9.08.3 Giant Covalent Substances

Syllabus

Substances that consist of giant covalent structures are solids with very high melting points. All of the atoms in these structures are linked to other atoms by strong covalent bonds. These bonds must be overcome to melt or boil these substances. Diamond and Graphite (forms of Carbon) and Silicon Dioxide (silica) are examples of giant covalent structures.
Students should be able to recognise giant covalent structures from diagrams showing their bonding and structure.
In Diamond, each Carbon atom forms four covalent bonds with other Carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity.
Students should be able to explain the properties of diamond in terms of its structure and bonding.
In Graphite, each Carbon atom forms three covalent bonds with three other Carbon atoms, forming layers of hexagonal rings which have no covalent bonds between the layers.
In Graphite, one electron from each Carbon atom is delocalised.
Students should be able to explain the properties of graphite in terms of its structure and bonding.
Students should know that graphite is similar to metals in that it has delocalised electrons.

What does this mean?

Many covalent compounds form tiny molecules.

These molecules are weakly attracted to their neighbours by intermolecular forces.

But some covalent substances do not make small molecules.

Every atom is covalently bonded to every near-neighbour and a giant covalent lattice forms.

We only need to know in detail about Diamond and Graphite, which are both allotropes (different arrangements) of Carbon.

And also Silicon Dioxide - which is sand

Diamond

Every Carbon atom in diamond has 4 covalent bonds.

These are all very strong and so it would take a very high temperature to provide enough energy to break them all.

For this reason, diamonds have a very high melting point and are very hard (meaning they can scratch any other substance).

This makes diamonds ideal for drills and saws which get very hot due to friction and need to be able to scratch away at the surface of whatever they are scratching or drilling through.

All the outer electrons in each Carbon atom are involved in covalent bonds. They are not free to move. So Diamond is a very poor electrical conductor.

Other properties of diamond are that it is transparent and very lustrous, making gem-quality diamonds useful in jewellery.

Graphite

Every Carbon atom in diamond has 3 covalent bonds.

These are all very strong and so it would take a very high temperature to provide enough energy to break them all.

For this reason, graphite also has a very high melting point but graphite is not hard. In fact, it is soft enough to be used as a lubricant in places where oil can't be used.

Only 3 of Carbon's 4 outer electrons are involved in covalent bonds. The others are free to move. So Graphite is an electrical conductor for the same reason as metals are.

These delocalised electrons form weak forces between the layers, allowing them to slide over each other.

This explains why graphite is soft and why parts of your pencil slide onto the paper as you write.

1 pencil can write a 35 mile line as layers slide off the end

And pencils conduct because delocalised electrons can flow through the Graphite

3D Animation of Graphite and Diamond structure

To see the structure of Graphite in a 3D form click the link below

Silicon Dioxide

Carbon Dioxide (CO2) forms simple molecules with weak intermolecular forces and a low melting point.

But Silicon Dioxide forms a giant lattice with a similar structure to Diamond.

We generally know Silicon Dioxide as Sand - it is hard, difficult to melt and an electrical insulator for exactly the same reasons as Diamond - covalent bonds take a lot of energy to melt and there are no delocalised electrons.

That's why it's useful in sandpaper to smooth down wood and metals.

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