Chemical Bonding (Ionic, Covalent and Metallic): Question 4
Syllabus 2.6
A specialist workshop cuts and polishes industrial glass panels using saw blades tipped with diamond, chosen because diamond is the hardest known natural material. In the same workshop, a furnace uses crucibles made from silicon dioxide, , to melt metal alloys at very high temperatures, and the moving parts of the cutting machine are lubricated with graphite powder.
(a) Describe the structure and bonding in diamond, and use it to explain why diamond is extremely hard. [3]
(b) Silicon dioxide has a giant covalent structure in which each silicon atom is bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms. Explain why silicon dioxide, like diamond, has a very high melting point. [2]
(c) Graphite has a different structure from diamond, even though both are forms of carbon.
(i) Describe the structure of graphite. [2]
(ii) Use the structure of graphite to explain why it is soft and slippery, and why it conducts electricity, unlike diamond. [3]
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Worked solution
Part (a): Structure and hardness of diamond
Diamond has a giant covalent structure: every carbon atom forms four strong covalent bonds to four neighbouring carbon atoms, building a rigid three-dimensional network that extends throughout the entire crystal.
To scratch or break diamond, a large number of these strong covalent bonds would have to be broken at the same time. Since covalent bonds require a great deal of energy to break, this makes diamond extremely hard (and gives it a very high melting point too).
Part (b): Why silicon dioxide also has a very high melting point
Silicon dioxide has the same kind of structure as diamond: a giant covalent lattice in which every silicon and oxygen atom is joined to its neighbours by strong covalent bonds, extending continuously through the whole structure. There are no separate, individual molecules.
To melt silicon dioxide, huge numbers of these strong covalent bonds throughout the lattice must be broken, which requires a very large amount of energy. This is why silicon dioxide, like diamond, has a very high melting point.
Part (c)(i): Structure of graphite
Graphite is also a giant covalent structure of carbon atoms, but arranged differently from diamond. The carbon atoms are arranged in flat layers of hexagonal rings. Within each layer, every carbon atom is covalently bonded to only three other carbon atoms (not four, as in diamond). The layers themselves are held together only by weak forces, much weaker than the covalent bonds within a layer.
Part (c)(ii): Explaining graphite’s properties
Soft and slippery: since the forces holding the layers together are weak, the layers can easily slide over one another without breaking the strong covalent bonds within each layer. This is why graphite feels slippery and works well as a lubricant.
Conducts electricity: because each carbon atom in graphite forms only three covalent bonds (using three of its four outer electrons), one outer electron per carbon atom is left over. This electron becomes delocalised, free to move along the layers rather than being fixed in a bond. These mobile, delocalised electrons allow graphite to conduct electricity.
Diamond, by contrast, has no delocalised electrons: every one of a carbon atom’s four outer electrons is used in a fixed covalent bond, so there are no free-moving charge carriers, and diamond does not conduct electricity.
Final answers
- (a) Giant covalent structure, each carbon bonded to 4 others; many strong covalent bonds must be broken, so diamond is extremely hard.
- (b) Also giant covalent, with strong bonds throughout; breaking many strong covalent bonds needs lots of energy, giving a very high melting point.
- (c)(i) Flat hexagonal layers, each carbon bonded to only 3 others; weak forces hold the layers together.
- (c)(ii) Weak forces let layers slide → soft/slippery; one delocalised electron per carbon atom → conducts electricity (unlike diamond, which has none).