Oxide	Structure	Bonding broken on melting
Na₂O, MgO	giant ionic lattice	strong electrostatic attraction between oppositely charged ions
Al₂O₃	giant ionic lattice with some covalent character	strong ionic attraction
SiO₂	macromolecular (giant covalent)	many strong covalent bonds
P₄O₁₀, SO₃ (and SO₂)	simple molecular	weak intermolecular forces only

Na₂O and MgO. Both are giant ionic lattices. Mg²⁺ has a higher charge than Na⁺ and is smaller, so the attraction between Mg²⁺ and O²⁻ ions is stronger. More energy is needed to overcome it, so MgO melts at a higher temperature.

Al₂O₃. Al³⁺ is small and highly charged. It distorts (polarises) the electron cloud of nearby O²⁻ ions, so the bonding has some covalent character. Its melting point is high but lower than that of MgO.

SiO₂. Each silicon atom is covalently bonded to four oxygen atoms and each oxygen to two silicon atoms in a giant structure. Melting breaks many strong covalent bonds, so the melting point is high.

P₄O₁₀, SO₃ and SO₂. These are simple molecules. Melting overcomes only the weak forces between molecules, not the covalent bonds inside them, so melting points are low. P₄O₁₀ molecules are much larger, with more electrons (140 per molecule against 40 in SO₃), so the van der Waals forces between them are stronger and P₄O₁₀ melts at a higher temperature than SO₃. SO₂ melts lower still.

Worked example 2

Explain why MgO has a much higher melting point than P₄O₁₀.

1. MgO is a giant ionic lattice. Melting it overcomes strong attractions between Mg²⁺ and O²⁻ ions, which needs a lot of energy.
2. P₄O₁₀ is simple molecular. Melting overcomes only weak van der Waals forces between molecules; no covalent bonds break.
3. Much less energy is needed, so P₄O₁₀ melts at a much lower temperature.

Reactions of the oxides with water

Oxide	Reaction with water	pH of solution (approx.)	Nature of oxide
Na₂O	Na₂O + H₂O → 2NaOH	13 to 14	basic
MgO	MgO + H₂O → Mg(OH)₂	9 to 10	basic
Al₂O₃	insoluble, no reaction	7	amphoteric
SiO₂	insoluble, no reaction	7	acidic
P₄O₁₀	P₄O₁₀ + 6H₂O → 4H₃PO₄	0 to 1	acidic
SO₂	SO₂ + H₂O → H₂SO₃	2 to 3	acidic
SO₃	SO₃ + H₂O → H₂SO₄	0 to 1	acidic

Explaining the trend by bonding.

• Ionic oxides (Na₂O, MgO) contain oxide ions. The oxide ion takes a proton from water to form hydroxide ions: O²⁻ + H₂O → 2OH⁻. The solution is alkaline. Mg(OH)₂ is only slightly soluble, so fewer OH⁻ ions are in solution and the pH is lower than for NaOH.
• Al₂O₃ (ionic with covalent character) and SiO₂ (macromolecular) have very strong bonding in giant structures and do not dissolve, so the water stays at pH 7.
• Simple molecular covalent oxides (P₄O₁₀, SO₂, SO₃) react with water to form acids, which release H⁺ ions. H₂SO₃ is a weak acid that only partly dissociates (H₂SO₃ ⇌ H⁺ + HSO₃⁻), so its pH is higher than that of H₂SO₄.

So across Period 3 the oxides change from basic (ionic) to amphoteric (Al₂O₃) to acidic (covalent).

Structures of the acids and anions

In the anions the negative charge is shown on single-bonded oxygen atoms.

