How to master ionic bonding

What do we want our students to say when we ask them to describe ionic bonding? Ideally, they would talk about the electrostatic attraction between oppositely charged ions within a giant lattice of alternating positive and negative ions. However, the way we’ve traditionally taught ionic bonding in pre-16 courses sometimes results in confusion. Students often discuss the electron transfer process that creates ions from the reaction between metal and non-metal atoms. This misconception can hinder their progress in more advanced courses. So, how can we help students better understand and distinguish these ideas?

Common misconceptions

Dot and cross diagrams are regularly used to explain ionic bonding. These show the transfer of the outer electron(s) from the metal atoms to fill the outer shell of the non-metal atoms. However, this transfer of electrons is what happens during the redox reaction between a metal and a non-metal, and is not a type of bonding.

The emphasis on dot and cross diagrams also leads to a reliance on the octet rule to explain why bonding occurs. This rule can lead to the idea that atoms are always trying to get a stable full outer shell of electrons. The full outer shell rule can be useful, but can also hinder students developing more sophisticated understanding if they see the rule as the driving force for reaction.

What you need to know

Students who go on to post-16 courses will study the energy transfers involved in forming ions. They will discover that the first ionisation enthalpy of a gaseous metal atom, such as sodium, is very endothermic. Atoms certainly do not ‘want’ to lose their outer electrons. An isolated metal atom is therefore far more stable with its outer electrons than as a positive ion. However, it is common for students to cling on to the idea that metal atoms lose their outer electrons because they are trying to get a full outer shell.

What about adding an electron to an isolated non-metal atom such as chlorine? This process, called the first electron affinity, is exothermic. The magnitude of this enthalpy change doesn’t make up for the highly endothermic removal of the electrons from the metal though. Also, transfer of a second electron to an outer shell is an endothermic process, due to the repulsion between the second electron and the already negative ion.

Whichever way you look at it, the electron transfer process shown by the dot and cross diagrams at pre-16 is not energetically feasible. Explaining that ionic bonding happens so that atoms get stable full outer shells just doesn’t fit the energetics. This begs the question, why do these ions form at all?

At post-16, students go on to look at Born–Haber enthalpy cycles, which describe the formation of ionic compounds from elements. Born–Haber cycles bring together all the energy-transfer processes involved. They show that to get to the isolated atoms depicted in the dot and cross diagrams, you need to break the metallic bonding in the solid metal and break the covalent bonding in the diatomic non-metal molecules. When this is added to the endothermic electron-transfer process, it means that a large amount of energy must be transferred to the reacting elements to form the gaseous ions.