Chemistry Revision: Bonding and Structure

• 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

• 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 

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

Diamond

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

Graphite

• 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. 

• 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