How does the structure of a solid explain its properties?
Giant ionic, giant covalent, simple molecular and metallic structures, and how bonding and structure explain melting point, conductivity, hardness and solubility.
A focused answer to WJEC A-Level Chemistry Unit 1, covering giant ionic, giant covalent, simple molecular and metallic lattices, and how each structure explains melting point, electrical conductivity, hardness and solubility.
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What this dot point is asking
WJEC wants you to describe the four main solid structures and to use bonding and structure to explain physical properties such as melting point, electrical conductivity, hardness and solubility.
The answer
The four structures
Explaining properties
Structure and conductivity
A summary helps:
| Structure | Melting point | Conducts? | Example |
|---|---|---|---|
| Giant ionic | High | Molten or aqueous only | |
| Giant covalent | Very high | Only graphite | Diamond, |
| Simple molecular | Low | No | , ice |
| Metallic | Variable, often high | Yes (solid and liquid) | Mg, Cu |
Diamond and graphite compared
Diamond and graphite are both giant covalent forms (allotropes) of carbon, yet behave very differently because of how the atoms are bonded. In diamond each carbon forms four covalent bonds in a rigid three-dimensional tetrahedral lattice, with no free electrons, making it extremely hard and an electrical insulator. In graphite each carbon forms only three covalent bonds in flat layers, leaving one delocalised electron per atom that can move along the layers, so graphite conducts electricity. Weak forces between the layers let them slide, making graphite soft and a good lubricant. The same element gives opposite properties purely through structure.
Predicting properties from structure
A reliable approach to any "explain the property" question is to identify the structure first, then reason from the forces involved. High melting point points to strong bonding throughout a giant structure; low melting point to weak intermolecular forces in a simple molecular solid. Electrical conduction requires mobile charge carriers, so ask whether there are free electrons (metals, graphite) or free ions (molten or dissolved ionic solids). Solubility in water suggests an ionic or polar molecular solid. Working from structure to property in this order keeps answers logical and complete.
Examples in context
Silicon dioxide in glass and electronics. is a giant covalent solid with a very high melting point and great hardness, used in glass, sandpaper and as the insulating layer in silicon chips. Metallic bonding and alloys. The delocalised electron model explains why metals conduct and are malleable, and why alloying (mixing ion sizes) hardens them by disrupting layer sliding, central to materials engineering.
Try this
Q1. State the type of structure present in solid iodine. [1 mark]
- Cue. Simple molecular.
Q2. Explain why magnesium has a higher melting point than sodium. [2 marks]
- Cue. has a higher charge and donates two delocalised electrons, so the metallic bonding is stronger than in sodium ().
Q3. Explain why simple molecular solids have low melting points. [1 mark]
- Cue. Only weak intermolecular forces are broken on melting, not the strong covalent bonds within molecules.
Q4. Explain why graphite is soft and slippery. [1 mark]
- Cue. Weak forces between its layers let them slide over one another.
Q5. State the property you would test to distinguish a metallic solid from a giant ionic solid. [1 mark]
- Cue. Electrical conductivity when solid (metals conduct; ionic solids do not).
Exam-style practice questions
Practice questions written in the style of WJEC exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
WJEC 20184 marksDiamond and graphite are both giant covalent forms of carbon. Explain why graphite conducts electricity but diamond does not, and why graphite is soft while diamond is hard.Show worked answer →
In graphite each carbon forms three covalent bonds in flat layers, leaving one delocalised electron per atom. These delocalised electrons move along the layers, so graphite conducts electricity.
In diamond each carbon forms four covalent bonds in a rigid tetrahedral lattice with no free electrons, so diamond does not conduct.
Graphite is soft because weak forces between layers let them slide over each other. Diamond is hard because every atom is locked by four strong covalent bonds in three dimensions.
Markers reward the delocalised electrons in graphite, the four bonds and no free electrons in diamond, and the layer-sliding versus rigid-network argument.
WJEC 20203 marksExplain why sodium chloride has a high melting point and conducts electricity when molten or in solution but not when solid.Show worked answer →
Sodium chloride is a giant ionic lattice held by strong electrostatic attraction between and ions, so a large amount of energy is needed to melt it, giving a high melting point.
When solid, the ions are fixed in the lattice and cannot move, so it does not conduct.
When molten or dissolved, the ions are free to move and carry charge, so it conducts.
Markers reward the strong electrostatic attraction for the melting point and free mobile ions for conduction.
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Sources & how we know this
- WJEC A-level Chemistry specification — WJEC (2015)