Types of Substance

1.31 Explain the formation of simple molecular, covalent substances, using dot and cross diagrams, including: hydrogen, hydrogen chloride, water, methane, oxygen, and carbon dioxide

1.32 Explain why elements and compounds can be classified as: ionic, simple molecular (covalent), giant covalent, metallic and how the structure and bonding of these types of substances results in different physical properties, including relative melting point and boiling point, relative solubility in water and ability to conduct electricity (as solids and in solution)
Properties of ionic compounds
  • Ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions.
  • They have high melting and boiling points, because a lot of energy is required to break the many strong bonds.
  • When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and carry current, and they do not conduct electricity as solids, because the ions are fixed and are not able to move, carrying charge with them.
    • Dissolve in water due to the polarity of the ionic bond – water is polar and like dissolves like
Properties of simple molecular compounds
  • Substances that consist of small molecules are usually gases or liquids that have low boiling and melting points.
  • Substances that consist of small molecules have weak intermolecular forces between the molecules. These are broken in boiling or melting, not the covalent bonds.
    • The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points.
  • Substances that consist of small molecules don’t conduct electricity, because small molecules do not have an overall electric charge.
Giant Covalent Structures
  • Substances that consist of giant covalent structures are solids with very high melting points.
    • All of the atoms in these structures are linked to other atoms by strong covalent bonds.
      • These bonds must be overcome to melt or boil these substances.
Properties of Metals
  • Metals consist of giant structures of atoms arranged in a regular pattern.
  • The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure.
  • The sharing of delocalised electrons gives rise to strong metallic bonds.

  • Metals have giant structures of atoms with strong metallic bonding.
    • Therefore, most metals have high melting and boiling points.
    • They can conduct heat and electricity because of the delocalised electrons in their structures.
    • Conduction depends on the ability for electrons to move throughout the metal.
    • The layers of atoms in metals are able to slide over each other, so metals can be bent and shaped.
1.33 Explain the properties of ionic compounds limited to: high melting points and boiling points, in terms of forces between ions and whether or not they conduct electricity as solids, when molten and in aqueous solution

*see 1.32

1.34 Explain the properties of typical covalent, simple molecular compounds limited to: low melting points and boiling points, in terms of forces between molecules (intermolecular forces) and poor conduction of electricity

*see 1.32

1.35 Recall that graphite and diamond are different forms of carbon and that they are examples of giant covalent substances
  • In diamond (right), each carbon is joined to 4 other carbons covalently.
    • It’s very hard, has a very high melting point and does not conduct electricity.

  • In graphite, each carbon is covalently bonded to 3 other carbons, forming layers of hexagonal rings, which have no covalent bonds between the layers.
    • The layers can slide over each other due to no covalent bonds between the layers, but weak intermolecular forces. Meaning that graphite is soft and slippery.
  • One electron from each carbon atom is delocalised.
    • This makes graphite similar to metals, because of its delocalised electrons.
    • It can conduct electricity – unlike Diamond.
1.36 Describe the structures of graphite and diamond

*see 1.35

1.37 Explain, in terms of structure and bonding, why graphite is used to make electrodes and as a lubricant, whereas diamond is used in cutting tools
  • Graphite uses
    • Electrodes – graphite can conduct electricity – unlike Diamond
    • Lubricant – weak intermolecular forces and no covalent bonds between the layers, therefore it is soft and slippery
  • Diamond uses
    • Cutting tools – very hard, due to its rigid structure 
1.38 Explain the properties of fullerenes including C60 and graphene in terms of their structures and bonding
  • Graphene
    • Single layer of graphite
    • Has properties that make it useful in electronics and composites
  • Carbon can also form fullerenes with different numbers of carbon atoms.
    • Molecules of carbon atoms with hollow shapes
    • They are based on hexagonal rings of carbon atoms, but they may also contain rings with five or seven carbon atoms
    • The first fullerene to be discovered was Buckminsterfullerene (C60), which has a spherical shape
  • Carbon nanotubes
    • Cylindrical fullerenes with very high length to diameter ratios
    • Their properties make them useful for nanotechnology, electronics and materials
  • Examples of uses
    • They can be used as lubricants, to deliver drugs in the body and catalysts.
    • Nanotubes can be used for reinforcing materials, for example tennis rackets.
1.39 Describe, using poly(ethene) as the example, that simple polymers consist of large molecules containing chains of carbon atoms
  • Polymers have very large molecules
  • Atoms in the polymer molecules are linked to other atoms by strong covalent bonds
  • Intermolecular forces between polymer molecules are relatively strong and so these substances are solids at room temperature

1.40 Explain the properties of metals, including malleability and the ability to conduct electricity

*see 1.32