The Periodic Table: Periodicity Across Period 3

This guide covers subtopics 9.1, Periodicity of physical properties of the elements in Period 3, 9.2, Periodicity of chemical properties of the elements in Period 3, and 9.3, Chemical periodicity of other elements, from Topic 9 of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. All three are AS Level content, grouped here because they tell one continuous story: the same underlying pattern (structure and bonding changing across a period) explains both the physical trends and the chemical trends, and 9.3 is simply that pattern applied predictively.

Before studying this

This resource assumes Atomic Structure: Particles, Radius and Isotopes — the explanation of atomic radius across a period in terms of nuclear charge and shielding is reused directly here — and Chemical Bonding: Electronegativity, Ionic and Metallic Bonds for electronegativity and predicting ionic vs covalent bonding. “Periodicity” means a property recurs in a regular, repeating pattern as you move through the Periodic Table — Period 3 (sodium to argon) is the specific period this topic uses to demonstrate it in depth.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 9

9.1 Periodicity of physical properties — describing qualitatively (and identifying the periodicity in) the variation in atomic radius, ionic radius, melting point and electrical conductivity across Period 3; explaining the variation in melting point and electrical conductivity in terms of structure and bonding.

9.2 Periodicity of chemical properties — describing and writing equations for the reactions of the Period 3 elements with oxygen, chlorine and water (sodium and magnesium only); stating and explaining the variation in oxidation number of the oxides and chlorides; describing and writing equations for the reactions of the oxides and chlorides with water, including the resulting pH; describing, explaining and writing equations for the acid/base behaviour of the oxides and hydroxides, including amphoteric behaviour; explaining these variations in terms of bonding and electronegativity; suggesting the type of bonding present from observed properties.

9.3 Chemical periodicity of other elements — predicting the characteristic properties of an element in a given group using knowledge of periodicity; deducing the possible identity and Periodic Table position of an unknown element from given data.

Physical periodicity across Period 3

Atomic radius decreases across the period, for the same reason it does in any period: nuclear charge increases while shielding (same outer shell throughout) stays roughly constant, pulling the outer electrons in more tightly.

Melting point and electrical conductivity both depend on structure and bonding, which changes abruptly partway through the period:

Element	Na	Mg	Al	Si	P	S	Cl₂	Ar
Structure/bonding	metallic	metallic	metallic	giant covalent	simple molecular	simple molecular	simple molecular	monatomic
Melting point (°C, approx.)	98	650	660	1410	44	115	−101	−189
Conducts electricity?	yes	yes	yes	semiconductor	no	no	no	no

Na → Mg → Al: melting point rises because metallic bonding strengthens — each successive element contributes one more delocalised electron and forms a smaller, more highly charged cation, both of which increase the electrostatic attraction holding the lattice together. All three conduct, because a metallic lattice has delocalised electrons free to move.

Si has by far the highest melting point in the period, because it forms a giant covalent (macromolecular) lattice — every atom is joined to its neighbours by strong covalent bonds extending through the whole structure, and melting means breaking a great many of those bonds at once. Silicon conducts only very weakly (it’s a semiconductor), since — unlike a metal — it has no delocalised electrons free to move; only a small number of electrons gain enough energy to become mobile.

P → S → Cl₂ → Ar: melting point drops sharply and none of these conduct, because from silicon onwards the elements exist as simple molecular (P₄, S₈, Cl₂) or monatomic (Ar) structures — held together in the solid only by weak instantaneous dipole–induced dipole (van der Waals) forces between molecules, while the strong covalent bonds within each molecule are irrelevant to melting. The small remaining variation (S₈ > P₄ > Cl₂ > Ar) follows molecule size and number of electrons, exactly as for any van der Waals comparison.

Chemical periodicity across Period 3

Reactions with oxygen (the elements typically studied: Na, Mg, Al, P, S):

4Na + O₂ → 2Na₂O

2Mg + O₂ → 2MgO

4Al + 3O₂ → 2Al₂O₃

P₄ + 5O₂ → P₄O₁₀

S + O₂ → SO₂

Reactions with chlorine:

2Na + Cl₂ → 2NaCl

Mg + Cl₂ → MgCl₂

2Al + 3Cl₂ → 2AlCl₃

Si + 2Cl₂ → SiCl₄

2P + 5Cl₂ → 2PCl₅

Reactions with water (sodium and magnesium only): sodium reacts vigorously with cold water, 2Na + 2H₂O → 2NaOH + H₂; magnesium reacts only extremely slowly with cold water, but readily with steam, Mg + H₂O(g) → MgO

• H₂.

Oxidation number across the period

The oxidation number of each element in its oxide and chloride rises steadily across the period — Na (+1), Mg (+2), Al (+3), Si (+4), P (+5), S (+6 in SO₃) — because each element is using progressively more of its outer-shell (valence-shell) electrons in bonding as the number of those electrons increases across the period.

Oxides and chlorides with water

This is the part worth understanding as a pattern, not a list to memorise mechanically: ionic oxides/chlorides (Na, Mg) behave very differently from covalent ones (Si, P, S).

