Bonding and physical properties

• Solids:

  • Have a regular arrangement of particles as they have definite shapes, while crystals have straight edges
  • Atoms are close together as they are not easily compressed
  • They vibrate in place as diffusion is slow and they expand when heated
• Liquids:

  • Random arrangement of particles as they move to fill the bottom of a container
  • Atoms are close together as they are not easily compressed
  • Move 'rapidly jostling' as liquid evaporate and diffusion is slow
• Gas:

  • Random arrangement of particles as they will move to fill a container
  • Are spaced far apart and are easily compressed
  • Rapidly move and diffusion is rapid. Gasses exert pressure
• Heating a solid


  The particles gain more energy and vibrate around a fixed position, making the solid expand
• Enthalpy is the heat energy change measured under constant pressure

  Temperature depends on the average kinetic energy of the particles
• Crystals are solids that have a regular repeating pattern of atoms and are held together by forces of attraction, either stronger ionic bonding or weaker intermolecular bonds. The strength affects its physical properties.

  e.g. stronger forces cause a higher boiling point
• There are 4 basic crystal types:

  • Ionic
  • Metallic
  • Molecular
  • Macromolecular
• Ionic crystals


  Strong electrostatic attractions between oppositely charged ions, can cause a high boiling point as a lot of energy is needed to break these bonds
• Metallic crystals


  Exist as a lattice of positive ions embedded in a sea of delocalised electrons, strong metallic bonds cause a high boiling point
• Molecular crystals


  Molecules held in a regular formation by intermolecular forces, covalent bonds within the molecules hold the atoms together, no covalent bonding between molecules, have a low boiling point as it is held together by intermolecular forces, soft and break easily, don't conduct electricity
• Macromolecular molecules 


  Covalent bonding extends throughout the crystal, high boiling point

  e.g. Diamond, graphite
• Diamond

  Carbon only, polymorph / allotrope of carbon

  Covalent bonding throughout with each carbon covalently bonded

  Each carbon forms 4 covalent bonds (tetrahedron 109.5)

  Very hard material

  Very high melting temperature

  Not a good conductor
• Graphite


  Strong covalent bonding and weaker Van der Waals

  Three single covalent bonds for each carbon (flat trigonal arrangement / trigonal planar, 120)

  2D arrangement of linked hexagons

  Delocalised electrons around the layer

  Can conduct electricity

  No covalent bonds between the layers and held with Van der Waals

  Soft and flakey due to weak intermolecular forces

  High melting temperature
• Electrical conductivity can let us know what kind of bonding a structure has