Period 3 Elements and Their Oxides

Period 3 (Na, Mg, Al, Si, P, S, Cl) is the textbook laboratory for watching bonding type change in real time. The metals on the left lose electrons to oxygen to give ionic, giant-lattice oxides; the non-metals on the right share electrons to give molecular or giant covalent oxides; aluminium sits awkwardly between the two camps. In this lesson we systematise the reactions of each element with oxygen and (where it occurs) with water, classify the resulting oxide by structure and bonding, and then watch the pH of the aqueous solution swing from ≈14 (Na₂O) through ≈7 (Al₂O₃, insoluble) down to ≈1 (SO₃, Cl₂O₇). The same logic governs the Period 3 chlorides, where the transition from ionic NaCl to covalent SiCl₄ and PCl₅ is read off by their behaviour with water (dissolution → partial hydrolysis → complete hydrolysis). Mastering this lesson means being able to predict, justify, and write balanced equations for any combination of these oxides or chlorides with water, acid, or base.

Spec mapping (AQA 7405): This lesson maps to §3.2.4 (Properties of Period 3 elements and their oxides). It is the third lesson of the Inorganic course and builds directly on L0 (periodicity — §3.2.1), L1 (Group 2, the alkaline earth metals — §3.2.2), and L2 (Group 7, the halogens — §3.2.3). The bonding-type underpinning of the oxide structures (ionic vs covalent network vs molecular covalent) is developed in §3.1.3 lesson 0 of the Physical/Bonding course, and the acid–base reactivity logic is reinforced in §3.1.12 (acids and bases) at A2. Refer to the official AQA specification document for the exact wording of each section.

Assessment objectives: Recall of the reactions of Period 3 elements with O₂ and with H₂O, the structures and bonding types of the oxides, and the equations for oxide reactions with water, acid, and base, are AO1 items. Writing correctly balanced equations for the formation, hydrolysis, and acid–base reactions of any specified Period 3 oxide or chloride is AO2 and features on every Paper 1. Rationalising the trend in oxide pH across the period from electronic structure, polarising power, and bonding-type arguments — and predicting products and pH for unfamiliar oxides or for chloride hydrolysis — is AO3 and dominates the high-tariff (5–6 mark) extended-prose questions.

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Reactions of Period 3 Elements with Oxygen

The Period 3 elements all combine with oxygen, but vigour, temperature required, and stoichiometry vary systematically. Vigour drops from sodium (room-temperature self-ignition once initiated) to silicon (only at >1000 °C); chlorine alone does not combine directly with oxygen under normal conditions (the chlorine oxides are made by indirect routes).

Sodium (Na)

4Na(s) + O₂(g) → 2Na₂O(s)

• Burns with a bright yellow-orange flame.
• Produces a white solid, sodium oxide.
• In excess oxygen sodium also forms the peroxide Na₂O₂ (pale yellow) by 2Na + O₂ → Na₂O₂. AQA expects the monoxide Na₂O as the principal product for spec purposes.

Magnesium (Mg)

2Mg(s) + O₂(g) → 2MgO(s)

• Ignites and burns with a brilliant, blinding white flame.
• Produces a fine white powder, magnesium oxide.
• Very exothermic (ΔH°_f ≈ −602 kJ mol⁻¹). The white smoke seen in school demonstrations is fine MgO aerosol.

Aluminium (Al)

4Al(s) + 3O₂(g) → 2Al₂O₃(s)

• Powdered Al burns with a brilliant white flame.
• Produces a white solid, aluminium oxide.
• Bulk aluminium is protected by a tough, thin, self-healing oxide layer (a few nm) that prevents further reaction at room temperature. This passivation is why aluminium appears unreactive despite sitting high in the reactivity series — a frequent A-Level synoptic question.

Silicon (Si)

Si(s) + O₂(g) → SiO₂(s)

• Reacts only at very high temperatures (>1000 °C).
• Produces silicon dioxide (silica), a giant covalent solid — the main component of sand, quartz, and most rocks.
• SiO₂ has a very high melting point (1713 °C) reflecting the giant covalent network.

Phosphorus (P)

P₄(s) + 5O₂(g) → P₄O₁₀(s)

• White phosphorus burns spontaneously in air with a bright white flame (hence its storage under water).
• Produces phosphorus(V) oxide, P₄O₁₀, a white deliquescent solid.
• In limited oxygen the lower oxide P₄O₆ forms: P₄(s) + 3O₂(g) → P₄O₆(s).

