Giant Covalent Structures (Diamond, Graphite, Silicon Dioxide)

Giant Covalent Structures Overview

Giant covalent structures are substances where atoms are bonded together by strong covalent bonds in a continuous network throughout the material. Unlike simple molecules, these structures do not consist of individual molecules but form a giant lattice of atoms.

Because of the many strong covalent bonds, giant covalent structures have very high melting and boiling points. A lot of energy is needed to break these bonds.

Common examples of giant covalent structures include:

• Diamond
• Graphite
• Silicon dioxide (quartz)

Diamond Structure and Properties

In diamond, each carbon atom forms four strong covalent bonds with four other carbon atoms in a tetrahedral arrangement. This creates a very rigid, three-dimensional giant covalent structure.

Key properties of diamond:

• Very hard and rigid: The strong covalent bonds in all directions make diamond extremely hard, used in cutting tools and jewellery.
• High melting point: It requires a large amount of energy to break the many covalent bonds, so diamond melts at very high temperatures.
• Does not conduct electricity: All four outer electrons of each carbon atom are used in bonding, so there are no free electrons or ions to carry charge.

For instance, diamond’s hardness is why it is used in drill bits and saws.

Example: Calculate the number of covalent bonds formed by 1 mole of carbon atoms in diamond.

Each carbon atom forms 4 covalent bonds. 1 mole contains $6.02 \times 10^{23}$ atoms.

Total bonds = $4 \times 6.02 \times 10^{23} = 2.41 \times 10^{24}$ bonds.

Graphite Structure and Properties

Graphite is another form of carbon but has a very different structure. Each carbon atom is bonded to three others in flat layers of hexagonal rings. These layers are held together by weak forces, allowing them to slide over each other.

Key properties of graphite:

• Layers slide easily: The weak forces between layers let them slide, making graphite soft and slippery. This is why graphite is used as a lubricant and in pencils.
• Conducts electricity: Each carbon atom has one electron not used in bonding, which becomes delocalised and free to move, allowing graphite to conduct electricity.
• High melting point: Strong covalent bonds within layers require a lot of energy to break, so graphite has a very high melting point.

The delocalised electrons in graphite are free to move and carry charge, which is why graphite conducts electricity, unlike diamond.

Example: Explain why graphite conducts electricity but diamond does not.

In graphite, each carbon atom bonds to three others, leaving one electron free to move (delocalised electrons). These free electrons carry charge. In diamond, all four electrons are used in bonds, so no free electrons are available to conduct electricity.

Silicon Dioxide (Quartz) Structure and Properties

Silicon dioxide, also called quartz, is a giant covalent structure made of silicon and oxygen atoms. Each silicon atom is covalently bonded to four oxygen atoms in a tetrahedral shape, and each oxygen atom links two silicon atoms, forming a giant 3D network. Quartz is a common mineral found in sand and rocks.

Key properties of silicon dioxide:

• Network of strong covalent bonds: The structure is very strong and stable.
• High melting point: Like diamond and graphite, silicon dioxide has a very high melting point due to the many strong covalent bonds.
• Does not conduct electricity: There are no free electrons or ions in the structure to carry charge.

Example: Why does silicon dioxide have a higher melting point than many other compounds?

Because it has a giant covalent structure with many strong covalent bonds throughout the lattice, a large amount of energy is needed to break these bonds and melt the substance.