We are all familiar with the classic picture of an ionic lattice – for example, for a compound such as sodium chloride where sodium and chlorine exist as discrete ions in a regular 3-D arrangement.
X-ray diffraction studies show that there is virtually no electron density between the nuclei. This is consistent with the model of ionic bonding in which an electron has been transferred completely from one atom to another forming oppositely charged ions.
If we bring two oppositely charged ions together they will attract. The strength of this electrostatic attraction is strongest the closer they are to each other – this is seen as a lowering of energy / exothermic release of energy.
However, we know that the ions are not going to get closer and closer until they collapse on top of each other! There must be a force of repulsion … this repulsion is the result of the electrons of the oppositely charged ions repelling each other, and is seen as a raising of energy / endothermic term.
If we put these two graphs together there is a balance between the repulsion and the attraction such that a minimum energy exists (an equilibrium) which defines the actual optimum distance between two oppositely charged ions.
In a lattice we need to consider that we don’t have just one ion pair but millions of ions interacting. On balance, the attractive interactions outweigh the repulsions between ions of like charge and so forming a lattice from gaseous ions is highly exothermic.
Na+(g) + Cl–(g) ⇾ NaCl(s) ΔrH⦵ = -790 kJ mol-1
Limitations of the ionic lattice model
We can calculate a theoretical value for the energy of a lattice and we can also determine an experimental value. Sometimes the two are very similar, sometimes they are decidedly different!
The ionic lattice model we have described assumes that the electron density between ionic nuclei falls to zero, but this is not the case – it falls to a very low value but there is always some electron density between the ions.
The model also assumes that the radius of an ion is always the same, regardless of the compound, and this is not true. The radius of an ion depends to a small degree on the way the lattice is packed (arranged) and it is also affected by the nature of the other ions present.
NaCl and KCl are examples of ionic compounds where the theoretical and experimental lattice energies are very similar. In AgI and TlBr the experimental lattice energy is far greater than the theoretical – the lattice is stronger than would be expected if all the interactions were electrostatic, the attractions are stronger and this is evidence that there is significant covalent character in the Ag-I bond and the Tl-Br bond.
How can we explain this?
Large anions such as I– are easily distorted or polarised as the outer electron density of the ion is further from the nucleus and less strongly held. Small, highly charged cations such as Be2+ and Al3+ are strongly polarising (they have a high charge density). Compounds between the two will have a strong covalent character – aluminium iodide is a good example.
Another way to think about it is in terms of electronegativity (the ability of an atom to attract a shared pair of electrons):
- The greater the difference in electronegativity between two elements, the greater the ionic character of the bond
- If both elements have a high electronegativity the compound tends to be covalent
- If both elements have a low electronegativity then the substance will be an alloy with metallic bonding
This can be summarised in a bond-type diagram, which plots the average electronegativity of the two elements against the difference in electronegativity between them.
The point is that we shouldn’t think of ionic, metallic and covalent bonds as exclusive categories but rather in terms of a continuum.