Periodic trends across Period 3

What you'll learn

The elements of Period 3 — sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine and argon — show a clear progression from reactive metals to reactive non-metals to an unreactive noble gas. Studying them reveals how structure and bonding change across a period, and how these changes control physical properties and the behaviour of the elements' oxides and chlorides. This guide covers trends in atomic radius, ionic radius, first ionisation energy, electronegativity, melting point and electrical conductivity; the reactions of the elements with oxygen, chlorine and water; the structure, bonding and acid–base character of the oxides; and the structure, bonding and behaviour in water of the chlorides.

CAPE questions frequently ask you to explain melting point trends, to write equations for reactions of oxides and chlorides with water, acids and bases, and to explain amphoteric behaviour, especially of aluminium oxide — a compound central to Jamaica's bauxite and alumina industry.

Key terms and definitions

Period — A horizontal row of the Periodic Table; Period 3 elements have outer electrons in the third shell.

Atomic radius — Half the distance between the nuclei of two bonded atoms of the same element.

Electronegativity — The ability of an atom to attract the bonding pair in a covalent bond.

Basic oxide — An oxide that reacts with acids to form a salt and water.

Acidic oxide — An oxide that reacts with bases to form a salt and water, often forming an acid with water.

Amphoteric oxide — An oxide that reacts with both acids and bases.

Hydrolysis — The reaction of a compound with water, often breaking bonds and producing H⁺ ions.

Giant metallic, giant covalent, simple molecular — Structure types that change across the period.

Core concepts

Structure and bonding of the elements

Physical properties

Atomic radius decreases from Na to Cl because nuclear charge increases while electrons are added to the same shell, with similar shielding, so the outer electrons are pulled closer.

Ionic radius: Na⁺, Mg²⁺ and Al³⁺ are much smaller than their atoms (they have lost the third shell), and decrease in size with increasing charge. P³⁻, S²⁻ and Cl⁻ are much larger than the metal ions (they have three shells), and also decrease in size across.

Electronegativity increases across the period (except argon, which does not usually form bonds).

First ionisation energy generally increases, with dips at aluminium (3p electron) and sulfur (paired 3p electrons).

Melting point:

• rises from Na to Al as the metallic bonding gets stronger — more delocalised electrons per atom and smaller, more highly charged ions;
• peaks at Si, which has a giant covalent structure with many strong covalent bonds to break;
• drops sharply to P, S and Cl, which are simple molecules held by weak London forces; S₈ has a higher melting point than P₄ because it has more electrons, and Cl₂ is lower still;
• is lowest for Ar, whose single atoms are held only by very weak London forces.

Electrical conductivity: high for Na, Mg and Al (delocalised electrons, increasing from Na to Al); silicon is a semiconductor; P, S, Cl and Ar do not conduct.

Reactions of the elements

With oxygen:

• 4Na + O₂ → 2Na₂O (bright yellow flame)
• 2Mg + O₂ → 2MgO (brilliant white flame)
• 4Al + 3O₂ → 2Al₂O₃ (protected by an oxide layer unless powdered)
• Si + O₂ → SiO₂ (on strong heating)
• P₄ + 5O₂ → P₄O₁₀ (white phosphorus ignites spontaneously; in excess oxygen)
• S + O₂ → SO₂ (blue flame)

With chlorine:

• 2Na + Cl₂ → 2NaCl
• Mg + Cl₂ → MgCl₂
• 2Al + 3Cl₂ → Al₂Cl₆ (or 2AlCl₃)
• Si + 2Cl₂ → SiCl₄
• P₄ + 10Cl₂ → 4PCl₅ (with excess chlorine; PCl₃ with limited chlorine)

With water:

• Sodium reacts vigorously with cold water: 2Na + 2H₂O → 2NaOH + H₂ (solution pH about 13–14).
• Magnesium reacts very slowly with cold water but readily with steam: Mg + H₂O → MgO + H₂.

The oxides

Reactions with water:

• Na₂O + H₂O → 2NaOH
• MgO + H₂O → Mg(OH)₂ (sparingly soluble, weakly alkaline)
• P₄O₁₀ + 6H₂O → 4H₃PO₄
• SO₂ + H₂O → H₂SO₃; SO₃ + H₂O → H₂SO₄

Acid–base reactions:

• Basic: MgO + 2HCl → MgCl₂ + H₂O
• Acidic: SiO₂ + 2NaOH → Na₂SiO₃ + H₂O (hot, concentrated NaOH)
• Amphoteric aluminium oxide:
  • with acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
  • with base: Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄ (sodium aluminate)

The Bayer process, used to refine Jamaican bauxite, depends on this amphoteric behaviour: hot sodium hydroxide dissolves the aluminium oxide as aluminate, leaving iron oxide impurities ("red mud") behind. Aluminium hydroxide is then precipitated and heated to give pure alumina.

The chlorides

Reactions with water:

• NaCl dissolves to form hydrated ions; the solution is neutral.
• MgCl₂ dissolves; slight hydrolysis gives a slightly acidic solution.
• Aluminium chloride hydrolyses: the small, highly charged Al³⁺ ion polarises the water molecules around it, releasing H⁺: [Al(H₂O)₆]³⁺ ⇌ [Al(H₂O)₅(OH)]²⁺ + H⁺
• SiCl₄ + 2H₂O → SiO₂ + 4HCl (steamy fumes of HCl)
• PCl₅ + 4H₂O → H₃PO₄ + 5HCl

The trend from ionic, neutral chlorides to covalent chlorides that hydrolyse to give acidic solutions mirrors the trend in the oxides.

