Giant Structures

Giant structures are extended, three-dimensional networks of atoms or ions held together by strong bonds throughout the entire structure. Unlike simple molecular substances, where weak intermolecular forces determine physical properties, the properties of giant structures are governed by the strong bonds within the lattice. This gives them characteristically high melting points and other distinctive properties.

There are two main types: giant ionic lattices and giant covalent structures.

---

Structure Type Identification Flowchart

When given data about an unknown substance, use this flowchart to identify the structure type:

flowchart TD
    A["Unknown substance"] --> B{"Melting point?"}
    B -->|"Low (< ~300°C)"| C["Simple molecular"]
    B -->|"High (> ~500°C)"| D{"Conducts electricity as solid?"}
    D -->|"Yes"| E{"Malleable?"}
    E -->|"Yes"| F["METALLIC"]
    E -->|"No (layered conductor)"| G["GIANT COVALENT - Graphite"]
    D -->|"No"| H{"Conducts when molten?"}
    H -->|"Yes"| I["GIANT IONIC"]
    H -->|"No"| J["GIANT COVALENT (e.g. diamond, SiO₂)"]
    C --> K{"Polar or H-bonding?"}
    K -->|"H bonded to N/O/F"| L["London + dipole + H-bonding"]
    K -->|"Polar, no H-bonding"| M["London + dipole-dipole"]
    K -->|"Non-polar"| N["London forces only"]

---

Giant Ionic Lattices

We covered ionic bonding in detail in Lesson 1, but here we revisit the structures with a focus on comparing the two key lattice types.

The NaCl Structure

In sodium chloride:

• Coordination number: 6:6 (each ion is surrounded by 6 of the opposite type)
• The ions arrange in a face-centred cubic lattice
• Na⁺ and Cl⁻ alternate at every position

The CsCl Structure

Caesium chloride has a different structure because Cs⁺ is much larger than Na⁺:

• Coordination number: 8:8 (each ion is surrounded by 8 of the opposite type)
• Cs⁺ sits at the centre of a cube of 8 Cl⁻ ions
• The larger Cs⁺ ion can accommodate 8 nearest neighbours (Na⁺ is too small)

The coordination number depends on the radius ratio — the ratio of the smaller ion to the larger ion. Larger ratios (where the ions are more similar in size) favour higher coordination numbers.

---

Giant Covalent Structures

Giant covalent structures (also called network covalent or macromolecular structures) consist of atoms bonded to each other by covalent bonds extending throughout the entire structure. There are no individual molecules — the whole sample is essentially one giant molecule.

Diamond

Diamond is a form (allotrope) of carbon with the following structure:

• Each carbon atom is bonded to 4 other carbon atoms in a tetrahedral arrangement
• All bonds are strong C–C single covalent bonds (sigma bonds)
• The bond angle is 109.5° (tetrahedral)
• The result is a rigid, three-dimensional network

Properties of diamond:

• Extremely high melting point (3550°C): Many very strong C–C covalent bonds must be broken to disrupt the lattice.
• Extremely hard: The rigid tetrahedral network resists deformation. Diamond is the hardest naturally occurring substance.
• Does not conduct electricity: All four outer electrons on each carbon are used in covalent bonds. There are no delocalised electrons or free ions to carry charge.
• Insoluble in all solvents: The strong covalent network cannot be disrupted by solvent interactions.

Graphite

Graphite is another allotrope of carbon with a very different structure:

• Each carbon atom is bonded to 3 other carbon atoms in a trigonal planar arrangement within flat layers
• The layers are called graphene sheets
• Within each layer, the bond angle is 120°
• Each carbon uses 3 of its 4 outer electrons for covalent bonds; the fourth electron is delocalised across the entire layer
• The layers are held together by weak London dispersion forces

Properties of graphite:

• Very high melting point (3652°C): Although the layers are weakly bonded, the strong covalent bonds within the layers still require enormous energy to break.
• Soft and slippery: The layers can slide over each other because only weak London forces hold them together. This is why graphite is used as a lubricant and in pencils.
• Conducts electricity (parallel to layers): The delocalised electrons within each layer can move and carry charge when a voltage is applied. Graphite only conducts along the plane of the layers, not perpendicular to them.
• Insoluble in water but can be dispersed: The strong covalent network prevents dissolution.

