Structures: Simple Molecular, Giant Covalent and Metallic

Spec Mapping — OCR H432 Module 2.2.2 — Bonding and structure, covering the four structural types of solids (simple molecular, giant covalent, giant ionic, giant metallic), the structures of diamond, graphite, silicon(IV) oxide, and metals, the physical properties of each structural type (melting point, hardness, electrical conductivity, solubility), and the relation between structure, bonding, and bulk properties (refer to the official OCR H432 specification document for exact wording).

This is the final synthesis lesson of Module 2. Bonding (ionic, covalent, dative, metallic) and intermolecular forces (London, dipole-dipole, hydrogen bonding) combine to determine the structure of a solid — and the structure determines the physical properties (melting point, hardness, conductivity, solubility, density, malleability). The four-way classification of solids — simple molecular, giant covalent (macromolecular), giant ionic, giant metallic — covers essentially every pure substance you will meet in the OCR specification. This lesson revisits each in turn, with particular attention to the carbon allotropes (diamond and graphite — same atom, totally different properties), silicon(IV) oxide (the giant covalent analogue of CO₂), metals (the third bonding model), and the property contrasts that distinguish the four types. Mastery of this lesson allows you to predict the structure and properties of any unfamiliar substance from its chemical formula alone.

Key Definitions:

Simple molecular structure — small, discrete molecules held in the solid state by intermolecular forces only; the covalent bonds within molecules are strong, but those between are weak.

Giant covalent (macromolecular) structure — a continuous 3D network of atoms held together by covalent bonds throughout; no discrete molecules.

Giant ionic structure — a 3D lattice of cations and anions held by electrostatic forces (covered in Lesson 6).

Giant metallic structure — a 3D lattice of cations held by a delocalised "sea" of valence electrons.

The four structural types — overview

flowchart TD
    A[Solid structural types] --> B[Simple molecular]
    A --> C[Giant covalent]
    A --> D[Giant ionic]
    A --> E[Giant metallic]
    B --> B1[I2 CO2 H2O ice S8 P4 C60]
    B --> B2[Held by IMFs only]
    B --> B3[Low m.p. soft no conductivity]
    C --> C1[Diamond graphite SiO2 SiC]
    C --> C2[Covalent bonds throughout]
    C --> C3[Very high m.p. very hard usually insulating]
    D --> D1[NaCl MgO CaF2 Al2O3]
    D --> D2[Ions in 3D lattice held by electrostatic forces]
    D --> D3[High m.p. brittle conducts when molten]
    E --> E1[Na Fe Cu Al W]
    E --> E2[Cations + delocalised electrons]
    E --> E3[High m.p. malleable conducts in solid state]

A solid's structural type is determined by the bonding within and between its constituent units. The classification cascades naturally from the bonding type:

Metallic elements (Na, Mg, Fe, Cu) → giant metallic structure.
Metal + non-metal (NaCl, MgO, Al₂O₃) → giant ionic structure.
Two non-metals → either giant covalent (if extended 3D network — diamond, SiO₂) or simple molecular (if discrete molecules — I₂, CO₂, H₂O).

The non-metal/non-metal distinction is the trickier one — see the section on diamond vs CO₂ below for the discriminating logic.

1. Simple molecular structures

A simple molecular solid contains small, discrete molecules. The strong covalent bonds within each molecule are intramolecular; the weak intermolecular forces between the molecules hold the solid together.

Examples

Iodine I₂ — purple-grey solid; sublimes readily at 184 °C.
Carbon dioxide CO₂ (dry ice) — sublimes at −78 °C (no liquid phase at atmospheric pressure).
Ice H₂O — distinctive hydrogen-bonded tetrahedral network (Lesson 11).
Sulfur S₈ — yellow rhombic crystals; m.p. 113 °C.
White phosphorus P₄ — soft, yellow waxy solid; m.p. 44 °C.
Buckminsterfullerene C₆₀ — black solid; m.p. ~600 °C (discrete C₆₀ molecules with strong London forces because each molecule has 60 atoms).
Organic crystals — naphthalene, sugar (sucrose), aspirin, paracetamol.