Sulfur (S)

S(s) + O₂(g) → SO₂(g)

• Burns with a characteristic pale blue flame.
• Produces sulfur dioxide (SO₂), a colourless, pungent, choking gas.
• SO₂ can be further oxidised to SO₃ in the Contact process over a V₂O₅ catalyst at ≈450 °C: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g). Under normal combustion conditions only ≈ 5–10% of the SO₂ produced is converted to SO₃ without catalysis.

Chlorine (Cl)

Cl₂ does not react directly with O₂ at ordinary temperatures or pressures. The chlorine oxides (Cl₂O, ClO₂, Cl₂O₇) are all prepared by indirect synthesis (e.g. dehydration of HClO₄ with P₄O₁₀ to give Cl₂O₇). For spec purposes Cl₂O₇ is named as the formal Period 3 chlorine oxide and its acidic behaviour examined, but no direct combustion equation is required.

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Reactions of Period 3 Elements with Water

The metals on the left of the period react with water at varying rates; the non-metals on the right do not. This is the second axis of the periodic trend and reinforces the metal/non-metal classification.

Sodium (Na)

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

• Vigorous, exothermic, visible reaction at room temperature. Sodium floats (ρ = 0.97 g cm⁻³), melts to a ball from the exotherm, fizzes as H₂ evolves, and the resulting solution turns universal indicator deep purple (pH ≈13–14).

Magnesium (Mg)

Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g) [very slow with cold water]

Mg(s) + H₂O(g) → MgO(s) + H₂(g) [with steam, rapid]

• Cold water: barely perceptible reaction; a few bubbles of H₂ after hours.
• Steam: vigorous, white flame, MgO formed (not Mg(OH)₂ — the temperature is above the dehydration point of Mg(OH)₂).

Aluminium (Al)

No reaction with water under normal conditions due to the passivating Al₂O₃ layer (see above). Once the oxide layer is removed (e.g. by mercury amalgamation), aluminium reacts vigorously with water: 2Al + 6H₂O → 2Al(OH)₃ + 3H₂.

Silicon, Phosphorus, Sulfur, Chlorine

Si, P, and S do not react with liquid water at room temperature in any meaningful sense. Cl₂ reacts reversibly with water in a disproportionation (Cl₂ + H₂O ⇌ HCl + HClO; covered in L2 — Group 7), but this is the halogen reacting, not Period 3 trend behaviour.

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Oxide Structures and Bonding

The bonding type of each Period 3 oxide is the single most important predictor of its acid–base behaviour. The transition runs ionic giant lattice → ionic with covalent character → giant covalent → molecular covalent.

• Na₂O — ionic giant lattice (antifluorite structure). Na⁺ and O²⁻; high lattice energy but the oxide ion is a strong base.
• MgO — ionic giant lattice (rock-salt structure). Mg²⁺ and O²⁻; very high lattice energy (≈ −3800 kJ mol⁻¹) giving an exceptionally high mp (≈2850 °C) — MgO is used as refractory furnace lining.
• Al₂O₃ — best described as ionic with significant covalent character. By Fajans's rules, the small, highly charged Al³⁺ cation (ionic radius 54 pm, charge +3) heavily polarises the larger O²⁻ anion (140 pm), distorting the electron cloud and giving partial covalent character to what would otherwise be a purely ionic bond. This mid-period "switching point" is why Al₂O₃ is amphoteric.
• SiO₂ — giant covalent network of corner-sharing SiO₄ tetrahedra. Each Si bonded to 4 O, each O bridging 2 Si. Strong, directional Si–O covalent bonds throughout the crystal. mp = 1713 °C.
• P₄O₁₀ — molecular covalent. Discrete P₄O₁₀ molecules with a tetrahedral cage of four P atoms each bonded to four O atoms (one terminal, three bridging). Weak London forces between molecules — mp ≈ 340 °C, sublimes.
• SO₂, SO₃ — molecular covalent. SO₂ is bent (V-shaped, ≈119°); SO₃ is trigonal planar in the gas phase (it polymerises to (SO₃)₃ trimers and chains in the solid). Both have low boiling points (SO₂ bp −10 °C; SO₃ bp 45 °C).
• Cl₂O₇ — molecular covalent. A liquid at room temperature (bp 82 °C); two ClO₃ groups bridged by one O atom; very weak intermolecular forces.