Why Graphite Conducts but Diamond Does Not

This is a crucial comparison. In diamond, all four outer electrons are localised in C–C sigma bonds — none are free to move. In graphite, each carbon only forms three sigma bonds, leaving one electron per atom delocalised across the layer. These delocalised electrons behave like the free electrons in a metal and can carry electrical charge.

Silicon Dioxide (SiO₂)

Silicon dioxide (silica) has a giant covalent structure similar to diamond:

• Each silicon atom is bonded to 4 oxygen atoms in a tetrahedral arrangement
• Each oxygen atom is bonded to 2 silicon atoms
• The formula SiO₂ represents the ratio, not a discrete molecule
• Strong Si–O covalent bonds extend throughout the structure

Properties:

• Very high melting point (1713°C) — strong covalent bonds throughout
• Hard but brittle
• Does not conduct electricity — no delocalised electrons
• Insoluble in water (but dissolves in HF and in hot concentrated NaOH)

Silicon dioxide is the main component of sand and quartz.

---

Comparing Giant Structures: Comprehensive Table

Property	Diamond	Graphite	SiO₂	NaCl	Copper
Structure type	Giant covalent	Giant covalent (layered)	Giant covalent	Giant ionic	Metallic
Melting point / °C	3550	3652	1713	801	1085
Hardness	Very hard	Soft (layers slide)	Hard, brittle	Hard, brittle	Malleable
Electrical conductivity	None	Yes (along layers)	None	Molten/dissolved	Yes (solid)
Bonding broken on melting	Covalent C–C	Covalent C–C	Covalent Si–O	Ionic	Metallic
Solubility in water	Insoluble	Insoluble	Insoluble	Soluble	Insoluble
Mobile charge carriers	None	Delocalised e⁻	None	Ions (when molten)	Delocalised e⁻

---

Graphene

Graphene is a single layer of graphite — a one-atom-thick sheet of carbon atoms arranged in a hexagonal pattern. It was first isolated in 2004 and has extraordinary properties:

• Strongest material ever measured: The covalent bonds within the sheet and the hexagonal arrangement create exceptional tensile strength
• Excellent electrical conductor: Delocalised electrons can move freely across the sheet (each carbon contributes one delocalised electron, as in graphite)
• Flexible and transparent: Being one atom thick, it is essentially a two-dimensional material
• Very high thermal conductivity: Delocalised electrons transfer thermal energy efficiently

Applications of graphene include electronics, composite materials, flexible screens, and filtration membranes.

---

Fullerenes and Nanotubes

Fullerenes

Fullerenes are closed cage-like structures of carbon atoms. The most famous is buckminsterfullerene (C₆₀), which consists of 60 carbon atoms arranged in pentagons and hexagons, like a football (soccer ball).

• Each carbon is bonded to 3 others (like graphite)
• The structure contains both hexagons and pentagons (the pentagons create the curvature)
• C₆₀ is a simple molecular substance (unlike graphite, it exists as discrete molecules)
• It dissolves in some non-polar solvents
• The hollow interior can encapsulate other atoms or small molecules (useful for drug delivery)

Carbon Nanotubes

Carbon nanotubes are cylinders of rolled graphene sheets. They can be single-walled (one layer) or multi-walled (concentric cylinders).

• Extremely strong (high tensile strength along the tube axis)
• Conduct electricity along their length (delocalised electrons from sp² carbon)
• Very high length-to-diameter ratio (essentially one-dimensional conductors)
• Applications include reinforcing composite materials, electronics, and drug delivery systems

---

Carbon Allotropes at a Glance

Allotrope	Bonding per C	Hybridisation	Delocalised e⁻?	Structure type	Conductivity
Diamond	4 C–C	sp³	No	Giant covalent	None
Graphite	3 C–C (layers)	sp²	Yes (in layers)	Giant covalent (layered)	Along layers
Graphene	3 C–C (sheet)	sp²	Yes (in sheet)	2D giant covalent	Excellent
C₆₀	3 C–C (cage)	sp²	Yes (within cage)	Simple molecular	Poor
Nanotube	3 C–C (cylinder)	sp²	Yes (along tube)	Cylindrical	Along tube

Common exam mistake: Students sometimes say "graphite has weak bonds." This is wrong. Graphite has very strong covalent bonds within its layers. It is the forces between layers that are weak (London forces). The slipperiness and softness are due to the weak interlayer forces, but the very high melting point is due to the strong intralayer covalent bonds that must be broken.

