Giant Structures
Summary: Comparison of giant ionic, giant covalent, and giant metallic structures — their bonding, properties, and uses explained by structure. Diamond vs graphite is a classic exam question: both are allotropes of carbon with vastly different properties due to different bonding arrangements. Tags: igcse chemistry bonding Created: 2026-07-14 Last Updated: 2026-07-18
A giant structure is an arrangement of very large numbers of particles — atoms, ions, or metal cations with delocalised electrons — held together by strong bonding that extends continuously in all directions, forming a single Giant Lattice rather than discrete molecules. Because there are no individual molecules, the chemical formula of a giant structure (such as NaCl or SiO2) expresses only the simplest ratio of particles present, not a countable molecular unit. Giant structures fall into three families: the Giant Ionic Lattice of alternating oppositely charged ions found in salts, the giant covalent networks exemplified by Diamond, Graphite, and Silicon(IV) Oxide, and the Giant Metallic Lattice of positive ions in a sea of delocalised electrons. All three share characteristically high melting and boiling points, since melting requires overcoming strong forces — electrostatic attraction, covalent bonds, or metallic bonding — throughout the entire lattice. Where they differ sharply is in electrical conductivity, hardness, and solubility, and explaining these differences in terms of structure and bonding is the central skill examined in this topic. The contrast between diamond and graphite, two allotropes of the same element with radically different properties, is the classic illustration of how structure, not composition alone, determines a substance’s behaviour.
Overview of Giant Structures
A giant structure (also called a giant lattice) is one where the bonding extends in all directions throughout the entire substance. There are no separate molecules — the whole crystal is essentially one continuous bonded network.
There are three types of giant structures:
| Type | Examples | Particles | Bonding | Key Properties |
|---|---|---|---|---|
| Giant Ionic | NaCl, MgO | Positive and negative ions | Strong electrostatic forces between oppositely charged ions in all directions | High mp/bp; conducts when molten/aqueous; brittle; often soluble in water |
| Giant Covalent | Diamond, graphite, SiO2 | Atoms | Strong covalent bonds between atoms in all directions | Very high mp/bp; hard (except graphite); do not conduct (except graphite); insoluble |
| Giant Metallic | Cu, Fe, Al | Positive ions + delocalised electrons | Electrostatic attraction between ions and delocalised electrons | High mp/bp; conduct electricity and heat; malleable/ductile |
Giant Ionic Structures (e.g., NaCl)
Structure:
- A regular, alternating 3D lattice of positive and negative ions.
- In NaCl: each Na+ ion is surrounded by 6 Cl- ions; each Cl- ion is surrounded by 6 Na+ ions (6:6 coordination).
- The ions are held by strong electrostatic forces of attraction in ALL directions.
- The formula NaCl represents the ratio of ions (1:1) in the giant lattice, NOT a molecule.
Properties explained by structure:
| Property | Explanation |
|---|---|
| High melting point and boiling point (NaCl mp: 801 degrees C) | Strong electrostatic forces between oppositely charged ions in all directions require a great deal of energy to overcome. |
| Conducts electricity when molten or in aqueous solution | When molten or dissolved, the ions are free to move and can carry charge to electrodes. |
| Does NOT conduct electricity when solid | Ions are held in fixed positions in the lattice and cannot move. |
| Brittle (shatters when hit) | When a force causes layers of ions to slide, like-charged ions become adjacent (+ next to + or - next to -). The strong repulsion between like charges causes the crystal to shatter. |
| Often soluble in water | Water molecules can attract and separate the ions (hydration), overcoming the lattice energy. (Not all ionic compounds are soluble — solubility depends on relative lattice energy vs hydration energy.) |
Giant Covalent Structure: Diamond
Diamond is one of the two key allotropes of carbon. An allotrope is a different structural form of the same element in the same physical state.
Allotropes of carbon for IGCSE:
- Diamond: giant covalent, each C bonded to 4 others
- Graphite: giant covalent, each C bonded to 3 others, with layers
Structure of diamond:
- Each carbon atom is covalently bonded to 4 other carbon atoms.
- The bonds are arranged in a tetrahedral shape (bond angle 109.5 degrees).
- This forms a rigid, continuous 3D network extending through the entire crystal.
- The whole diamond crystal is essentially one giant molecule.
Properties of diamond explained by structure:
| Property | Explanation |
|---|---|
| Very hard (hardest natural substance) | Each C atom is bonded to 4 others by strong covalent bonds in a rigid tetrahedral network. To scratch diamond, many covalent bonds must be broken. Used in cutting tools, drill bits. |
| Very high melting point (>3500 degrees C) | All atoms are held by strong covalent bonds. To melt diamond, covalent bonds must be broken throughout the entire giant structure — this requires enormous energy. |
| Does NOT conduct electricity | All 4 outer electrons of each carbon atom are used in covalent bonds. There are no delocalised electrons and no free ions. Diamond is an electrical insulator. |
| Insoluble in water | Covalent bonds are too strong for water molecules to break. There is no attraction between water molecules and the carbon atoms. |
| High density | The tetrahedral network packs atoms efficiently in 3D. |
| Transparent, sparkles | Regular structure allows light to pass through and refract. |
Giant Covalent Structure: Graphite
Structure of graphite:
- Each carbon atom is covalently bonded to 3 other carbon atoms.