Phosphoric(V) acid, H₃PO₄         Phosphate ion, PO₄³⁻
          O                                 O
          ‖                                 ‖
   H–O – P – O–H                     ⁻O – P – O⁻
          |                                 |
          O–H                               O⁻

Sulfuric(IV) acid, H₂SO₃          Sulfate(IV) ion, SO₃²⁻
          O                                 O
          ‖                                 ‖
   H–O – S – O–H   (lone pair on S)   ⁻O – S – O⁻   (lone pair on S)

Sulfuric(VI) acid, H₂SO₄          Sulfate(VI) ion, SO₄²⁻
          O                                 O
          ‖                                 ‖
   H–O – S – O–H                     ⁻O – S – O⁻
          ‖                                 ‖
          O                                 O

Partly neutralised ions keep some O–H groups: H₂PO₄⁻ and HPO₄²⁻ from phosphoric(V) acid, HSO₃⁻ from sulfuric(IV) acid and HSO₄⁻ from sulfuric(VI) acid. Phosphorus and sulfur(VI) have four bonding regions, so H₃PO₄, PO₄³⁻, H₂SO₄ and SO₄²⁻ are tetrahedral around the central atom. In H₂SO₃ and SO₃²⁻ sulfur has three bonding regions and a lone pair, so the shape is pyramidal.

Sulfuric(IV) acid is also called sulfurous acid, and the sulfate(IV) ion is the sulfite ion.

Reactions of the oxides with acids and bases

Basic oxides react with acids to form a salt and water.

Na₂O + 2HCl → 2NaCl + H₂O          MgO + H₂SO₄ → MgSO₄ + H₂O
ionic: O²⁻ + 2H⁺ → H₂O

Amphoteric Al₂O₃ reacts with both acids and bases.

with acid:  Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
            Al₂O₃ + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂O
with base:  Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄
ionic:      Al₂O₃ + 2OH⁻ + 3H₂O → 2Al(OH)₄⁻

Acidic oxides react with bases. SiO₂ does not react with water, but it reacts with hot, concentrated sodium hydroxide and with basic oxides such as CaO.

SiO₂ + 2NaOH → Na₂SiO₃ + H₂O        SiO₂ + CaO → CaSiO₃
P₄O₁₀ + 12NaOH → 4Na₃PO₄ + 6H₂O
SO₂ + 2NaOH → Na₂SO₃ + H₂O          SO₂ + CaO → CaSO₃
SO₃ + 2NaOH → Na₂SO₄ + H₂O          SO₃ + MgO → MgSO₄

Worked example 3

Calculate the mass of magnesium oxide that reacts exactly with 25.0 cm³ of 2.00 mol dm⁻³ hydrochloric acid. (Ar: Mg = 24.3, O = 16.0)

MgO + 2HCl → MgCl₂ + H₂O
moles of HCl = 25.0 / 1000 × 2.00 = 0.0500 mol
moles of MgO = 0.0500 / 2 = 0.0250 mol
Mr of MgO    = 24.3 + 16.0 = 40.3
mass of MgO  = 0.0250 × 40.3 = 1.01 g (3 s.f.)

Practical: burning elements and testing the oxides

A typical method:

1. Heat a small piece of the element on a combustion spoon and lower it into a gas jar of oxygen.
2. When burning stops, add a little distilled water and shake.
3. Measure the pH with universal indicator (compare with a pH chart) or a pH meter (AT c).
4. Record the results in a table of element, observation, oxide formed and pH (PS 2.2).

Handle corrosive and toxic substances with care (AT k). Burn phosphorus and sulfur in a fume cupboard, because their oxides are toxic and irritant.

Common errors

• Writing Mg + H₂O → MgO + H₂ for cold water. That equation is for steam; cold water gives Mg(OH)₂.
• Giving pH 14 for MgO. Mg(OH)₂ is only slightly soluble, so the pH is about 9 to 10.
• Saying Al₂O₃ or SiO₂ “dissolve to form a neutral solution”. They are insoluble; the water stays neutral.
• Saying covalent bonds break when P₄O₁₀ or SO₃ melts. Only intermolecular forces are overcome.
• Giving SiO₂ “strong intermolecular forces”; it has covalent bonds throughout.
• Writing SO₂ + H₂O → H₂SO₄. Sulfur stays +4, so the product is H₂SO₃.
• Forgetting that Al₂O₃ reacts with bases as well as acids.

Official syllabus

AQA, AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards: section 3.2.4 Properties of Period 3 elements and their oxides (A-level only).