Structure, Bonding, and Melting Points of Period 3 Oxides Revision Notes | A Level AQA Chemistry

1. Why Period 3 Oxides Are Worth Understanding

Across Period 3, eight elements form oxides with wildly different melting points — from sodium oxide melting above 1000 °C to sulphur trioxide melting at around 17 °C. That is not random. Every single difference traces back to one question: what type of structure does the oxide adopt?

This topic sits at the intersection of bonding, structure, and reactivity — three themes that run through the entire A-level course. Master it here and you will find the same reasoning patterns appearing in questions on lattice enthalpies, acid-base chemistry, and industrial processes.

2. The Oxides of Period 3: An Overview

The Period 3 elements sodium through to sulphur each form a highest oxide. Moving left to right: $\text{Na}_2\text{O}$, $\text{MgO}$, $\text{Al}_2\text{O}_3$, $\text{SiO}_2$, $\text{P}_4\text{O}_{10}$, $\text{SO}_3$. Chlorine also forms oxides but the AQA specification focuses on Na–S for this trend, so that is our range.

The key insight is that as you move from left to right across Period 3, the bonding character of the oxides shifts from ionic to covalent, and the structures shift from giant lattices to simple molecules. That shift drives everything else — melting points, acid-base behaviour, and solubility.

3. Melting Points: The Structural Explanation

Here is where the lightbulb moment lives. Melting point tells you how much energy is needed to break the forces holding a structure together. The type of force — and how many of them there are — determines the melting point.

Think of it in three tiers. Giant ionic lattices (Na, Mg, Al oxides) have strong electrostatic attractions between ions throughout the whole structure — high melting points. Giant covalent structures (Si oxide) have strong covalent bonds throughout — very high melting point. Simple molecular structures (P and S oxides) have only weak intermolecular forces between molecules (van der Waals forces, and additionally dipole-dipole forces for polar molecules such as $\text{SO}_3$) — low melting points.

Notice the pattern in the chart: melting points are high for the ionic oxides on the left, peak at MgO (2852 °C), then decrease slightly for $\text{Al}_2\text{O}_3$ (2072 °C), remain high for $\text{SiO}_2$, then drop sharply for $\text{P}_4\text{O}_{10}$ and $\text{SO}_3$. That sharp drop at silicon → phosphorus marks the transition from giant to simple molecular structure.

Now let us work through each oxide in turn, because the examiner will expect you to justify the melting point of a specific oxide — not just quote the general trend.

3a. Sodium Oxide and Magnesium Oxide — Giant Ionic Lattices

Both $\text{Na}_2\text{O}$ and $\text{MgO}$ adopt giant ionic lattice structures. To melt them, you must supply enough energy to overcome the electrostatic attractions between all the oppositely charged ions in the lattice — a large energy requirement, hence high melting points.

MgO has a higher melting point than $\text{Na}_2\text{O}$ (2852 °C vs 1132 °C). The reason is charge density. $\text{Mg}^{2+}$ carries a 2+ charge and is smaller than $\text{Na}^+$, so the electrostatic attraction between $\text{Mg}^{2+}$ and $\text{O}^{2-}$ is stronger than between $\text{Na}^+$ and $\text{O}^{2-}$. More energy is needed to break the MgO lattice.

3b. Aluminium Oxide — Still a Giant Lattice, but with a Twist

$\text{Al}_2\text{O}_3$ is also a giant ionic lattice with a high melting point (2072 °C). The $\text{Al}^{3+}$ ion carries a 3+ charge, which you might expect to give an even higher melting point than MgO. In practice, the melting point is slightly lower than MgO, partly because the bonding in $\text{Al}_2\text{O}_3$ has some covalent character — the high charge density of $\text{Al}^{3+}$ polarises the electron cloud of $\text{O}^{2-}$, distorting the purely ionic model. The melting point remains very high, however, because the lattice is still giant and the forces are still strong.

3c. Silicon Dioxide — Giant Covalent Structure

$\text{SiO}_2$ does not consist of discrete $\text{SiO}_2$ molecules. Each silicon atom is covalently bonded to four oxygen atoms in a continuous three-dimensional network — the same structural principle as diamond. To melt $\text{SiO}_2$, you must break many strong Si–O covalent bonds throughout the entire lattice, requiring a very large amount of energy. The melting point is 1710 °C.

