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9.1 Periodicity of physical properties of the elements in Period 3

Syllabus
9701–2028–2029
Topic
9.1
Level
AS

Period 3 shows four distinct physical-property patterns

Property Qualitative variation across Na, Mg, Al, Si, P, S, Cl, Ar
atomic radius decreases from Na to Cl as atoms use the same principal shell while nuclear charge increases; Ar has no directly comparable covalent radius in the usual dataset
ionic radius decreases across Na⁺ → Mg²⁺ → Al³⁺ → Si⁴⁺; jumps to a much larger value at P³⁻; then decreases P³⁻ → S²⁻ → Cl⁻
melting point rises Na → Mg → Al → Si, with Si highest; drops sharply at P; rises at S; then falls through Cl to Ar
electrical conductivity increases across the metals Na → Mg → Al; drops enormously at Si; P, S, Cl and Ar are effectively non-conducting

The cation sequence Na⁺, Mg²⁺, Al³⁺ and Si⁴⁺ is isoelectronic with 10 electrons, so increasing proton number pulls the same electron arrangement inward. P³⁻, S²⁻ and Cl⁻ are isoelectronic with 18 electrons and also shrink as proton number rises. The jump occurs because the anions have an occupied third shell, whereas the cations have lost it.

These linked but non-uniform variations are periodicity: patterns recur because outer-shell configuration and resulting structure change systematically across a period.

Do not draw one smooth trend through the ionic radii: cations and anions belong to different electron-shell groups. Do not describe melting point as simply decreasing after Si—S melts above P because S₈ particles are larger than P₄ particles.

Structure controls what must move for melting and for conduction

Elements Structure and bonding Melting-point explanation Electrical-conductivity explanation
Na, Mg, Al giant metallic lattice: positive ions attracted to delocalised electrons generally rises Na → Al as ion charge, number of delocalised electrons and charge density increase, strengthening metallic bonding all conduct through mobile delocalised electrons; conductivity increases overall toward Al as more electrons are contributed per atom
Si giant covalent network highest: many strong Si–Si covalent bonds must be broken to melt the network semiconductor with conductivity far below the metals because only a small number of mobile charge carriers are available
P₄, S₈, Cl₂ simple molecular; instantaneous dipole–induced dipole forces between molecules much lower than Si because melting overcomes intermolecular forces, not internal covalent bonds; S₈ > P₄ > Cl₂ as molecular size and polarisability increase essentially do not conduct because the molecules provide no mobile charged particles
Ar monatomic particles with instantaneous dipole–induced dipole attractions lowest: small Ar atoms have very weak attractions between them does not conduct because there are no mobile charged particles

For melting, identify the particles separated and the attractions between those particles. For conduction, ask whether the structure contains charge carriers that can move through the solid. These are separate tests: strong bonding can produce a high melting point without producing conductivity.

Silicon and aluminium are both extended lattices, but Al contains mobile delocalised electrons whereas pure Si has far fewer mobile carriers. Sulfur and phosphorus contain strong covalent bonds inside S₈ and P₄, yet their low melting points reflect the weaker forces between intact molecules.

Argon is monatomic, not a simple molecule. Do not say that melting P₄, S₈ or Cl₂ breaks covalent bonds, and do not infer electrical conduction solely from a high melting point.