- The bonds are arranged in a trigonal planar shape (bond angle 120 degrees), forming flat hexagonal layers.
- Between the layers: there are weak intermolecular forces (London dispersion forces) — NOT covalent bonds.
- The fourth outer electron of each carbon atom is delocalised and free to move between the layers.
Properties of graphite explained by structure:
| Property | Explanation |
|---|---|
| Slippery/soft (can be used as a lubricant) | The layers of carbon atoms are held together by weak intermolecular forces, not covalent bonds. The layers can slide over each other easily. This is why graphite feels greasy and is used in pencil “lead” (layers rub off onto paper). |
| Very high melting point (>3600 degrees C — higher than diamond) | Within each layer, carbon atoms are held by strong covalent bonds. To melt graphite, these covalent bonds must be broken. The strong intralayer bonds give a very high mp. |
| Conducts electricity (along the layers) | Each C atom contributes 1 delocalised electron. These electrons are free to move within and between the layers, carrying electrical current. This makes graphite unusual — a non-metal that conducts electricity. Used in electrodes and electric motor brushes. |
| Low density | The relatively large gaps between layers make graphite less dense than diamond (graphite: 2.25 g/cm3 vs diamond: 3.51 g/cm3). |
| Insoluble in water | Covalent bonds are too strong for water to disrupt. |
| Opaque, black, shiny | Delocalised electrons absorb and re-emit visible light. |
Diamond vs Graphite — Classic Comparison
Both are allotropes of carbon (same element, different structural forms), yet they have dramatically different properties.
| Property | Diamond | Graphite | Reason for Difference |
|---|---|---|---|
| Hardness | Very hard (hardest natural material) | Soft, slippery | Diamond: 4 bonds per C, rigid 3D tetrahedral network. Graphite: 3 bonds per C, layered structure with weak forces between layers. |
| Electrical conductivity | None (insulator) | Good conductor along layers | Diamond: all electrons in covalent bonds. Graphite: 1 delocalised electron per C atom. |
| Melting point | Very high (~3550 degrees C) | Very high (~3650 degrees C) | Both have strong covalent bonds within their structures (C-C bonds in diamond; C-C bonds within graphite layers). Both require enormous energy to melt. |
| Appearance | Transparent, sparkling | Black, opaque, shiny | Diamond: regular tetrahedral arrangement, light passes through. Graphite: delocalised electrons absorb light. |
| Density | Higher (3.51 g/cm3) | Lower (2.25 g/cm3) | Diamond: tightly packed 3D network. Graphite: large gaps between layers. |
| Uses | Cutting tools, drill bits, jewellery | Pencil lead, lubricant, electrodes | Properties determine uses — hardness for cutting; conductivity and slipperiness for graphite applications. |
Giant Covalent Structure: Silicon Dioxide (SiO2)
Silicon dioxide (silica) is the main component of sand and quartz.
Structure:
- Silicon dioxide has a structure similar to diamond.
- Each Si atom is covalently bonded to 4 O atoms (tetrahedral).
- Each O atom is covalently bonded to 2 Si atoms.
- This forms a rigid 3D network of alternating Si and O atoms.
- The formula SiO2 represents the ratio of atoms (1 Si : 2 O) in the giant lattice — there are no SiO2 molecules.
Properties:
| Property | Explanation |
|---|---|
| Very high melting point (~1710 degrees C) | Covalent Si-O bonds throughout the giant structure require enormous energy to break. |
| Hard | The rigid 3D network of strong covalent bonds resists deformation. Sandpaper and grinding wheels use silicon dioxide/silicon carbide abrasives. |
| Does NOT conduct electricity | All electrons are in covalent bonds — no delocalised electrons, no ions. Pure SiO2 is an excellent insulator. |
| Insoluble in water | Covalent bonds are too strong for water molecules to disrupt. |
| Main component of sand, quartz, flint | These are naturally occurring forms of SiO2. |
SiO2 is a covalent compound despite containing a metal-like element (Si). Silicon is a metalloid, and Si-O bonds are polar covalent bonds. It is classified as giant covalent.
Giant Metallic Structures
Covered in detail in Metallic Bonding. Key points for comparison:
- Structure: regular lattice of positive metal ions surrounded by a sea of delocalised electrons.
- Properties: high mp/bp (varies with number of delocalised electrons); conduct electricity and heat; malleable and ductile; insoluble.
- Contrast with ionic: metals are malleable (layers slide, bonding maintained); ionic compounds are brittle (layers slide → like charges repel).