Summary Table

Oxide	Formula	Structure / bonding	mp / bp	Acid–base character
Sodium oxide	Na₂O	Ionic giant lattice	mp ≈1132 °C	Strongly basic
Magnesium oxide	MgO	Ionic giant lattice	mp ≈2850 °C	Basic
Aluminium oxide	Al₂O₃	Ionic with covalent character	mp ≈2072 °C	Amphoteric
Silicon dioxide	SiO₂	Giant covalent	mp 1713 °C	Acidic (weak)
Phosphorus(V) oxide	P₄O₁₀	Molecular covalent	mp 340 °C	Acidic
Sulfur dioxide	SO₂	Molecular covalent	bp −10 °C	Acidic
Sulfur trioxide	SO₃	Molecular covalent	bp 45 °C	Strongly acidic
Dichlorine heptoxide	Cl₂O₇	Molecular covalent	bp 82 °C	Strongly acidic

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Reactions of the Oxides with Water (pH Trend)

Moving across the period, the oxides change from basic (pH ≈14) through insoluble/amphoteric (pH ≈7, neutral water film) to acidic (pH ≈1). The unifying mechanism: ionic oxide ions (O²⁻) react with water as a strong base (O²⁻ + H₂O → 2OH⁻), while covalent non-metal oxides react with water to form oxyacids whose protons dissociate.

Na₂O — strongly basic (pH ≈14)

Na₂O(s) + H₂O(l) → 2NaOH(aq)

• Dissolves vigorously to give a strongly alkaline solution (pH ≈14).
• Mechanism: O²⁻ + H₂O → 2OH⁻. The oxide ion is the conjugate base of OH⁻ and is a stronger base than OH⁻ itself — hence quantitative reaction.

MgO — basic, sparingly soluble (pH ≈9–10)

MgO(s) + H₂O(l) → Mg(OH)₂(aq, sparingly)

• Reacts slowly with water; Mg(OH)₂ has low solubility (K_sp ≈ 5.6 × 10⁻¹²).
• Produces a weakly alkaline solution (pH ≈9–10) — the basis of "milk of magnesia" antacid.
• The higher lattice energy of MgO (Mg²⁺, O²⁻ both +/−2) compared to Na₂O (Na⁺, O²⁻) makes the oxide much harder to break apart, limiting dissolution.

Al₂O₃ — amphoteric, insoluble in pure water

Al₂O₃ is insoluble in pure water and the surrounding water film is neutral (pH ≈7). However, the oxide reacts with both acids and bases — the defining test of amphoteric behaviour.

With acid (Al₂O₃ behaves as a base):

Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)

Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l)

With base (Al₂O₃ behaves as an acid):

Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq)

• The product with NaOH is sodium tetrahydroxoaluminate(III), containing the [Al(OH)₄]⁻ complex ion. Al³⁺ is too charge-dense to exist as a bare ion in solution; it is always coordinated by 4 or 6 water/hydroxide ligands.
• This dual reactivity is the textbook definition of amphoteric behaviour and the key Period 3 datum at the metal/non-metal boundary.

SiO₂ — acidic, insoluble in water

SiO₂ does not react with water (the giant covalent network is too stable; no surface acid–base sites accessible to liquid water under normal conditions). It is therefore classed as insoluble, but is still acidic in character because it reacts with hot, concentrated NaOH:

SiO₂(s) + 2NaOH(aq) → Na₂SiO₃(aq) + H₂O(l)

• Produces sodium silicate ("water glass"), a useful industrial reagent.
• This is the synoptic A-Level point that the most common test of "acidic oxide" — reaction with water to give pH < 7 — fails for SiO₂. The robust test is reaction with a base, which SiO₂ passes.
• The same logic explains why HF (and not HCl) is needed to etch SiO₂ glass: SiO₂ + 4HF → SiF₄ + 2H₂O; this is the basis of microelectronics fabrication.

P₄O₁₀ — strongly acidic (pH ≈1)

P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)

• Reacts vigorously, sometimes violently, with water to produce phosphoric(V) acid (orthophosphoric acid).
• pH ≈1 in concentrated solution; H₃PO₄ is a moderately strong triprotic acid (pK_a1 = 2.15, pK_a2 = 7.20, pK_a3 = 12.35).
• P₄O₁₀ is hygroscopic and is used industrially as a powerful dehydrating agent.

SO₂ — acidic (pH ≈3)

SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq) [sulfurous acid]

• pH ≈3 in saturated solution.
• H₂SO₃ is a weak diprotic acid (pK_a1 = 1.85, pK_a2 = 7.2). The equilibrium lies well to the left in dilute solution — most "dissolved SO₂" is actually solvated SO₂(aq), not true H₂SO₃.

SO₃ — very strongly acidic (pH ≈1)

SO₃(g) + H₂O(l) → H₂SO₄(aq)

• Reacts extremely exothermically with water (the reaction is so vigorous that in the Contact process SO₃ is dissolved in concentrated H₂SO₄, not water, to form oleum H₂S₂O₇, which is then carefully diluted).
• Sulfuric acid is a strong diprotic acid; in 1 mol dm⁻³ solution pH ≈ −0.3 (the first dissociation is essentially complete).

