1.5 - structures

• Boiling temperature


  The temperature at which a substance changes from a liquid state to a gaseous state
• Delocalised electrons


  The electrons that are not contained within a single atom or a covalent bond
• Diamond


  • A type of giant covalent structure composed of carbon atoms
  • Each carbon atom is bonded to four other carbon atoms in a tetrahedral structure, making diamond very hard
• Giant covalent structure


  Large structures containing lots of atoms that are covalently bonded to each other, they are usually arranged in a regular lattice
• Giant covalent structure


  • Diamond
• Giant ionic lattice


  A regular repeating structure made up of oppositely charged ions
• Graphite


  • A type of giant covalent structure composed of carbon atoms
  • Each carbon atom is only bonded to three other carbon atoms in flat hexagonal sheets
  • There is one delocalised electron per carbon atom, so graphite can conduct electricity
  • The forces between the hexagonal layers in graphite are weak and can slide over each other
• Melting temperature


  The melting point of a substance is the temperature at which it changes from solid state to liquid state
• Metallic bonding


  Strong electrostatic attraction between positive metal ions and the sea of delocalised electrons that surround them
• Metallic structure


  Layers of positive metal ions surrounded by a 'sea' of delocalised electrons
• Simple molecular substances


  • The structures formed by covalent molecules with weak intermolecular forces between molecules
  • These substances generally have low melting and boiling points
• 1.6: The Periodic Table
• Ionic solids (crystals)


  Giant lattices of positive and negative ions
• Ionic solids


  • Structures are made of the same base unit repeated over and over again
  • The structure of the crystal depends on the relative number of ions and their sizes
• Cs+ ion is larger than Na+ ion


  More Cl– ions can fit around it
• Physical properties of ionic solids

  • High melting and boiling temperatures
  • Often soluble in water
  • Hard but brittle
  • Poor electrical conductivity when solid, but good when molten or dissolved
• High melting and boiling temperatures


  It takes a large amount of energy to overcome the strong electrostatic forces between the oppositely charged ions
• Soluble in water


  The oxygen end of the water molecules is attracted to the positive ions, and the hydrogen ends of the water molecules are attracted to the negative ions
• Hard but brittle


  When force is applied, layers of ions slide over each other causing ions of the same charge to be next to each other; the ions repel each other and the crystal shatters
• Poor electrical conductivity when solid, good when molten or dissolved


  In the solid state, the ions are fixed in position by the strong ionic bonds; however, when molten or dissolved, the ions are free to move and will move to the electrode of opposite charge, so will carry the current
• Giant covalent solids


  Networks of covalently bonded atoms arranged into giant lattices
• Diamond


  • Each carbon atom is joined to four others by strong covalent bonds
  • Atoms arrange themselves in a tetrahedral shape
  • Very hard
• Diamond


  • Very high melting temperature - a lot of energy needed to break the numerous strong covalent bonds
  • Does not conduct electricity - there are no free electrons or ions present
• Graphite


  • Consists of hexagonal layers
  • Each carbon is joined to three others by strong covalent bonds
  • The extra electrons are delocalised within the layer
  • The layers are held together by weak van der Waals forces
• Graphite


  • Very high melting temperature - it has strong covalent bonds in the hexagon layers
  • Soft and slippery - the weak forces between the layers are easily broken, so the layers can slide over each other
  • Good conductor of electricity - the delocalised electrons are free to move along the layers so an electric current can flow
• Metals


  • Metal atoms bond together to form a giant metallic structure
  • Metals consist of a regular arrangement of metal cations (a lattice) surrounded by a 'sea' of delocalised electrons
  • The strong metallic bond is due to the electrostatic forces of attraction between the nucleus of the cations and the delocalised electrons
• Physical properties of metals


  • High melting temperatures
  • Hard
  • Good conductors of electricity both in the solid and molten state
  • Good thermal conductors
  • Malleable and ductile
• High melting temperatures


  A large energy is needed to overcome the strong forces of attraction between the nuclei of the metal cations and the delocalised electrons; the melting temperature is affected by the number of delocalised electrons per cation and the size of the cation
• Good conductors of electricity

  The delocalised electrons can carry a current because, when a potential difference is applied across the ends of a metal, they will be attracted to and move towards the positive terminal of the cell
• Good thermal conductors


  The delocalised electrons can pass kinetic energy to each other
• Malleable and ductile


  When a force is applied to a metal, the layers of cations can slide over each other; however, the delocalised electrons move with the cations and prevent forces of repulsion forming between the layers
• Simple molecular solids


  Consist of covalently bonded molecules held together by weak intermolecular forces
• Physical properties of simple molecular solids


  • Low melting and boiling temperatures
  • Poor conductors of electricity
• Low melting and boiling temperatures


  Although the covalent bonds within the molecules are strong, the intermolecular forces holding the molecules together are weak and do not need much energy to break
• Simple molecular solids


  • Iodine
  • Ice
• Iodine


  • Atoms are covalently bonded in pairs to form diatomic I2 molecules
  • These molecules are held together by weak van der Waals forces and are arranged in a regular pattern
• Ice


  • Molecules of water are arranged in rings of six held together by hydrogen bonds
  • In this ordered structure, the water molecules are further apart than they are in the liquid state
  • Since there are large areas of open space inside the rings, ice is less dense than liquid water at 0˚C