Period 3 Structure and Bonding: Na to Ar

Spec Mapping — OCR H432 Module 3.1.1 — Periodicity, covering the structures and bonding of the Period 3 elements Na, Mg, Al, Si, P, S, Cl, Ar, the link between structure type (giant metallic, giant covalent, simple molecular, monatomic) and physical properties (melting point, boiling point, electrical conductivity), and the explanation of the sharp drop in melting point between Si and P (refer to the official OCR H432 specification document for exact wording).

Period 3 is OCR's canonical case study in structure-property relationships: traverse the row from sodium to argon and you cross three distinct structural categories — giant metallic, giant covalent, simple molecular (plus monatomic at the very end) — and watch the melting point swing from 98 °C (Na) up to 1414 °C (Si, the peak) and back down to −189 °C (Ar). Every twist of the curve has a structural explanation rooted in the type of force broken on melting, and OCR examines this lesson aggressively in Paper 1 short-answer questions. The single most distinctive feature is the ~1370 °C drop between Si and P, where the structure changes abruptly from a giant covalent network to discrete P₄ molecules held together only by weak London dispersion forces. Get the structure right and the physical-property reasoning falls out automatically — the goal of this lesson is to make the structure classification, and the reasoning that follows from it, automatic.

Key Definitions:

Giant metallic structure — a lattice of positive metal cations held together by a sea of delocalised valence electrons.

Giant covalent (network) structure — a continuous three-dimensional lattice of atoms covalently bonded to neighbours, with no discrete molecules.

Simple molecular structure — discrete small molecules held to one another by weak intermolecular forces (London dispersion in non-polar molecules).

Monatomic — the substance consists of single atoms (the noble gases Ne, Ar, Kr, Xe, Rn).

Overview of Period 3

Period 3 contains three structure types in sequence:

flowchart LR
  A[Na] --> B[Mg] --> C[Al] --> D[Si] --> E[P] --> F[S] --> G[Cl] --> H[Ar]
  A -.giant metallic.-> C
  D -.giant covalent.-> D
  E -.simple molecular P4.-> G
  H -.monatomic.-> H

Element	Z	Configuration	Structure	Bonding	m.p. / °C	b.p. / °C
Na	11	[Ne] 3s¹	Giant metallic	Metallic (1 delocalised e⁻ per atom)	98	883
Mg	12	[Ne] 3s²	Giant metallic	Metallic (2 delocalised e⁻)	650	1090
Al	13	[Ne] 3s² 3p¹	Giant metallic	Metallic (3 delocalised e⁻)	660	2470
Si	14	[Ne] 3s² 3p²	Giant covalent	Covalent network (4 single bonds per atom)	1414	3265
P	15	[Ne] 3s² 3p³	Simple molecular (P₄ white)	London forces	44	280
S	16	[Ne] 3s² 3p⁴	Simple molecular (S₈)	London forces	113	445
Cl	17	[Ne] 3s² 3p⁵	Simple molecular (Cl₂)	London forces	−101	−34
Ar	18	[Ne] 3s² 3p⁶	Monatomic	London forces	−189	−186

The melting-point curve has three distinct regions: a metallic rise (Na → Mg → Al), a giant-covalent peak (Si), and a simple-molecular trough (P → Ar). Each region has its own physical explanation.

Na, Mg, Al — giant metallic

All three are giant metallic lattices: a regular three-dimensional array of positive metal cations occupying lattice sites, with the valence electrons fully delocalised through the entire crystal as a "sea". This sea of mobile electrons binds the cations together (metallic bonding) and is responsible for the electrical conductivity, thermal conductivity, and malleability of metals.

Moving Na → Mg → Al, three quantities all change in the same direction:

Cation charge increases: Na⁺, Mg²⁺, Al³⁺.
Number of delocalised electrons per atom increases: 1, 2, 3 (because Na has 1 valence electron, Mg has 2, Al has 3).
Ionic radius decreases: Na⁺ ~95 pm, Mg²⁺ ~65 pm, Al³⁺ ~50 pm.

All three effects strengthen the metallic bond (cations are smaller, more positively charged, more strongly attracted to a denser electron sea), so the energy to overcome the bond rises rapidly. The melting point jumps from 98 °C (Na) to 650 °C (Mg) to 660 °C (Al). (The Mg→Al rise is small because the additional delocalised electron is offset by Al's slightly larger cation; the boiling-point comparison is more dramatic because boiling requires complete dispersal of the electron sea.)

Metal	Cation	Outer e⁻ delocalised per atom	m.p. / °C	b.p. / °C
Na	Na⁺	1	98	883
Mg	Mg²⁺	2	650	1090
Al	Al³⁺	3	660	2470

The boiling points (which require complete vaporisation of the lattice) show the trend more cleanly than the melting points, because melting only loosens the lattice (the cations still sit in the electron sea after melting), while boiling completely disperses both cations and electrons.

Si — giant covalent

Silicon adopts the diamond-like structure also seen in carbon (diamond), germanium and grey tin: each Si atom is covalently bonded to four neighbouring Si atoms in a tetrahedral arrangement, building up a continuous three-dimensional network with no discrete molecules. Every atom is part of one giant supramolecule.

To melt silicon, you must break strong covalent bonds throughout the lattice. Each Si–Si bond requires ~226 kJ mol⁻¹ to break, and each silicon atom is involved in four such bonds, so the energy per mole of atoms melted is enormous. The melting point is consequently ~1414 °C — the peak of the Period 3 m.p. graph — and the boiling point is even higher at 3265 °C.

