Chapter 7: Allotropes of Carbon and Metallic Bonding
7.1-Allotropes of Carbon
Diamond is very hard
- 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
Graphite contains sheets of hexagons
- 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 one layer of graphite
- 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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