Properties of simple molecular solids

Property	Observation	Explanation
Melting / boiling point	Low (≤ ~100 °C typical)	Only weak intermolecular forces between molecules; easily overcome thermally
Hardness	Soft or crumbly	Weak attractions between molecules
Electrical conductivity	None (insulator)	No mobile charge carriers (no ions, no delocalised electrons)
Thermal conductivity	Poor	No free electrons; phonon transport is inefficient
Solubility in water	Varies	Polar molecules with H-bonding (ethanol, glucose) soluble; non-polar (I₂, naphthalene) insoluble
Solubility in non-polar solvents	Often yes	Non-polar molecules dissolve in non-polar solvents (I₂ in hexane)
Density	Low to moderate	Loose packing; cavities between molecules

The principle: when a simple molecular solid melts or boils, you break only intermolecular forces (typically 5–40 kJ mol⁻¹), not the covalent bonds within molecules (150–1000 kJ mol⁻¹). This gives the characteristic low melting points.

2. Giant covalent (macromolecular) structures

A giant covalent solid contains a continuous 3D (or 2D layered) network of atoms held together by covalent bonds throughout the entire crystal. There are no discrete molecules — you can think of the whole crystal as one enormous "molecule". The OCR canonical examples are diamond, graphite, and silicon(IV) oxide.

Diamond — sp³ carbon, tetrahedral 3D network

Each C atom is covalently bonded to 4 others in a perfect tetrahedral arrangement (109.5° angles).
All four valence electrons of each C are in localised C-C σ bonds.
The result is a rigid 3D network of strong covalent bonds.

Diamond — sp³, tetrahedral 3D 4 covalent bonds per C, 109.5° No delocalised electrons → insulator m.p. ~3700 °C; hardest natural substance

Graphite — sp², 2D hexagonal layers

3 covalent bonds per C, 120° (in plane) 1 delocalised electron → conducts in plane Layers held by London forces only → slippery

Properties of diamond:

Very high m.p. (~3700 °C / 4000 K) — strong covalent bonds throughout.
Hardest natural substance — rigid 3D bonding resists deformation.
Insulator (electrically) — no delocalised electrons; all 4 valence electrons per C in localised σ bonds.
Insoluble in water and organic solvents — cannot break the lattice.
Excellent thermal conductor — phonons transmit efficiently through the rigid lattice. Surprisingly, diamond conducts heat better than most metals.

Uses: Cutting tools (drill bits, surgical blades), abrasives (sandpaper, grinding wheels), jewellery, heat sinks for high-power electronics.

Graphite — sp² carbon, 2D layered network with delocalised electrons

Each C atom is covalently bonded to 3 others in a planar hexagonal arrangement (120° angles, sp² hybridised).
The fourth valence electron per C is delocalised in a π-orbital system over the layer.
Adjacent layers are held together by weak London forces only (no covalent bonds between layers).

Properties of graphite:

Very high m.p. (~3650 °C) — strong covalent bonds within layers (you cannot easily melt graphite even though layers slide).
Soft and slippery — layers slide over each other easily because only weak London forces bind them.
Conducts electricity in plane — delocalised electrons free to move within each layer.
Does not conduct perpendicular to layers — no electron mobility between layers.
Lower density than diamond (2.3 vs 3.5 g cm⁻³) — layers are widely spaced.
Insoluble in water and organic solvents.

Uses: Pencil "lead" (graphite + clay), electrodes (esp. in electrolysis), lubricant (e.g. for door locks where oil would attract dust), brushes in electric motors.

Diamond vs graphite — the canonical OCR comparison

Property	Diamond	Graphite
Bonds per C	4 (sp³, tetrahedral)	3 (sp², trigonal planar)
Bond angle	109.5°	120°
Structure	3D tetrahedral network	2D hexagonal layers, weakly stacked
Delocalised electrons	None	1 per C (delocalised π system)
Conductivity	Insulator	Conducts in plane only
Hardness	Hardest natural substance	Soft, slippery
Density (g cm⁻³)	3.5	2.3
Melting point (°C)	~3700	~3650 (sublimes)
Cleavage	Crystalline, perfect	Layered, easily flaked

Both are allotropes of carbon — same element, different structural arrangement. The dramatic difference in properties arises entirely from the bonding pattern (sp³ vs sp²) and the consequent network connectivity.

Silicon(IV) oxide (silica, SiO₂)

Structure analogous to diamond: each Si atom is bonded to 4 O atoms, each O atom is bonded to 2 Si atoms, in a 3D tetrahedral network.
Found naturally as quartz, sand, and (with impurities) as glass.

Properties:

Very high m.p. (~1710 °C) — strong covalent Si-O bonds throughout.
Hard — rigid 3D bonding.
Insulator — no free electrons.
Insoluble — cannot disrupt the lattice without breaking covalent bonds.

Uses: Glass (silicates with Na, Ca, etc.), silicon wafers (after carbon reduction), abrasives.

Why does CO₂ form a molecular solid while SiO₂ forms a giant covalent solid?