3d. Phosphorus(V) Oxide and Sulphur Trioxide — Simple Molecular Structures

$\text{P}_4\text{O}_{10}$ and $\text{SO}_3$ are both simple molecular. The covalent bonds within each molecule are strong, but the forces between molecules are only weak van der Waals forces. Melting requires overcoming only these weak intermolecular forces, not the covalent bonds themselves. The result is dramatically lower melting points: $\text{P}_4\text{O}_{10}$ melts at around 340 °C and $\text{SO}_3$ at only 17 °C.

$\text{SO}_3$ has a lower melting point than $\text{P}_4\text{O}_{10}$ because its molecules are smaller and have fewer electrons, so the van der Waals forces between $\text{SO}_3$ molecules are weaker than those between the larger $\text{P}_4\text{O}_{10}$ molecules.

4. Acid-Base Character of Period 3 Oxides

Now that the structures are clear, we can connect them to reactivity. The acid-base character of an oxide follows a clear pattern: metallic oxides on the left are basic, non-metallic oxides on the right are acidic, and aluminium oxide in the middle is amphoteric.

The underlying reason is electronegativity. In ionic oxides, the $\text{O}^{2-}$ ion is available to accept protons (acting as a base). In covalent oxides, the element–oxygen bond is polarised so that the oxide dissolves in water to release $\text{H}^+$ ions (acting as an acid). The more electronegative the Period 3 element, the more acidic its oxide.

4a. Basic Oxides: Na₂O and MgO

$\text{Na}_2\text{O}$ and $\text{MgO}$ are basic oxides. They react with acids to form a salt and water. $\text{Na}_2\text{O}$ also reacts vigorously with water to form sodium hydroxide solution.

• $\text{Na}_2\text{O}(s) + \text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq)$
• $\text{Na}_2\text{O}(s) + 2\text{HCl}(aq) \rightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l)$
• $\text{MgO}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2\text{O}(l)$
• $\text{MgO}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{MgSO}_4(aq) + \text{H}_2\text{O}(l)$

MgO reacts only slowly with water at room temperature, so its reaction with water is not listed as a key equation in the same way as $\text{Na}_2\text{O}$. Its basic character is demonstrated primarily through its reactions with acids.

4b. Amphoteric Oxide: Al₂O₃

Aluminium oxide sits at the boundary between metallic and non-metallic character in Period 3. It is amphoteric — it reacts with both acids and bases. This is the behaviour that examiners test most frequently for this topic, so make sure you can write both equations.

• Reaction with acid: $\text{Al}_2\text{O}_3(s) + 6\text{HCl}(aq) \rightarrow 2\text{AlCl}_3(aq) + 3\text{H}_2\text{O}(l)$
• Reaction with base: $\text{Al}_2\text{O}_3(s) + 2\text{NaOH}(aq) + 3\text{H}_2\text{O}(l) \rightarrow 2\text{NaAl(OH)}_4(aq)$

Al2O3 acts as an acid when it reacts with a base such as NaOH. In this reaction the Al3+ centre accepts hydroxide ions to form the tetrahydroxoaluminate(III) ion, Al(OH)4-.

4c. Acidic Oxides: SiO₂, P₄O₁₀, and SO₃

The non-metal oxides on the right of Period 3 are acidic. They react with bases (alkalis) but not with acids. $\text{SiO}_2$ is only weakly acidic and does not dissolve in water at room temperature, but it does react with hot concentrated sodium hydroxide solution.

$\text{P}_4\text{O}_{10}$ and $\text{SO}_3$ react readily with water to form acids — phosphoric(V) acid and sulphuric acid respectively. These dissolved acids then react with bases in the usual way.

• $\text{SiO}_2(s) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{SiO}_3(aq) + \text{H}_2\text{O}(l)$
• $\text{P}_4\text{O}_{10}(s) + 6\text{H}_2\text{O}(l) \rightarrow 4\text{H}_3\text{PO}_4(aq)$
• $\text{P}_4\text{O}_{10}(s) + 12\text{NaOH}(aq) \rightarrow 4\text{Na}_3\text{PO}_4(aq) + 6\text{H}_2\text{O}(l)$
• $\text{SO}_3(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_4(aq)$
• $\text{SO}_3(g) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{SO}_4(aq) + \text{H}_2\text{O}(l)$

Notice that none of the acidic oxides react with hydrochloric acid or sulphuric acid — they only react with bases. This is the defining test of an acidic oxide, and it is the mirror image of the basic oxides which react with acids but not with bases.

5. Putting the Trend Together: Structure, Bonding, and Melting Point

You now have all the pieces. The trend in melting points across Period 3 oxides is not a smooth decrease — it has a distinctive shape that reflects the change in structure type. The step-by-step walkthrough below models how to construct a full exam-quality explanation of this trend.