Oxide	Reaction with water	Resulting pH
Na₂O	Na₂O + H₂O → 2NaOH	~13–14 (strongly alkaline)
MgO	MgO + H₂O → Mg(OH)₂ (sparingly soluble)	~9–10 (weakly alkaline)
Al₂O₃	no reaction (insoluble)	—
SiO₂	no reaction (insoluble)	—
P₄O₁₀	P₄O₁₀ + 6H₂O → 4H₃PO₄	~0–1 (strongly acidic)
SO₂	SO₂ + H₂O ⇌ H₂SO₃	~2–3 (weakly acidic)
SO₃	SO₃ + H₂O → H₂SO₄	~0 (strongly acidic)

Acid/base character: Na₂O and MgO are basic oxides. P₄O₁₀, SO₂ and SO₃ are acidic oxides. Al₂O₃ is amphoteric — it does not react with water, but reacts with both an acid and (specifically) sodium hydroxide:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

The corresponding hydroxides follow the same pattern: NaOH is a strong base, Mg(OH)₂ a weak (sparingly soluble) base, and Al(OH)₃ amphoteric, reacting with both HCl and NaOH in the same way as Al₂O₃.

Chlorides with water split the same way, but the explanation is different — it’s about ionic vs covalent bonding, not acid/base oxide character:

• NaCl simply dissolves; Na⁺ and Cl⁻ don’t react with water. pH ≈ 7.
• MgCl₂ dissolves and is weakly acidic (pH ≈ 6): the small, doubly charged Mg²⁺ ion polarises coordinated water molecules enough to release some H⁺.
• AlCl₃ is more strongly acidic still (pH ≈ 3), for the same reason taken further — Al³⁺ is smaller and more highly charged, so it polarises coordinated water even more.
• SiCl₄ and PCl₅ hydrolyse completely and violently (“fume” in moist air), because they are covalent molecules that react directly with water:

SiCl₄ + 2H₂O → SiO₂ + 4HCl

PCl₅ + 4H₂O → H₃PO₄ + 5HCl

Why the split happens where it does: electronegativity increases across the period while chlorine’s stays fixed, so the electronegativity difference between the element and chlorine shrinks — Na–Cl and Mg–Cl bonds are ionic enough that the chloride simply dissolves as ions, while Si–Cl and P–Cl bonds are covalent enough that water attacks the bond directly, hydrolysing it.

Predicting properties from periodicity (9.3)

Because these trends repeat in every period, you can use the pattern itself to predict or deduce things you haven’t been told directly.

Worked example. An unknown Period 3-like element X reacts with water to give a strongly acidic solution and, on analysis, its chloride is found to hydrolyse violently in moist air. Suggest whereabouts in a period X is likely to sit, and justify your answer.

A violently hydrolysing, “fuming” chloride indicates a covalent chloride — the kind formed towards the right-hand side of a period (by analogy with Si and P in Period 3), not the ionic chlorides typical of the left-hand side. A strongly acidic oxide similarly points to a non-metal towards the right of the period, since acidic oxides (like P₄O₁₀, SO₂, SO₃) sit on that side while basic oxides (like Na₂O, MgO) sit on the left. Both pieces of evidence place X on the right-hand side of its period.

Common mistakes

• Explaining Si’s high melting point as “strong covalent bonds,” full stop. The key comparison is that Si’s bonds extend through a giant lattice, so melting breaks many strong bonds — whereas P₄, S₈ and Cl₂ also have strong covalent bonds within each molecule, but melting only needs to overcome the weak forces between molecules.
• Assuming all Period 3 chlorides behave like NaCl. Only the metallic elements’ chlorides are ionic and simply dissolve; from silicon onwards, the chlorides are covalent and hydrolyse.
• Forgetting that Al₂O₃ doesn’t react with water at all, despite reacting readily with both acids and NaOH — amphoteric means reacting with acid and base, not “reacts with everything.”
• Mixing up which oxides are acidic and which are basic. The pattern follows metallic character: metals (left) give basic oxides, non-metals (right) give acidic oxides, with aluminium’s amphoteric oxide marking the transition.

Quick revision checklist

• Physical trends: atomic radius (nuclear charge/shielding), melting point and conductivity (metallic → giant covalent → simple molecular/monatomic)
• Reactions of Na, Mg, Al, P, S with oxygen and chlorine; Na and Mg with water
• Oxidation number rising across the period, following valence electrons
• Oxide + water reactions and pH: basic (Na, Mg) → amphoteric (Al) → insoluble (Si) → acidic (P, S)
• Chloride + water reactions: ionic chlorides dissolve, covalent chlorides (Si, P) hydrolyse violently
• Explaining the ionic/covalent chloride split via shrinking electronegativity difference
• Using the pattern predictively for unknown elements (9.3)

Related resources

• Atomic Structure: Particles, Radius and Isotopes — the atomic radius trend this resource extends
• Chemical Bonding: Electronegativity, Ionic and Metallic Bonds — electronegativity and predicting bond type
• Group 2: The Alkaline Earth Metals — periodicity applied down a single group instead of across a period
• The Periodic Table: Groups and Trends — the IGCSE/O Level Group I, VII and VIII treatment this resource builds on
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.