4.2.5 Giant Covalent Structures

Giant Covalent Structures

Covalent bonds are bonds between nonmetals in which electrons are shared between the atoms
In some cases, it is not possible to satisfy the bonding capacity of a substance in the form of a molecule; the bonds between atoms continue indefinitely, and a large lattice is formed. There are no individual molecules and covalent bonding exists between all adjacent atoms
Such substances are called giant covalent substances, and the most important examples are C and SiO₂
Graphite, diamond, buckminsterfullerene and graphene are allotropes of carbon

Diamond is a giant lattice of carbon atoms
Each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of 109.5^o
The result is a giant lattice with strong bonds in all directions
Diamond is the hardest substance known
- For this reason it is used in drills and glass-cutting tools

The structure of diamond

In graphite, each carbon atom is bonded to three others in a layered structure
The layers are made of hexagons with a bond angle of 120^o
The spare electron is delocalised and occupies the space in between the layers
All atoms in the same layer are held together by strong covalent bonds, and the different layers are held together by weak intermolecular forces

The structure of graphite

Buckminsterfullerene is one type of fullerene, named after Buckminster Fuller, the American architect who designed domes like the Epcot Centre in Florida
It contains 60 carbon atoms, each of which is bonded to three others by single covalent bonds
The fourth electron is delocalised so the electrons can migrate throughout the structure making the buckyball a semi-conductor
It has exactly the same shape as a soccer ball, hence the nickname the football molecule

The structure of buckminsterfullerene

Some substances contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers. Graphene is an example
Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons
Graphene is one million times thinner than paper; so thin that it is actually considered two dimensional

The structure of graphene

Silicon(IV)oxide is also known as silicon dioxide, but you will be more familiar with it as the white stuff on beaches!
Silicon(IV)oxide adopts the same structure as diamond - a giant structure made of tetrahedral units all bonded by strong covalent bonds
Each silicon is shared by four oxygens and each oxygen is shared by two silicons
This gives an empirical formula of SiO₂

4-1-11-the-structure-of-silicon

The structure of silicon dioxide

Different types of structure and bonding have different effects on the physical properties of substances such as their melting and boiling points, electrical conductivity and solubility

Giant covalent lattices have very high melting and boiling points
- These compounds have a large number of covalent bonds linking the whole structure
- A lot of energy is required to break the lattice
The compounds can be hard or soft
- Graphite is soft as the forces between the carbon layers are weak
- Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds
- Graphene is strong, flexible and transparent which it makes it potentially a very useful material
Most compounds are insoluble with water
Most compounds do not conduct electricity however some do
- Graphite has delocalised electrons between the carbon layers which can move along the layers when a voltage is applied
- Graphene is an excellent conductors of electricity due to the delocalised electrons
- Buckminsterfullerene is a semi-conductor
- Diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom is involved in a covalent bond so there are no free electrons available

Characteristics of Giant Covalent Structures Table

Although buckminsterfullerene is included in this section it is not classified as a giant structure as it has a fixed formula, C₆₀