Chapter 7: Allotropes of Carbon and Metallic Bonding

7.1-Allotropes of Carbon

Diamond has a giant covalent structure, made up of carbon atoms that each form four covalent bonds.
- This makes diamond really hard
Those strong covalent bonds take a lot of energy to break and give diamonds a very high melting point
It doesn’t conduct electricity because it has no free electrons or ions

In graphite, each carbon atom only forms three covalent bonds creating sheets of carbon atoms arranged in hexagons
There aren’t any covalent bonds between the layers-they’re only held together weakly, so they’re free to move over
- This makes graphite soft and slippery, so it’s ideal as a lubricating material
Graphite’s got a high melting point-the covalent bonds in the layers need loads of energy to break
Only three out of each carbon’s four covalent electrons are used in bonds, so each carbon atom has one electron that’s delocalised(free) and can move
So graphite conducts electricity and thermal energy

Graphene is a sheet of carbon atoms joined together in hexagons
The sheet is just one atom thick, making it two-dimensional compound
The network of covalent bonds makes it very strong
- It’s also incredibly light, so can be added to composite materials to improve their strength without adding much weight
Like graphite, it contains delocalised electrons so can conduct electricity through the whole structure
- This means it has the potential to be used in electrons

Fullerenes are molecules of carbon, shaped like closed tubes or hollow balls
They’re mainly made up of carbon atoms arranged in hexagons, but can also contain pentagons(rings of five carbon atoms) or heptagons(rings of seven carbon atoms).
Fullerenes can be used to ‘cage’ other molecules
- The fullerenes structure forms around another atom or molecule, which is then trapped inside
- This could be used to deliver a drug into the body
Fullerenes have a surface area, so they could make great industrial catalysts, individual catalyst molecules could be attached to the fullerenes

Metals also consist of a giant structure
The electrons in the outer shell of the metal atoms are delocalised
- There are strong forces of electrostatic attraction between the positive metal ions and the shared negative electrons
These forces of attraction hold the atoms together in a regular structure and are known as metallic bonding
- Metallic bonding is very strong
Substances that are held together by metallic bonding include metallic elements and alloys
It’s the delocalised electrons in the metallic bonds which produce all the properties of metals

The electrostatic forces between the metal atoms and the delocalised sea of electrons are very strong, so need lots of energy to be broken
This means that most compounds with metallic bonds have very high melting and boiling points, so they’re generally solid at room temperature

The delocalised electrons carry electrical current and thermal energy through the whole structure, so metals are good conductors of electricity and heat

The layers of atoms in a metal can slide over each other, making metals malleable, this means that they can be bent or hammered or rolled into flat sheets

Pure metals often aren’t quite right for certain jobs, they’re often too soft when they’re pure so are mixed with other metals to make them harder
Most of the metals we use everyday are alloys, a mixture of two or more metals or a metal and another element
- Alloys are harder and so more useful than pure metals
- Different elements have different sized atoms, so when another elements is mixed with a pure metal, the new metal will distort the layers of metal atoms, making it more difficult for them to slide over each other

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