Summary Comparison Table
| Giant Ionic | Giant Covalent (Diamond) | Giant Covalent (Graphite) | Giant Metallic | |
|---|---|---|---|---|
| Particles | + and - ions | Atoms | Atoms | + ions + e- |
| Bonding | Ionic (electrostatic) | Covalent | Covalent + delocalised e- | Metallic |
| Melting/boiling point | High | Very high | Very high | High (varies) |
| Hardness | Hard (but brittle) | Very hard | Soft/slippery | Malleable/ductile |
| Conducts electricity (solid)? | No | No | Yes (along layers) | Yes |
| Conducts when molten/aqueous? | Yes | — | — | Yes (already solid) |
| Soluble in water? | Often yes | No | No | No |
Key Points
- Three types of giant structures: ionic, covalent, metallic
- Giant ionic: alternating + and - ions, strong electrostatic forces, high mp, brittle, conducts only when molten/aqueous
- Diamond: each C bonded to 4 others (tetrahedral), very hard, very high mp, does NOT conduct (insulator), all electrons in bonds
- Graphite: each C bonded to 3 others (trigonal planar), hexagonal layers, soft/slippery (weak forces between layers), very high mp, CONDUCTS electricity (1 delocalised electron per C atom)
- SiO2: similar to diamond, each Si bonded to 4 O, each O bonded to 2 Si, very high mp, does not conduct, main component of sand/quartz
- Diamond vs graphite: SAME element (C), DIFFERENT bonding/structure → DIFFERENT properties
- Ionic compounds brittle: layers slide → like charges repel → shatters
- Metals malleable: layers slide → delocalised electrons move with ions → bonding maintained
Key Concepts from Past Papers
- Allotrope: different structural forms of the same element in the same physical state
- Giant structure / giant lattice: a structure in which bonding extends in all directions throughout the entire substance; there are no separate molecules
- Delocalised electrons: electrons that are not associated with any single atom or bond and are free to move through the structure
- In diamond, each carbon atom is bonded to 4 other carbon atoms in a tetrahedral arrangement
- In graphite, each carbon atom is bonded to 3 others in hexagonal layers
- Graphite conducts electricity because each carbon atom contributes one delocalised electron
Keywords from Past Papers
graphite, reacts, calcium, oxide, diamond, platinum, giant, covalent, down, forms, silicon, silica, high, temperature, catalyst
Related Notes
- Atomic Structure
- Ions and Ionic Bonds
- Covalent Bonding
- Metallic Bonding
- Diamond
- Graphite
- Silicon(IV) Oxide
- Giant Lattice
- Giant Ionic Lattice
- Giant Metallic Lattice
Sources
- OpenStax Chemistry 2e — Chapter 10: Liquids and Solids, Rice University (free, CC BY 4.0)
- BBC Bitesize GCSE Chemistry — Giant Covalent Structures, BBC (free educational resource)
- Cambridge IGCSE Chemistry 0620 — Syllabus Section 2.6, Cambridge Assessment International Education
- CK-12 Chemistry for High School — Giant Covalent Structures, CK-12 Foundation (free, CC BY-NC 3.0)
Past Paper Sources
- 0620/32 Feb/March 2017: Q22(b)(iii) (1m), Q22(d)(i) (1m)
- 0620/32 Feb/March 2020: Q22(b)(ii) (3m)
- 0620/32 Feb/March 2021: Q77(b)(i) (1m)
- 0620/32 Feb/March 2022: Q66(c)(i) (1m)
- 0620/32 May/June 2020: Q55(b)(ii) (1m)
- 0620/33 May/June 2021: Q88(b)(iii) (1m)
- 0620/33 May/June 2024: Q33(a)(ii) (1m)
- 0620/33 October/November 2015: Q44(b)(i) (2m)
- 0620/33 October/November 2017: Q44(f)(i) (0m)
- 0620/33 October/November 2023: Q55(c)(ii) (1m)
- 0971/31 May/June 2021: Q44(c)(i) (2m)
- 0971/32 May/June 2020: Q55(b)(ii) (1m)
Common Misconceptions
| Misconception | Reality |
|---|---|
| ”Diamond and graphite are different elements” | Both are allotropes of carbon — the same element in different structural forms. The difference in properties is entirely due to different bonding arrangements. |
| ”Graphite has a low melting point because it is soft” | Graphite has a very high melting point (even higher than diamond’s) because the covalent bonds within each layer are very strong. Softness is due to weak forces between layers — a different type of bond. |
| ”Diamond conducts electricity because it’s shiny” | Diamond is an electrical insulator — all electrons are locked in covalent bonds. Graphite conducts because of delocalised electrons between layers. |
| ”SiO2 has simple molecules because of its formula” | The formula SiO2 represents the ratio of atoms (1:2) in the giant covalent lattice. There are no separate SiO2 molecules — the entire crystal is one continuous network. |
| ”NaCl is a molecule” | NaCl is a giant ionic lattice. The formula represents the ratio. There is no “molecule of NaCl” — each Na+ is surrounded by 6 Cl- and vice versa. |
| ”All ionic compounds are soluble” | Many are, but some (e.g., AgCl, BaSO4) are almost completely insoluble. Solubility depends on the balance between lattice energy and hydration energy. |