Chemistry Revision: Bonding and Structure

Ionic Bonding

  • The loss or gain of electrons in order to fill the outer shell

Ionic Structures:

  • Ions are packed in a regular lattice arrangement
  • Very strong chemical bonds between all the ions
  • High melting points and boiling points, due to the strong chemical bonds
  • When in solution, they will conduct electricity as the ions are free to move
  • The same can be said for when they are molten
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Covalent Bonding: Simple Molecular substances

  • A shared pair of electrons in order to fill both outer shells

Important examples are: Hydrogen,Chlorine,Hydrogen Chloride,Ammonia,Methane,Water,Oxygen

Simple Molecular Substances

  • Very weak forces of attraction between the molecules
  • As a result the melting and boiling points are very low as the molecules are easily parted from each other.
  • Do not conduct electricity as there are no ions
  • Primarily gases/liquids at room temperature 
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Giant Covalent Structures

  • Similar to giant ionic structures, except no ions
  • Very high melting/boiling points
  • Usually insoluble in water


  • Each carbon atom has 4 covalent bonds in a very rigid structure, making diamond the hardest natural substance


  • Each carbon atom has 3 covalent bonds, creating layers which can slide over each other. They can also rub off (hence pencils) which leaves free electrons, making graphite a good non-metal electrical conductor. 
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Metallic Bonding

  • If the metal is in Group 1, there is 1 free electron, in Group 2 there are 2, and so on. 
  • Metallic bonding relies on free electrons, as the outer shells of the metals overlap, creating these free, or delocalised electrons.
  • There is electrostatic attraction between the cations and delocalised electrons. These delocalised electrons allow all metals to conduct electricity. 
  • They also "hold" the atoms in a regular structure
  • The atoms are also able to slide over each other, hence why pure metals are often very malleable and can be manipulated easily
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