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