Worked examples

Example 1: Melting point of silicon

Explain why silicon has the highest melting point in Period 3. (3 marks)

Silicon has a giant covalent (macromolecular) structure (1) in which each atom is joined to four others by strong covalent bonds (1). A large amount of energy is needed to break these many bonds (1).

Example 2: Amphoteric behaviour

Write equations to show that aluminium oxide is amphoteric.

With acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O. With base: Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄.

Example 3: Chlorides with water

Explain why adding SiCl₄ to water produces a strongly acidic solution while NaCl gives a neutral solution.

NaCl is ionic and simply dissolves, forming Na⁺(aq) and Cl⁻(aq), which do not react with water. SiCl₄ is covalent and is hydrolysed by water: SiCl₄ + 2H₂O → SiO₂ + 4HCl, and the HCl produced dissociates to give a low pH.

Example 4: Comparing sulfur and phosphorus

Explain why sulfur has a higher melting point than phosphorus.

Both are simple molecular, held by London forces. Sulfur exists as S₈ molecules, which have more electrons than P₄ molecules, so the London forces between S₈ molecules are stronger and more energy is needed to separate them.

Example 5: Predicting pH when compounds are added to water

Predict the approximate pH when each is added to water: (a) Na₂O, (b) SO₃, (c) NaCl, (d) PCl₅, (e) MgO.

(a) pH 13–14: Na₂O + H₂O → 2NaOH, a strong alkali. (b) pH 1: SO₃ + H₂O → H₂SO₄, a strong acid. (c) pH 7: NaCl dissolves to form hydrated ions without reacting. (d) pH 1–2: PCl₅ is hydrolysed to H₃PO₄ and HCl. (e) pH about 9: MgO forms sparingly soluble Mg(OH)₂, giving a weakly alkaline solution.

Example 6: Reactions of oxides with acids and bases

Write equations to show the reactions of (a) magnesium oxide with hydrochloric acid, (b) silicon dioxide with hot concentrated sodium hydroxide, (c) sulfur dioxide with sodium hydroxide.

(a) MgO + 2HCl → MgCl₂ + H₂O (basic oxide). (b) SiO₂ + 2NaOH → Na₂SiO₃ + H₂O (acidic oxide; forms sodium silicate). (c) SO₂ + 2NaOH → Na₂SO₃ + H₂O (acidic oxide; forms sodium sulfite).

Example 7: Explaining conductivity across the period

Explain the trend in electrical conductivity from sodium to argon.

Sodium, magnesium and aluminium are metals with delocalised electrons; conductivity increases from Na to Al because the number of delocalised electrons per atom increases from one to three. Silicon is a semiconductor: its electrons are held in covalent bonds, but a few can be released, especially when heated or doped. Phosphorus, sulfur, chlorine and argon are non-metals with no free electrons or ions, so they are insulators.

Example 8: The aluminium chloride dimer

Explain why aluminium chloride exists as Al₂Cl₆ in the vapour at moderate temperatures.

In AlCl₃, aluminium has only six electrons in its outer shell and is electron-deficient. A lone pair on a chlorine atom of one AlCl₃ molecule forms a dative bond to the aluminium of another, and vice versa, so two molecules join into a dimer in which each aluminium has a complete octet. At higher temperatures, the dimer separates into AlCl₃ molecules.

Oxide and chloride trends summary

The change from ionic to covalent bonding across the period, driven by increasing electronegativity of the element, explains the change from basic to acidic behaviour.

Common mistakes and how to avoid them

• Saying silicon's high melting point is due to metallic bonding. It is giant covalent.
• Saying P, S and Cl have low melting points because covalent bonds are weak. It is the intermolecular forces that are weak.
• Forgetting to balance equations, especially for P₄O₁₀ and Al₂O₃.
• Missing the amphoteric nature of Al₂O₃.
• Writing that SiO₂ dissolves in water — it is insoluble, though acidic.

How this links to your Internal Assessment

Practical work on Period 3 might involve testing the pH of solutions formed when oxides and chlorides are added to water (under supervision, using small quantities and a fume cupboard for hydrolysing chlorides), or investigating the amphoteric behaviour of aluminium hydroxide by adding acid and alkali to precipitates. An SBA could also explore the chemistry of bauxite processing using a simulated Bayer-style dissolution of aluminium hydroxide in sodium hydroxide. Record observations carefully, write equations, and explain results in terms of bonding and structure.

Exam technique for Periodic trends across Period 3

• Link every physical property to structure and bonding.
• Use the oxide and chloride tables to organise answers: bonding → structure → behaviour with water → pH.
• Write balanced equations with state symbols where appropriate.
• When explaining acidity of aluminium chloride solutions, mention polarisation by the small, highly charged Al³⁺ ion.

Quick revision summary

• Structures: metallic (Na, Mg, Al) → giant covalent (Si) → simple molecular (P₄, S₈, Cl₂) → monatomic (Ar).
• Atomic radius decreases; electronegativity and ionisation energy generally increase across.
• Melting point rises to Si, then drops sharply; conductivity high for metals, Si semiconductor.
• Oxides: basic (Na₂O, MgO) → amphoteric (Al₂O₃) → acidic (SiO₂, P₄O₁₀, SO₂, SO₃).
• Chlorides: ionic and neutral (NaCl) → covalent and hydrolysed to acidic solutions (AlCl₃, SiCl₄, PCl₅).
• Amphoteric Al₂O₃ underpins the Bayer process used in Jamaica's alumina industry.