Cl₂O₇ — very strongly acidic (pH ≈1)

Cl₂O₇(l) + H₂O(l) → 2HClO₄(aq)

• Produces perchloric acid HClO₄, one of the strongest mineral acids (pK_a ≈ −10). pH ≈1 in moderately dilute solution; HClO₄ dissociates completely.

pH Trend at a Glance

Oxide	pH of aqueous environment	Notes
Na₂O	≈14	Quantitative reaction with H₂O
MgO	≈9–10	Sparingly soluble; saturated Mg(OH)₂
Al₂O₃	≈7	Insoluble; amphoteric on contact with acid/base
SiO₂	≈7	Insoluble in H₂O; acidic with hot conc. NaOH
P₄O₁₀	≈1–2	H₃PO₄(aq)
SO₂	≈3	Weak acid H₂SO₃
SO₃	≈0–1	Strong acid H₂SO₄
Cl₂O₇	≈0–1	Strong acid HClO₄

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Reactions of Period 3 Oxides with Acids and Bases

A useful spec-aligned summary: every oxide is classified by its reactivity profile.

Basic Oxides with Acid

Na₂O(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l)

Na₂O(s) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O(l)

MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)

MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l)

Amphoteric Oxide — Al₂O₃

With acid:

Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)

With base:

Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq)

Acidic Oxides with Base

SiO₂(s) + 2NaOH(aq) → Na₂SiO₃(aq) + H₂O(l) [hot, conc.]

P₄O₁₀(s) + 12NaOH(aq) → 4Na₃PO₄(aq) + 6H₂O(l)

SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l)

SO₃(g) + 2NaOH(aq) → Na₂SO₄(aq) + H₂O(l)

Exam Tip: Amphoteric means "reacts with both acids and bases." Al₂O₃ is the only Period 3 example. Learn both equations: with 6HCl giving 2AlCl₃ + 3H₂O, and with 2NaOH + 3H₂O giving 2NaAl(OH)₄.

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Period 3 Chlorides — Bonding and Hydrolysis

The same ionic→covalent transition runs through the Period 3 chlorides, and the simplest practical readout is what happens when they meet water.

Bonding and structure

• NaCl — ionic giant lattice (rock salt). Na⁺ and Cl⁻.
• MgCl₂ — predominantly ionic, but Mg²⁺ has enough polarising power to introduce small covalent character (Fajans). Layered CdCl₂-type lattice.
• AlCl₃ — covalent. Al³⁺ is so polarising that the bond becomes substantially covalent. Solid AlCl₃ sublimes at ≈180 °C as the dimer Al₂Cl₆ (two Al atoms bridged by two Cl atoms, with each Al sp³-hybridised and tetrahedrally coordinated). The dimer dissociates to monomeric AlCl₃ above ≈400 °C.
• SiCl₄ — molecular covalent. Tetrahedral SiCl₄ molecules; liquid at RT (bp 58 °C).
• PCl₃, PCl₅ — molecular covalent. PCl₃ is pyramidal (bp 76 °C); PCl₅ is trigonal bipyramidal in the gas phase but ionic [PCl₄]⁺[PCl₆]⁻ in the solid.

Hydrolysis

Chloride	Behaviour with water	Solution
NaCl	Dissolves; no hydrolysis	Neutral (pH 7)
MgCl₂	Dissolves; very slight hydrolysis of [Mg(H₂O)₆]²⁺	pH ≈6.5
AlCl₃	Partial hydrolysis; [Al(H₂O)₆]³⁺ is acidic	pH ≈3
SiCl₄	Complete hydrolysis — fumes in moist air	Strongly acidic
PCl₃	Complete hydrolysis	Strongly acidic
PCl₅	Complete hydrolysis (in 2 steps)	Strongly acidic

Key equations:

AlCl₃ (partial hydrolysis via hexaaqua complex):

[Al(H₂O)₆]³⁺(aq) ⇌ [Al(H₂O)₅(OH)]²⁺(aq) + H⁺(aq)

SiCl₄ (complete hydrolysis):

SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g)

(In practice, hydrated SiO₂·xH₂O forms as a gelatinous solid.)

PCl₃, PCl₅ hydrolysis:

PCl₃(l) + 3H₂O(l) → H₃PO₃(aq) + 3HCl(aq)

PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(aq)

Mechanistic logic: as the chloride becomes more covalent, the chlorine atoms become more accessible to nucleophilic attack by water, and the central atom (Al, Si, P) accepts oxygen ligands from water in place of chlorine. Ionic NaCl simply dissolves; covalent SiCl₄ reacts.

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Practical-Skills Box — Testing Oxide pH