---

A-Level Deep Dive: Giant Structures

Spec mapping

Edexcel 9CH0 specification, Topic 2: Bonding and Structure, sub-topic 2.7 — giant ionic and giant covalent structures. Candidates are required to describe the structures of giant ionic lattices (e.g. NaCl, MgO) and giant covalent solids (diamond, graphite, silicon, silicon dioxide), and to explain physical properties — melting and boiling points, electrical conductivity in solid, molten and aqueous states, hardness, brittleness, and solubility — in terms of the structure and the bonding present.

Cross-paper assessment is unusually heavy here because structure-type reasoning is genuinely synoptic. Paper 1 Topic 4 (periodicity) revisits silicon dioxide on Period 3 — a giant covalent oxide with a very high melting point and effective insolubility in water — alongside the molecular oxides P₄O₁₀ and SO₃. Paper 1 Topic 15 (transition metals) brings back giant covalent and giant ionic motifs in the context of refractory ceramics: tungsten carbide (WC) and titanium carbide (TiC) are giant covalent compounds whose hardness and high melting points are exploited industrially. Paper 3 (synoptic) routinely sets data-response items in which candidates infer structure type from a small table of physical properties — exactly the skill that the worked example below rehearses.

(refer to the official Pearson Edexcel specification document for exact wording).

---

Worked example with full mark scheme

Question (8 marks):

The table below gives selected physical properties of four substances, A, B, C and D.

Substance	m.p. / °C	Conducts as solid?	Conducts when molten?	Soluble in water?
A	801	No	Yes	Yes
B	1610	No	No	No
C	1085	Yes	Yes	No
D	−7	No	No	No

Identify each substance as giant ionic, giant covalent, simple molecular, or metallic, and justify each identification using the data. (8)

Solution with mark scheme:

Substance A (m.p. 801 °C, non-conducting solid, conducting melt, water-soluble) — giant ionic.

M1 — recognises that a very high melting point combined with no conduction in the solid state is consistent with a giant ionic lattice rather than a metal. The reasoning must be explicit: "the very high melting point is due to strong electrostatic attraction between oppositely charged ions throughout the lattice, which requires a large amount of energy to overcome." Bare statement that "it's ionic because it dissolves" earns nothing.

A1 — explains conduction in the melt and aqueous solution by reference to mobile ions that are free to move and carry charge once the lattice has been disrupted. The candidate must connect why solid ionic compounds do not conduct (ions held in fixed positions by the lattice) with why the melt does (ions free to move).

Substance B (m.p. 1610 °C, non-conducting throughout, water-insoluble) — giant covalent.

M1 — identifies that a very high melting point with no electrical conductivity in any state and no water solubility points to a giant covalent network with no mobile charge carriers. Examples worth naming include silicon dioxide and diamond.

A1 — explains that to melt the substance many strong covalent bonds throughout the entire structure must be broken, and that there are no charged particles or delocalised electrons available to conduct electricity.

Substance C (m.p. 1085 °C, conducting in both solid and molten states, water-insoluble) — metallic.

M1 — recognises that conduction in the solid state is a fingerprint of metallic bonding (or, much more rarely, graphite — but graphite is layered and behaves differently in other respects).

A1 — attributes solid-state conductivity to delocalised electrons free to move throughout the lattice of positive metal ions, and notes that metals are typically insoluble in water because there is no favourable energy pathway to break apart the metallic lattice into hydrated species.

Substance D (m.p. −7 °C, non-conducting throughout, water-insoluble) — simple molecular.

M1 — recognises that a low melting point indicates that only weak intermolecular forces are being overcome on melting, not strong covalent or ionic bonds.

