What are giant covalent and metallic structures, and how does metallic bonding explain the properties of metals?
Giant covalent structures such as diamond and graphite, metallic bonding, and how each structure explains the properties of the substance.
A focused answer to the WJEC GCSE Science Double Award Unit 5 topic on giant structures, covering giant covalent structures such as diamond and graphite, metallic bonding, and how each structure explains the properties of the substance.
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What this dot point is asking
WJEC Double Award Unit 5 wants you to describe giant covalent structures such as diamond and graphite, describe metallic bonding, and explain how each structure gives its properties.
Giant covalent structures
The main examples are diamond and graphite (both forms of carbon) and silicon dioxide (sand).
Diamond and graphite
Diamond and graphite are both made only of carbon but have very different properties because their structures differ:
- Diamond: each carbon atom forms four strong covalent bonds in a rigid 3D structure. This makes diamond very hard and gives it a very high melting point. It does not conduct electricity (no free electrons).
- Graphite: each carbon forms three bonds, making flat layers. The layers are held together by weak forces, so they can slide over each other, making graphite soft and slippery. Each carbon has one spare (delocalised) electron, so graphite conducts electricity.
Metallic bonding
The outer electrons of the metal atoms are free to move through the whole structure.
Properties of metals
- Conduct electricity and heat: the delocalised electrons are free to move and carry charge and energy.
- Malleable (can be bent and shaped): the layers of ions can slide over each other without breaking the metallic bonds.
- High melting points: the strong attraction between the ions and the sea of electrons needs a lot of energy to overcome.
Uses linked to structure
The structures explain why each material is used in particular ways. Diamond is used in cutting tools and drill tips because it is extremely hard. Graphite is used in pencils (the layers rub off) and as a lubricant and in electrodes because it is slippery and conducts. Metals are used in wiring (good conductors), saucepans (conduct heat) and structures (strong but malleable). Being able to link a use to the property, and the property to the structure, is a common application question that ties this whole module together.
Graphene and other carbon forms
Carbon has other useful structures worth knowing. Graphene is a single layer of graphite, just one atom thick; it is very strong, very light and an excellent conductor, making it useful in electronics and strong materials. Graphite and diamond, along with graphene, are all forms of the same element carbon but with very different properties because their structures differ. This is a clear example of how structure, not just the element, decides the properties of a substance.
Try this
Q1. Why is diamond very hard? [1 mark]
- Cue. Each carbon forms four strong covalent bonds in a rigid 3D structure.
Q2. What carries the charge when a metal conducts electricity? [1 mark]
- Cue. The delocalised (free) electrons.
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 style4 marksDiamond and graphite are both made only of carbon. Explain why diamond is very hard but graphite is soft and slippery.Show worked answer →
A Unit 5 explain question worth 4 marks. Reward: in diamond, each carbon atom forms four strong covalent bonds in a rigid 3D structure, so it is very hard (2); in graphite, each carbon forms three bonds, making layers that can slide over each other (because the forces between layers are weak), so it is soft and slippery (2). Markers credit four bonds and rigid for diamond, and layers that slide for graphite. A common error is to say graphite has weak covalent bonds (the covalent bonds are strong; the forces between layers are weak).
WJEC style4 marksExplain, using metallic bonding, why metals conduct electricity and can be bent into shape.Show worked answer →
A Unit 5 explain question worth 4 marks. Reward: metals have a structure of positive ions surrounded by a sea of delocalised (free) electrons (1); the free electrons can move through the metal, carrying charge, so metals conduct electricity (1); the layers of ions can slide over each other when a force is applied (1), so metals can be bent and shaped (are malleable) (1). Markers credit free electrons for conduction and sliding layers for malleability. A common error is to say metals conduct because of ions moving.
Related dot points
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A focused answer to the WJEC GCSE Science Double Award Unit 5 topic on ionic bonding, covering how electrons are transferred between metals and non-metals to form ions, the giant ionic lattice, and the properties of ionic compounds.
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A focused answer to the WJEC GCSE Science Double Award Unit 5 topic on covalent bonding, covering the sharing of electrons between non-metal atoms, simple molecules, and why simple molecular substances have low melting points and do not conduct.
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A focused answer to the WJEC GCSE Science Double Award Unit 5 topic linking structure and bonding to properties, covering how melting point, conductivity and hardness reveal whether a substance is ionic, simple molecular, giant covalent or metallic.