Crystal structures
CIE A-Level Chemistry· Unit 4: States of Matter· 30 min read
1. Ionic Crystal Structures★★☆☆☆⏱ 15 min
Ionic Crystal
A giant crystalline lattice of oppositely charged ions held together by strong electrostatic ionic bonds.
Example:
Sodium chloride (NaCl), magnesium oxide (MgO)
Ions pack to maximize electrostatic attraction between opposite charges, forming regular lattices. Coordination number describes how many adjacent oppositely charged ions surround a given ion. For NaCl, coordination number is 6 for both ions.
Calculate the formula of NaCl from its unit cell, where Na⁺ is at edge centers and body center, and Cl⁻ is at corners and face centers.
- 1
Count and weight Na⁺ ions: 12 edge ions × 1/4 + 1 body ion = 3 + 1 = 4 total Na⁺ per unit cell
- 2
Count and weight Cl⁻ ions: 8 corner ions × 1/8 + 6 face ions × 1/2 = 1 + 3 = 4 total Cl⁻ per unit cell
- 3
Ratio of Na⁺:Cl⁻ = 4:4 = 1:1, so the formula is NaCl, matching expected value.
2. Metallic Crystal Structures★★☆☆☆⏱ 15 min
Metallic Crystal
A giant lattice of positively charged metal ions surrounded by a delocalized sea of free electrons, held together by metallic bonding.
Example:
Copper, iron, magnesium
Metals pack tightly to form dense lattices. Delocalized electrons move freely through the structure, which gives metals their characteristic thermal and electrical conductivity. Layers of metal ions can slide past each other without breaking bonding, making metals malleable and ductile.
Explain why metallic crystals conduct electricity but solid ionic crystals do not.
- 1
In metallic crystals, delocalized free electrons are present throughout the lattice. These electrons can move freely to carry electric charge when a potential difference is applied.
- 2
In solid ionic crystals, all ions are fixed in place in the lattice and cannot move. There are no free charged particles to carry charge.
- 3
This is why metals conduct electricity in solid state, but solid ionic crystals do not.
3. Covalent Network Crystals★★★☆☆⏱ 20 min
Covalent Network (Macromolecular) Crystal
A continuous giant crystal structure where all atoms are bonded to adjacent atoms by strong covalent bonds, with no discrete molecules.
Example:
Diamond, graphite, silicon(IV) oxide
Covalent network crystals have very high melting points because all bonds between atoms are strong covalent bonds that require large amounts of energy to break. Most are non-conductive, with graphite being a key exception.
Explain why graphite is soft and can be used as a lubricant, while diamond is very hard.
- 1
In diamond, each carbon atom bonds covalently to 4 other carbons in a rigid 3D tetrahedral network. No layers can slide, so diamond is extremely hard.
- 2
In graphite, carbon atoms bond to 3 other carbons to form flat hexagonal layers. Only weak intermolecular forces hold the layers together.
- 3
Layers can slide easily past each other when force is applied, making graphite soft and suitable as a lubricant.
4. Molecular Crystals★★☆☆☆⏱ 10 min
Molecular Crystal
A crystalline solid made of discrete small molecules, held together by weak intermolecular forces (not covalent bonds between molecules).
Example:
Ice, solid iodine, solid carbon dioxide
Molecular crystals have much lower melting points than other crystal types because weak intermolecular forces require little energy to overcome. All electrons are held in covalent bonds within molecules, so molecular crystals do not conduct electricity.
Iodine has strong covalent bonds within I₂ molecules, but its melting point is low. Explain why.
- 1
When solid iodine melts, only the weak intermolecular forces between separate I₂ molecules are broken.
- 2
The strong covalent bonds holding the two iodine atoms together within each I₂ molecule do not break.
- 3
Little energy is needed to overcome the weak intermolecular forces, so iodine has a low melting point despite strong intramolecular covalent bonds.
5. Common Pitfalls
Wrong move:
Claiming covalent bonds break when molecular crystals melt.
Why:
Students confuse intermolecular and intramolecular bonding, which loses easy marks.
Correct move:
Only weak intermolecular forces between molecules break when molecular crystals melt; covalent bonds inside molecules remain intact.
Wrong move:
Stating all covalent crystals cannot conduct electricity.
Why:
This ignores the common exception of graphite, which is regularly tested.
Correct move:
All covalent network crystals except graphite do not conduct electricity; graphite has delocalized electrons that allow conduction.
Wrong move:
Counting all ions in a unit cell as full ions to get the wrong formula.
Why:
Most ions are shared between multiple adjacent unit cells, so they must be weighted correctly.
Correct move:
Apply the fraction rule: corner = 1/8, edge = 1/4, face = 1/2, fully internal = 1, to get the correct ion ratio.
Wrong move:
Claiming solid ionic crystals conduct electricity.
Why:
Students mix up conductivity of solid and molten/dissolved ionic compounds.
Correct move:
Solid ionic crystals only have fixed ions, so they do not conduct electricity; they only conduct when molten or dissolved.
6. Quick Reference Cheatsheet
Crystal Type | Bonding Between Particles | Melting Point | Solid Conductivity | Example |
|---|---|---|---|---|
Ionic | Strong ionic (electrostatic) | High | No | NaCl |
Metallic | Metallic (delocalized electrons) | High | Yes | Copper |
Covalent Network | Strong covalent | Very High | No (except graphite) | Diamond |
Molecular | Weak intermolecular | Low | No | Ice |
When this came up on past exams
AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.
- 2022 · 21
Compare ionic and metallic crystal properties
- 2023 · 12
Describe the structure of diamond
- 2021 · 23
Explain metallic conductivity
Going deeper
What's Next
Understanding crystal structures is the foundation for explaining the physical properties of solids, a recurring exam theme that connects to bonding, energetics, and materials chemistry across AS and A-Level. You will build on this knowledge when studying polymers, ceramics, and nanomaterials in later A2 units. Mastering the distinctions between the four crystal types and their properties secures easy marks in multiple choice and structured questions, making this a high-value topic to review. Explore the linked topics below to extend your understanding.