A1 — connects the absence of conductivity to the absence of mobile ions or delocalised electrons, and the limited water solubility to the non-polar (or weakly polar) nature of typical simple molecular substances.

Total: 8 marks (4 × M1 A1). Each identification requires both the correct label and a property-grounded reason. Bare labels score zero.

---

Specimen question modelled on the Edexcel 9CH0 Paper 1 format

Question (6 marks): Diamond and graphite are both pure carbon, yet diamond is one of the hardest known substances and a poor electrical conductor, while graphite is soft and conducts electricity along its layers. Compare and contrast the structures and bonding of diamond and graphite, and use these differences to explain the physical properties described. (6)

Mark scheme decomposition by AO:

• M1 (AO1.1) — diamond: each carbon atom is sp³ hybridised and forms four equal covalent bonds to four other carbon atoms in a tetrahedral arrangement, generating a three-dimensional giant covalent network in which all four valence electrons of every carbon are localised in σ bonds.
• M1 (AO1.2) — graphite: each carbon atom is sp² hybridised and forms three covalent bonds to three other carbons, generating planar hexagonal layers; the fourth electron occupies a p-orbital perpendicular to the plane and is delocalised into a π system shared across the layer.
• M1 (AO1.2) — graphite layers are held to one another by weak London (induced dipole–induced dipole) forces, with no covalent bonds between layers.
• M1 (AO2.1) — diamond is very hard because rigid covalent bonds extend in all three dimensions, so deforming the lattice requires breaking many strong σ bonds simultaneously; diamond is a poor electrical conductor because all valence electrons are tied up in localised σ bonds and there are no delocalised charge carriers.
• M1 (AO2.1) — graphite conducts in plane because delocalised π electrons in each layer are free to move parallel to the layer and carry charge; conductivity perpendicular to the layers is much poorer.
• M1 (AO2.2) — graphite is soft because adjacent layers can slide past one another; this requires only the weak London forces between layers to be overcome, not the strong covalent bonds within a layer.

Total: 6 marks split AO1 = 3, AO2 = 3. This is a balanced AO1/AO2 question — half the marks reward recall of the structural facts, half reward the explanatory step that links structure to property. Candidates who write "diamond has covalent bonds and graphite has covalent bonds" without distinguishing the two architectures cap at 1 or 2 marks.

---

Synoptic links

Connects to:

1. Topic 1 — atomic structure and the periodic table. The bonding capacity of an atom is set by its number of valence electrons. Carbon and silicon both sit in Group 14 with four valence electrons, and both can form four equivalent covalent bonds in tetrahedral arrangements — which is why both are capable of giant covalent networks (diamond and silicon respectively). Phosphorus, sulfur and chlorine, with five, six and seven valence electrons, do not form four equivalent bonds and so do not generate analogous three-dimensional networks; they prefer molecular allotropes such as P₄, S₈ and Cl₂.

2. Topic 4 — periodicity (Period 3 oxides). SiO₂ is giant covalent, with each Si bonded to four O and each O bridging two Si in a three-dimensional network. P₄O₁₀ and SO₃ are simple molecular oxides. The contrast — same period, neighbouring elements, very different oxide structure types — is precisely the diagnostic that periodicity questions exploit when they ask candidates to compare melting points across a row of oxides.

3. Topic 8 — energetics (Born–Haber cycles). The lattice energy of a giant ionic compound is the energy released when one mole of the solid forms from gaseous ions. Born–Haber cycles allow lattice energies to be calculated indirectly from enthalpies of atomisation, ionisation, electron affinity and formation. Trends in lattice energy across a series (e.g. NaF → NaCl → NaBr → NaI) illuminate why melting points of giant ionic solids fall as anion radius increases.

4. Topic 15 — transition metals. Several transition metal oxides (TiO₂, Al₂O₃, Cr₂O₃) and carbides (TiC, WC) form giant ionic or giant covalent networks. Their extreme hardness and very high melting points make them industrially important refractory ceramics, used in cutting tools and high-temperature furnaces. The same structural reasoning that explains diamond's hardness explains tungsten carbide's.