Worked example: Explaining the Melting Point Trend Across Period 3 Oxides

Explain the trend in the melting points of the highest oxides of the elements Na to S. [6 marks]

1. Na₂O, MgO, Al₂O₃ → giant ionic lattices SiO₂ → giant covalent lattice P₄O₁₀, SO₃ → simple molecular structures — Categorising by structure type is the essential first move. Every melting point explanation must start here.
2. Na₂O and MgO: strong electrostatic attractions between oppositely charged ions throughout the giant lattice require large amounts of energy to overcome → high melting points. — Use the phrase 'electrostatic attractions' rather than just 'ionic bonds' — it is more precise and scores the mark.
3. Mg²⁺ has a higher charge and smaller ionic radius than Na⁺, so the electrostatic attraction between Mg²⁺ and O²⁻ is stronger than between Na⁺ and O²⁻ → MgO has a higher melting point than Na₂O. — Charge density is the key concept here. Higher charge and smaller size = stronger attraction = higher melting point.
4. SiO₂ has a giant covalent structure. Each Si is bonded to four O atoms by strong covalent bonds throughout the lattice. Melting requires breaking many strong covalent bonds → very high energy required → high melting point (1710 °C). — Emphasise 'throughout the lattice' — this distinguishes giant covalent from simple molecular.
5. P₄O₁₀ and SO₃ are simple molecular. Strong covalent bonds exist within each molecule, but only weak van der Waals forces act between molecules. Melting requires overcoming only these weak intermolecular forces → low melting points. — The trap here is saying the covalent bonds are broken on melting — they are not. Only the intermolecular forces are overcome.
6. SO₃ molecules are smaller than P₄O₁₀ molecules and have fewer electrons, so the van der Waals forces between SO₃ molecules are weaker → lower melting point than P₄O₁₀. — Molecular size and electron count determine the strength of van der Waals forces between simple molecules.

Answer: Melting points are high for Na₂O, MgO, and Al₂O₃ (giant ionic lattices with strong electrostatic attractions), very high for SiO₂ (giant covalent lattice requiring many strong covalent bonds to be broken), then drop sharply for P₄O₁₀ and SO₃ (simple molecular structures with only weak van der Waals forces between molecules). SO₃ has the lowest melting point because its molecules are smallest with fewest electrons, giving the weakest van der Waals forces.

6. Writing Equations: A Systematic Approach

The examiner can ask you to write equations for any Period 3 oxide reacting with a given acid or base. There is a reliable method for constructing these equations without memorising every single one.

Worked example: Writing an Equation for a Period 3 Oxide Reacting with an Acid or Base

Write a balanced equation, with state symbols, for the reaction of aluminium oxide with dilute sulphuric acid.

1. Al₂O₃ is amphoteric. Reacting with an acid → Al₂O₃ acts as a base. — Knowing the acid-base character tells you what products to expect.
2. Base + acid → salt + water\nAl₂O₃ + H₂SO₄ → Al₂(SO₄)₃ + H₂O — The salt is aluminium sulphate. Check the formula: Al³⁺ and SO₄²⁻ → Al₂(SO₄)₃.
3. Al₂O₃ + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂O\nLeft: 2 Al, 3 S, 15 O (3 from Al₂O₃ + 12 from 3H₂SO₄), 6 H\nRight: 2 Al, 3 S, 15 O (12 from Al₂(SO₄)₃ + 3 from 3H₂O), 6 H ✓ — Count every atom on both sides. Oxygen is the trickiest — count it last.
4. Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l) — Al₂O₃ is a solid. The acid is aqueous. The salt dissolves in water, so it is aqueous. Water is liquid.

Answer: $\text{Al}_2\text{O}_3(s) + 3\text{H}_2\text{SO}_4(aq) \rightarrow \text{Al}_2(\text{SO}_4)_3(aq) + 3\text{H}_2\text{O}(l)$

7. Exam Technique and Common Mistakes

This topic rewards precise language. Examiners are looking for specific terms — 'electrostatic attractions', 'giant covalent lattice', 'van der Waals forces between molecules' — and will not award marks for vague phrasing like 'strong bonds' without qualification.

The most common error in melting point questions is saying that covalent bonds are broken when simple molecular substances melt. They are not. The covalent bonds within $\text{SO}_3$ or $\text{P}_4\text{O}_{10}$ molecules remain intact during melting — only the intermolecular forces between molecules are overcome. Save 'breaking covalent bonds' for giant covalent structures like $\text{SiO}_2$.