Why do diamond, graphite and silicon dioxide behave so differently?
Giant covalent structures, the structures and properties of diamond, graphite and silicon dioxide, and how bonding explains hardness, melting point and electrical conductivity.
A CCEA GCSE Chemistry answer on giant covalent structures, covering the structures of diamond, graphite and silicon dioxide, and how their covalent bonding explains very high melting points, hardness, and why graphite conducts electricity while diamond does not.
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What this dot point is asking
CCEA wants you to describe the giant covalent structures of diamond, graphite and silicon dioxide, and explain their physical properties (very high melting points, hardness, and electrical conductivity) in terms of their bonding.
What makes a giant covalent structure
Because melting or boiling means breaking strong covalent bonds throughout the structure, giant covalent substances have very high melting and boiling points. This is the key shared property of diamond, graphite and silicon dioxide.
Diamond
Every atom is locked by four strong bonds in all directions, so diamond is the hardest natural substance and has an extremely high melting point. Because all four outer electrons are used in bonding, there are no free electrons, so diamond does not conduct electricity. Its hardness makes it ideal for cutting tools and drill tips.
Graphite
This layered structure explains two contrasting properties. The weak forces let the layers slide, so graphite is soft and slippery and is used as a lubricant and in pencils. The delocalised electrons are free to move, so graphite conducts electricity and is used for electrodes.
Silicon dioxide
Silicon dioxide (silica, the main part of sand and quartz) is a giant covalent lattice in which each silicon atom is bonded to four oxygen atoms and each oxygen to two silicon atoms. Like diamond it is hard, has a very high melting point and does not conduct, because all the bonds are strong covalent bonds with no free electrons. It is used in glass and as an abrasive.
Worked example
Examples in context
Example 1. Diamond-tipped cutting tools. Industrial saws and drills use diamond edges because the rigid 3D lattice resists being scratched or worn. The hardness that makes diamond a prized gem also makes it the material of choice for cutting hard stone and metal.
Example 2. Graphite electrodes in industry. Aluminium extraction and many electrolysis cells use graphite electrodes because graphite conducts electricity yet withstands high temperatures. Its rare combination of conductivity and a high melting point, both from its layered structure, makes it ideal.
Try this
Q1. State why diamond has a very high melting point. [1 mark]
- Cue. Many strong covalent bonds in a giant lattice must be broken.
Q2. Explain why graphite is used as a lubricant. [2 marks]
- Cue. Its layers are held by weak forces, so they slide over each other.
Exam-style practice questions
Practice questions written in the style of CCEA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
CCEA 20194 marksDiamond and graphite are both forms of carbon. Explain why diamond is very hard but graphite is soft and slippery.Show worked answer →
Markers want each property linked to the structure.
In diamond each carbon atom is covalently bonded to four others in a rigid 3D giant lattice. Every atom is held firmly by strong covalent bonds in all directions, so diamond is extremely hard.
In graphite each carbon is bonded to only three others, forming layers of hexagonal rings. The layers are held together only by weak forces between layers, so the layers can slide over each other, making graphite soft and slippery.
Markers reward four bonds in a rigid 3D lattice for diamond's hardness, and layers held by weak forces that slide for graphite being soft.
CCEA 20213 marksExplain why graphite conducts electricity but diamond does not.Show worked answer →
The marks are for the free electrons in graphite and their absence in diamond.
In graphite each carbon atom uses only three of its four outer electrons in bonding. The fourth electron is delocalised (free to move) between the layers.
These delocalised electrons can move through the structure and carry charge, so graphite conducts electricity.
In diamond all four of each carbon's electrons are used in covalent bonds, so there are no free electrons to carry charge, and diamond does not conduct.
Markers reward one delocalised electron per carbon in graphite that carries charge, and all four electrons bonded in diamond so none are free.
Related dot points
- Covalent bonding as the sharing of electron pairs between non-metal atoms, drawing dot-and-cross diagrams for simple molecules, and the properties of simple molecular substances.
A CCEA GCSE Chemistry answer on covalent bonding, covering how non-metal atoms share pairs of electrons to fill their outer shells, dot-and-cross diagrams for molecules such as hydrogen, water, ammonia and methane, and why simple molecular substances have low melting points and do not conduct.
- Ionic bonding as the transfer of electrons to form charged ions, drawing dot-and-cross diagrams, the giant ionic lattice, and how the structure explains the properties of ionic compounds.
A CCEA GCSE Chemistry answer on ionic bonding, covering how electrons transfer from metals to non-metals to form ions, dot-and-cross diagrams, the giant ionic lattice, and how this structure explains the high melting points, conductivity and solubility of ionic compounds.
- Metallic bonding as positive ions in a sea of delocalised electrons, how this explains the properties of metals, and why alloys are harder than the pure metal.
A CCEA GCSE Chemistry answer on metallic bonding, covering the model of positive ions in a sea of delocalised electrons, how it explains conductivity, malleability and high melting points, and why alloys are harder than the pure metal.
- Nanoparticles and nanoscience, the large surface area to volume ratio of nanoparticles, their uses, and the benefits and risks of nanotechnology.
A CCEA GCSE Chemistry answer on nanoparticles, covering their size, the large surface area to volume ratio that makes them so reactive and useful, their applications in sunscreens, catalysts and medicine, and the benefits and possible risks of nanotechnology.
- Electron arrangement in shells for the first 20 elements, writing electron configurations, and the link between outer-shell electrons, the group number and chemical reactivity.
A CCEA GCSE Chemistry answer on electron arrangement, covering how electrons fill shells for the first 20 elements, how to write electron configurations, and how the number of outer-shell electrons links to the group, the period and the chemical reactivity of an element.
Sources & how we know this
- CCEA GCSE Chemistry specification (1110) — CCEA (2017)