C3 - Structure and Bonding

• The particle model assumes:


  • there are no forces between the particles
  • that all particles in a substance are spherical
  • that all the spheres are sollid
• • The amount of energy needed to change the state of a substance depends on the forces between the particles
  • The stonger the forces between the particles, the higher the melting or boiling point of the substance
• Covalent bonding


  • Bonds between non-metal atoms
  • When electrons are shared between non-metal atoms
  • The number of electrons shared depends on how many extra electrons an electrons needs to make a full outer shell
• Single covalent bond


  Each atom shares one pair of electrons
• Double bond


  Each atom shares two pairs of electrons
• Most covalent structures do not conduct electricity because they do not have delocalised electrons that are free enough to carry charge
• The three types of covalent structure


  • Giant covalent
  • e.g. diamond
  • Small molecules
  • e.g. water
  • Large molecules
  • e.g. polymers
• Giant covalent structures


  • Billions of atoms, each with a strong covalent bond to a number of others
  • e.g. diamond
• Small covalent molecules


  • Only a few atoms with strong covalent bonds
  • Weak intermolecular forces
  • e.g. water
• Large covalent molecules


  • Many repeating units joined by covalent bonds to form a chain
  • Intermolecular forces are stronger than small molecules
  • e.g. polymers
• Properties of giant covalent structures


  • High melting and boiling points
  • Strong covalent bonds need more energy to break
  • Solid at room temperature
• Properties of small molecules

  • Low melting and boiling points
  • only weak intermolecular forces need to be overcome which does not require a lot of energy
  • Normally a gas or liquid at room temperature
• Properties of large molecules


  • Higher melting and boiling points than small molecules, but lower than giant covalent structures
  • Stronger intermolecular forces require more energy to overcome
  • Normally solid at room temperature
• Graphite


  • Structure
  • Giant covalent structure
  • Made only of carbon
  • Each carbon atom bonds to three others and forms hexagonal rings in layers
  • Each carbon atom has one delocalised electron
  • Hardness
  • The layers are not covalently bonded so they can slide
  • Therefore it is softer than diamond
  • Conductivity
  • The delocalised electrons can move
  • They can carry charges and therefore can conduct electricity
• Graphene


  • A single layer of graphite
  • Strong covalent bonds mean it is strong and can conduct electricity
  • Used in composites and high tech electronics
• Fullerenes


  • Hollow cages of carbon atoms bonded in one molecule
  • Can be a sphere or tube
  • Intermolecular forces are wek so they can slide over each other
• Fullerene spheres


  • Buckminsterfullerene was the first fullerene to be discovered, and has 60 carbon atoms
  • Fullerenns wwith different numbers of carbons exist with rings that form hollow shapes
  • Fullerenes like this are used as lubricants and in drug delivery
• Nanotubes


  • Carbon atoms arranged in cylindrical tubes
  • Their high tensile strength makes them useful in electronics
• Ions


  • Atoms can gain or lose electrons to give them a full outer shell
  • The number of protons is then different to the number of neutrons
  • The resulting particle is charged and called an ion
• Ionic bonding


  • Metal atoms react with non-metal atoms, and transfer electrons to the non-metal atom
  • The metal atoms do NOT share the electrons
• In ionic bonding, metals lose electrons to become positive ions
• In ionic bonding, non-metals gain electrons to become negative ions
• Giant ionic lattices 


  • When metal atoms transfer electrons to non-metal atoms, positive and negative ions are left
  • These are attracted to each other by strong electrostatic forces of attraction
  • This is ionic bonding
• Melting points of ionic substances


  • Ionic substances have high melting points
  • This is because the electrostatic force of attraction is strong and requires a lot of energy to break
• Conductivity of ionic substances


  • Solid ionic substances fo not conduct electricity because the ions are fixed in position and are not free to carry charge
  • When melted or dissolved in water, ionic substances do conduct electricity, because the ions are free to move and carry charge
• Metals


  • The atoms that make up metals form layers
  • The electrons in the outer shells of the atoms are delocalised
  • This means that they are free to move through the whole structure
  • The positive metal ions are then attracted to these delocalised electrons by the electrostatic force of attraction
• Pure metals are malleable because the layers can slide over each other
• Metals are good conductors of electricity and heat because delocalised electrons are free to move through the whole structure
• Metals have high melting and boiling points because the electrostatic forces of attraction between metal ions and delocalised electrons are strong and so lots of energy is needed to break them
• Alloys 


  • Pure metals are often too soft to use as they are
  • Adding atoms of a different element can make the resulting mixture harder because the new atoms will be a different size to the pure metal's atoms
  • This disturbs the regular arrangement of the layers, preventing them from sliding over each other
  • This harder mixture is called an alloy