Silicon's electrical conductivity is also distinctive: it is a semiconductor, with conductivity intermediate between metals and insulators. The valence electrons are localised in covalent bonds (so they cannot flow freely as in a metal), but the band gap is small enough that thermal excitation or doping can promote some electrons into a conduction band. This is the principle behind silicon's central role in microelectronics. At A-Level you simply note "semiconductor / poor conductor at room temperature".

P, S, Cl, Ar — simple molecular (and monatomic)

From phosphorus onwards, the structure changes abruptly to simple molecular:

Element	Molecular formula	Shape	Notes
P	P₄	Tetrahedral (4 P at corners of a tetrahedron)	"White phosphorus"; another allotrope (red P) is polymeric
S	S₈	Crown-shaped puckered 8-ring	The stable allotrope at room T (rhombic sulfur)
Cl	Cl₂	Linear diatomic	Yellow-green gas
Ar	Ar	Single spherical atom	Monatomic

Within each molecule, atoms are held by strong covalent bonds (P–P, S–S, Cl–Cl). Between molecules, only weak London dispersion forces (instantaneous-dipole–induced-dipole forces, induced fluctuations in electron density) act. Melting requires only overcoming these weak intermolecular forces — not the covalent bonds within the molecules — so the melting points are all low.

Within the simple-molecular group, the order is: $\text{S}_8\ (113) > \text{P}_4\ (44) > \text{Cl}_2\ (-101) > \text{Ar}\ (-189)$

The order tracks the number of electrons per molecule (electron count → polarisability → strength of London forces):

Species	Electrons per molecule	London-force strength	m.p. / °C
S₈	128	Strongest	113
P₄	60	Strong	44
Cl₂	34	Weak	−101
Ar	18 (per atom)	Weakest	−189

The larger the electron cloud, the more polarisable it is and the stronger the instantaneous-dipole / induced-dipole interactions between neighbouring molecules. S₈ has 128 electrons per molecule — the highest in the period — and therefore the strongest London forces, despite being further to the right than P₄ in the period.

Key point — and the single most-tested fact in this lesson: the covalent bonds within P₄, S₈, Cl₂ are not broken on melting. Only the weak intermolecular forces are. Stating otherwise loses marks immediately.

The Si → P drop

The sharp decrease in melting point from Si (1414 °C) to P₄ (44 °C) is the most dramatic feature of the Period 3 trend. The ~1370 °C drop in a single step occurs because:

Si has a giant covalent lattice — melting requires breaking strong Si–Si covalent bonds throughout the entire crystal.
P has a simple molecular structure — melting only requires overcoming the weak London dispersion forces between P₄ molecules; the P–P covalent bonds within each P₄ tetrahedron remain intact.

The energy difference between "breaking covalent bonds throughout a crystal" and "overcoming London forces between molecules" is roughly two orders of magnitude, and this is exactly the magnitude of the m.p. drop.

flowchart TD
  A[Si melting point ~1414 C] -->|Giant covalent lattice| B[Break covalent bonds across whole crystal]
  C[P4 melting point ~44 C] -->|Simple molecular| D[Overcome London forces between P4 molecules only]
  B --> E[Very high energy: m.p. is high]
  D --> F[Low energy: m.p. is low]
  E --> G[Si is m.p. peak in Period 3]
  F --> H[Drop from Si to P is ~1370 C]

Electrical conductivity across Period 3

Element	Electrical conductivity
Na, Mg, Al	Good — delocalised valence electrons in metallic bonding act as charge carriers
Si	Semiconductor — covalent electrons localised, but small band gap allows thermal/doped excitation
P, S, Cl, Ar	Poor / zero — no delocalised electrons and no free ions to carry charge

The transition mirrors the structural shift: metallic conductivity (delocalised sea), then semiconductor (filled valence band, small gap), then insulator (electrons locked in covalent bonds or single-atom orbitals).

Boiling points: a cleaner trend

If the Period 3 melting-point graph is a "double dip" (rise to Al, peak at Si, crash at P, low for the rest), the boiling-point graph is a smoother version of the same story:

Element	b.p. / °C
Na	883
Mg	1090
Al	2470
Si	3265
P	280
S	445
Cl	−34
Ar	−186

The boiling-point peak is even more pronounced (Si peaks at 3265 °C), and the Si → P drop is even more dramatic (~3000 °C). This is because boiling requires complete vaporisation — totally pulling apart the lattice — while melting only loosens it. For metals, melting still leaves the cation-electron-sea liquid largely intact; for giant covalent solids, melting requires almost the same energy as boiling because covalent bonds must be broken either way.

Worked examples

Example 1 — comparing Na, Mg, Si

Q: Explain why the melting point of magnesium is higher than that of sodium, but much lower than that of silicon.

Answer:

Na vs Mg: Both have giant metallic structures. Mg has a Mg²⁺ cation and contributes 2 delocalised electrons per atom, compared with Na⁺ and 1 delocalised electron for sodium. The metallic bond in Mg is therefore stronger (higher charge density, more delocalised electrons, smaller cation), so more energy is required to overcome it → higher melting point.

Mg vs Si: Mg is giant metallic; Si is giant covalent. Melting Si requires breaking strong covalent bonds throughout the diamond-type lattice (each Si atom is bonded to 4 others, ~226 kJ mol⁻¹ per Si–Si bond), which needs much more energy than overcoming the metallic bonds in Mg (where melting only loosens the lattice — the cations still sit in the electron sea after melting).

Period 3 Structure and Bonding