A profound A-Level question with a deep answer. C and Si are both Group 14 elements. In CO₂ the C and O form two π bonds via favourable 2p-2p sideways overlap, giving discrete O=C=O molecules. In SiO₂ the Si 3p orbitals are too large to overlap effectively in π-bonding with O 2p, so Si forms only single bonds — and since each Si needs 4 bonds (to fill its valence shell) and each O needs 2 bonds, the result is a 3D network rather than discrete molecules. This is one of the most striking demonstrations of how subtle differences in orbital overlap can drive dramatically different macroscopic structures.

Graphene — a single layer of graphite

Single layers of graphite (one atom thick) are called graphene, isolated in 2004 by Geim and Novoselov (Nobel Prize 2010). Graphene has:

Tensile strength ~100× that of steel per unit weight.
Excellent electrical and thermal conductivity (better than graphite or copper).
High flexibility and transparency.

Graphene is an emerging material in flexible electronics, transparent conductors, batteries, and photovoltaics.

3. Giant metallic structures

A giant metallic solid consists of a 3D lattice of positive metal ions (cations) in a "sea" of delocalised valence electrons. The metallic bond is the electrostatic attraction between the cation cores and the delocalised electron sea.

The metallic bonding model

When metal atoms come together in the solid state, their outer (valence) electrons are no longer attached to specific atoms but become delocalised — free to move through the entire crystal. Each metal atom contributes its valence electrons (1 for Na, 2 for Mg, 3 for Al) to the sea, leaving behind a positive cation core.

Each Na atom contributes its 3s¹ electron to the delocalised sea (small blue dots) Na⁺ + e⁻ + e⁻ + .. = strong non-directional metallic bond

The metallic bond is non-directional (unlike covalent bonds, which point along specific axes). This non-directionality is what gives metals their malleability and ductility — layers of cations can slide past each other while the delocalised electron sea adapts seamlessly.

Properties of metals

Property	Observation	Explanation
Melting / boiling point	Mostly high (Na +98 °C; Fe 1538 °C; W 3422 °C)	Strong electrostatic attraction between cations and delocalised electrons
Electrical conductivity (solid AND liquid)	Excellent	Delocalised electrons free to move under applied potential difference
Thermal conductivity	Excellent	Delocalised electrons transfer kinetic energy quickly
Malleability	Good (can be hammered into sheets)	Layers slide; non-directional bond accommodates
Ductility	Good (can be drawn into wires)	Same mechanism
Lustre	Shiny when polished	Delocalised electrons reflect light at all visible frequencies
Density	Mostly high	Close packing of metal cations
Solubility in water	Insoluble (some react chemically; e.g. Na + H₂O)	Metallic bond does not dissolve; water cannot polarise metallic-bonded crystals

Period 3 metals — melting point trend

Metal	Configuration	Cation	Electrons donated	m.p. (°C)
Na	[Ne] 3s¹	Na⁺	1	98
Mg	[Ne] 3s²	Mg²⁺	2	650
Al	[Ne] 3s² 3p¹	Al³⁺	3	660

Across the period, the cation charge and the number of delocalised electrons per cation both increase, strengthening the metallic bond. The melting point rises sharply from Na to Mg and Al. (The trend breaks at Si onward, where the bonding type shifts to giant covalent (Si) and then simple molecular (P₄, S₈, Cl₂, Ar).)

Alloys

Alloys are metallic solids in which two or more metals (or a metal and a non-metal) are mixed at the atomic level. The substitution of one metal atom for another (e.g. Cu by Zn in brass, Fe by C in steel) disrupts the regular slip planes, making the alloy harder and stronger than either pure component. Pure metals are typically too soft for engineering use; almost all "metals" we use are actually alloys.

4. Giant ionic structures (review from Lesson 6)

Covered in detail in Lesson 6 — included here for the four-way comparison:

Structures: Simple Molecular, Giant Covalent and Metallic

Structures: Simple Molecular, Giant Covalent and Metallic

The four structural types — overview

1. Simple molecular structures

Examples

Properties of simple molecular solids

2. Giant covalent (macromolecular) structures

Diamond — sp³ carbon, tetrahedral 3D network

Graphite — sp² carbon, 2D layered network with delocalised electrons

Diamond vs graphite — the canonical OCR comparison

Silicon(IV) oxide (silica, SiO₂)

Why does CO₂ form a molecular solid while SiO₂ forms a giant covalent solid?

Graphene — a single layer of graphite

3. Giant metallic structures

The metallic bonding model

Properties of metals

Period 3 metals — melting point trend

Alloys

4. Giant ionic structures (review from Lesson 6)

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