A GIANT IONIC LATTICE – explaining its properties

• All ionic compounds have a very similar structure and therefore very similar properties.
• The diagram on the right is typical of the giant ionic crystal structure of ionic compounds like sodium chloride and magnesium oxide.
• Solid ionic compounds consist of a giant lattice of closely packed ions which are all combine together to form a crystal. You can see in the right–hand diagram of sodium chloride, there is one positive ion to one negative ion, giving the empirical formula NaCl.
• The (+) and (-) ions are held together by electrostatic attraction.
• Same for KCl, MgO, CaO
• The ball and stick model is shown in the diagram on the right. Note that the thin lines are NOT bonds, they just indicate the geometry of the crystal structure. The electrostatic attractive force acts in ALL directions and that's what the ionic bond is.
• This type of diagram does not show any electronic detail of the ions, nor does it indicate the relative size of the ions (governed by the space occupied by the electron clouds), what it does show clearly is where the centres of the ions are positioned in the crystal lattice - the 3D spatial arrangement of the ions with respect to each other.
• The alternate positive and negative ions in an ionic solid are arranged in an orderly or regular way in a giant ionic lattice structure eg shown on the right.
• The ionic bond is the strong electrical attraction between the oppositely charged positive and negative ions next to each other in the lattice, so the electrostatic force of attraction in ionic compounds acts in all directions.
• The ionic bonding extends throughout the crystal in all directions.
• Salts and metal oxides are typical ionic compounds.
• This strong bonding force between the oppositely charged ions makes the structure hard (if brittle) and have high melting and very high boiling points, so they are not very volatile!
• A relatively large amount of energy is needed to melt or boil ionic compounds to reduce/overcome the strong bonding forces.
  • The more energy needed, the higher the melting point and boiling point, so most ionic compounds only melt and boil at relatively high temperatures – a direct consequence of the strong chemical bonding in ionic compounds.
  • Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.
• The bigger the charges on the ions the stronger the bonding attraction e.g. magnesium oxide Mg²⁺O^2– has a much higher melting point than sodium chloride Na⁺Cl^–.
  • The ions of magnesium oxide are both doubly charged so the electrostatic attraction is much greater (its actually about 4x as strong attractive force).
    • As it happens in this case, the ions in magnesium oxide are smaller than the ions in sodium chloride, so the ions in magnesium oxide can pack closer together and this also increase the attractive bonding force.
  • This double effect results in a much stronger ionic bond in magnesium oxide, so a much greater thermal kinetic energy i.e. a much greater temperature, is required to weaken the giant ionic lattice and melt the crystals of magnesium oxide compared to sodium chloride.
  • Simple experimental evidence – sodium chloride melts at 801^oC, whereas magnesium oxide melts much higher at 2852^oC.
• Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.
• They are hard but brittle, when stressed the bonds are broken along planes of ions which shear away.
  • They are NOT malleable like metals.
• Many ionic compounds are soluble in water, but not all, so don't make this assumption.
  • Salts can dissolve in water because the ions can separate and become surrounded by water molecules which weakly bond to the ions (see diagrams below).
  • This reduces the attractive forces between the ions, preventing the crystal structure to exist.
  • Evaporating the water from a salt solution will eventually allow the ionic crystal lattice to reform.
• The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current.
  • However, if the ionic compound is melted or dissolved in water, the liquid or solution will now conduct electricity.
  • This is because the ion particles are now free to move and carry the electric current in the molten salt or the solution of the salt in aqueous solution (see diagrams below).
    • An electric current is the flow of charged particles (ions or electrons).
    • This electrical conduction under these conditions is evidence for the existence of ions in this type of compound.

Need diagram to show conduction  - simple circuit diagram battery bulb electrodes solution

An 'advanced level' particle picture of sodium chloride dissolving in water

BUT, in reality there wouldn't be as much space between H₂O molecules ...

(the partial electrical charges δ+ and δ– are for advanced A level students only)

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solid sodium chloride ==> molten sodium chloride (from fixed ions to free moving ions)