Giant covalent and metallic structures  

​​In a nutshell 

Giant covalent structures contain strong covalent bonds. Metallic structures consist entirely of metal atoms with delocalised electrons.

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Giant covalent structures

Giant covalent structures contain strong covalent bonds between atoms. Electrostatic forces between atoms makes these stronger than regular covalently bonded structures. Giant covalent compounds contain giant lattice structures. 

Example 1

Carbon can form a giant covalent structure known as diamond.  

Each carbon atom is bonded to four other carbon atoms. These are strong covalent bonds. Diamond is the hardest known carbon substance, with very high melting and boiling points. 

Note: Allotropes are different forms of the same element. Carbon can form three allotropes: diamond, graphite and graphene. 

Example 2 

Silicon dioxide can form a giant covalent structure.   

There are strong covalent bonds between the silicon and oxygen atoms. 

Properties of giant covalent substances

• Very high melting and boiling points, as high amounts of energy is needed to break the strong covalent bonds. 
  
• Giant covalent structures do not contain any charged or free particles, so they do not conduct electricity. 
  
• Good thermal conductors as vibrations can travel through the lattice structure. 
• Not soluble in water.
• Hard due to the strong covalent bonds. 
  

Note: The exceptions to the conductivity rule are graphite and graphene. These are giant covalent substances that will conduct electricity. 

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Graphite

Graphite is made from sheets of carbon atoms. Each carbon atom forms bonds with three other carbon atoms. There is a free electron for each carbon atom. This enables graphite to conduct electricity. 

Graphene 

Graphene is one layer of graphite and is two-dimensional. As graphene is made from graphite, it is also able to conduct electricity. Graphene is strong, transparent and light. 

Metallic bonding  

Metallic bonding consists of electrostatic forces of attraction between the positive metal ions and their delocalised electrons. A delocalised electron (2.) is an electron that is free to move around. Metal structures consist of layers of positive metal ions (1.)

Properties of metals   

Melting and boiling points

• Strong electrostatic forces between opposite charges that require high amounts of energy to break.
  
• Metals contain high melting and boiling points this makes them shiny at room temperature.
  
• The more delocalised electrons in the structure, the higher the melting and boiling point. Therefore, metals containing higher charged ions will have a higher melting and boiling point. 

​​Solubility 

• Not soluble in water.
  

Material strength

• The layers in a pure metal can slide past one another easily.
• Metals are malleable (the shape can be changed).

Conductivity

• The free delocalised electrons can carry an electrical charge, so metals conduct electricity.
  
• The free delocalised electrons can carry heat energy, so metals can conduct heat.