Chemistry c5

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  • Created by: Spiderpig
  • Created on: 01-04-16 14:26

Properties of metals

Metals are made up of atoms which are arranged in a regular pattern packed closely together.Their particles are held by strong metallic bonds.

  • metals can conduct electricity due to free moving electrons which can carry charge around to different places and when a potential difference is applied they move together allowing an electric current to flow through the metal
  • solid at room temperature
  • High melting point and boiling points
  • not soluble
  • dense
  • hard
  • malleable ( can change their shape ; you can bend them without them breaking)
  • conduct heat
  • expand when heated
  • ductile ( can be stretched to make a wire)
  • shiny

Metals are shiny because when their free electrons come into contact with light, they vibrate producing their own light.The light is then reflected back ; creating a shiny ( lustrous) appearance.

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metallic bonding

A metallic bond is the electostatic force between free electrons and metal ions.Metallic bonds are strong which is why metals maintain their regular lattic structure and have high melting and boiling points ( because lots of energy is required to break the bonds holding the particles together)

The process

  • Solid metal atoms lose their free electrons from their outer shell and become positively charged.
  • A sea of 'free' delocalised (negative) electrons drift freely between metal atoms.
  • The attraction of the 'sea' of negaive and positively charged metal ions hold the structure together.
  • Overall the metal crystal is not charged because the negative charge on the electrons balances the total charge o the metal ions.
  • The electrons can move freely through the structure which allows an electric current to flow from one end of the wire to another.
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Giant covalent structures

Giant covalent structures contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds. The atoms are usually arranged into giant regular lattices - extremely strong structures because of the many bonds involved.

Examples of giant covalent structures and their properties

Silicon dioxide (Sio₂) - commonly known as glass. Silicon and oxygen are both non-metals so they make covalent bonds.

  • cannot conduct electricity
  • insoluble
  • high melting and boiling points
  • hard
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Simple covalent molecules

Simple covalent molecules only contain a few atoms held together by strong covalent bonds. An example is carbon dioxide (CO2), the molecules of which contain one atom of carbon bonded with two atoms of oxygen.

Simple covalent substances have low melting and boiling points.This is because although they have strong bonds between the atoms ; there are weak forces holding the molecules together.When the simple substance melts or boils the weak intermolecular forces break,not the strong covalent bonds therefore they have low melting and boiling points.

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Covalent bonding

Covalent bonding occurs between non-metals.It is when atoms share 1 or more pairs of electons to make them stable ( have a complete outer shell )

Covalently bonded substances fall into two main types: simple molecules and giant covalent structures.

Most non-metals elements are molecular and most of these consist of molecules with two atoms joined together.E.g. nitrogen(N₂) oxygen (O₂).

Most compounds formed between non-metals are also molecular. E.g. nitrogen dioxide (NO₂) and carbon dioxide (CO₂).

Properties of Simple covalent molecules

  • Low melting and boiling points
  • usually liquid or gas at room temperature
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Allotropes of Carbon

Allotropes - structurally different forms of the same element.


  • made from only carbon atoms
  • every carbon makes four covalent bonds to acheive a full outer shell
  • every carbon atom is bonded to four other carbon atoms
  • therefore the structure carries on growing


  • Hard to scratch
  • very high melting and boiling points
  • when diamonds are cut light can be reflected of the diamond to make them shiny
  • doesn't conduct electricity.
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Allotropes of Carbon


Graphite is made only from carbon atoms. Every carbon atom makes four covalent bonds to achieve a full outer shell. Every carbon atom is bonded to three other atoms.

The covalent bonds form a hexagon pattern. The hexagons are flat forming sheets which stack in layers. Between the sheets are delocalised electrons ( which are free to move ) allowing the thin layers to slide over each other making the substance soft and slippery ( also so as a pencil it can make marks on paper).


  • high melting and boiling points
  • insoluble
  • conducts electricity due to free moving electron
  • soft (used as pencil)
  • slippery (used as lubricant for machinary)
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difference between the two allotropes of carbon

Graphite is soft whereas diamond is hard because graphite has covalent  bonds which joins each carbon atom to three of its neighbours.The three bonds are distributed in two dimensions, making flat sheets of hexagons. Each atom has one outer shell electron which has not been used to make a covalent bond. This free electron can drift between the gaps in the layers.Therefore there is a weak bond holding the layers together so that it is soft. ( the weak forces allow layers to slide of each other allowing it to conduct electricity but diamond has no charged particles that are free to move so it is an insulator diamond has strong covalent bonds holding its carbon atoms together - this makes it hard to break and cut.

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The Earth's atmosphere

  • Litosphere- land
  • Hydrosphere - water
  • atmosphere - air

useful materials we extract  from the Lithosphere are : coal, oil, gas , aluminium ,zinc, iron, copper,diamond , graphite and building materials such as clay , stone and gravel.

Relative atomic mass = the mass of a atom compared to the mass of an atom of carbon which is giventhe value of 12.

Relative formula mass = the sum of the realtive atomic masses of all the atoms or ions shown its formula.

e.g. hydrogen atomic mass = 1 and oxygen = 16 so formula mass = 2 x 1 + 16 = 18

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