2.2.2 (a) Ionic Bonding and Dot-Cross Diagrams
Syllabus
(a) ionic bonding as electrostatic attraction between positive and negative ions, and the construction of 'dot-and-cross' diagrams.
What does this mean?
Your GCSE definition of Ionic Bonding is still find.
"Ionic Bonds are the electrostatic attractions between oppositely charged ions."
And your GCSE Dot-Cross diagrams are also fine.
The usual example
The most commonly asked for is still Sodium Chloride (NaCl)
It will probably only earn you 1 mark now.
You will undoubtedly be asked to show "only outside shells".
Adding extra shells is more work, looks worse and gains you nothing.
The diagram above would gain you the marks because:
i) it shows the (+) and (-) charges of the Cation and Anion.
ii) It is clear that one of the electrons on the Chloride ion has come from the Sodium atom.
But how do you know what to draw?
Sodium has Atomic Number 11 (so has electronic configuration 2.8.1)
Or it is in Group 1 so its electronic configuration must end with a 1
Either way it needs to lose 1 electron to have a full outer shell (or a complete octet).
Losing negative electrons must make you a positive Cation.
Losing one electron must make you a 1+ ion.
Chlorine has Atomic Number 17 (so has electronic configuration 2.8.7)
Or it is in Group 7 so its electronic configuration must end with a 7
Either way it needs to gain 1 electron to have a full outer shell (or a complete octet).
Gaining negative electrons must make you a negative Anion.
Gaining one electron must make you a 1- ion.
So, the diagram shows the Sodium atom losing the only electron in its outer shell, the Chlorine atom adding one electron to the seven it already has, and the charges of the two ions formed with Sodium having an empty outer shell and the Chloride having a full outer shell.
Other examples
Magnesium Oxide
Magnesium has Atomic Number 12 (electronic structure 2.8.2)
Or it is in Grp 2. So, its electronic configuration must end with 2
Either way it needs to lose 2 electrons to have a full outer shell (or a complete octet).
Losing two electrons must make you a 2+ ion.
Oxygen has Atomic Number 8 (electronic configuration 2.6)
Or it is in Grp 6. So, its electronic configuration must end with 6
Either way it needs to gain 2 electrons to have a full outer shell
Gaining two electrons must make you a 2- ion.
The first diagram shows all the shells (unnecessary).
More importantly, it shows the outside shells, the electron transfers and the resulting charges.
The second diagram would be equally acceptable at A level.
It simply suggests that 2 Cl- ions are formed in the same way.
And it is a slightly more obvious way of showing the formula of the resulting compound as MgCl2.
Sodium Oxide.
Sodium atoms still need to lose one electron each for the reason set out above.
But Oxygen atoms need to gain two electrons.
So we need two Sodium atoms for each Oxygen atom as below.
The electron transfers are obvious, the charges are correct, the outer electrons on the ions formed are also correct.
Aluminium Fluoride
Aluminium is in Group 3 - it needs to lose three electrons.
Fluorine is in in Group 7 - needs to gain only 1 electron.
So we'll need three Fluorine atoms for each Oxygen.
This diagram shows the outer shell of the 3+ Aluminium ion empty.
And it shows that we need 1- Chloride ions which contain seven electrons plus the one that they have gained.
Aluminium Oxide
Aluminium ions need to lose 3 electrons each. (2.8.3)
Oxygen atoms need to gain 2 electrons each. (2.6)
The common multiple of 2 and 3 is 6.
It takes two Aluminium atoms to lose 6 electrons.
And three Oxygen atoms to gain them.
The first diagram is okay, although an examiner might want to see dots and crosses rather than just dots.
The second diagram is fine as it shows the correct charges and the correct final outer shell and suggests that each Oxygen gained two electrons (crosses) to add to the six (dots) that it already had.
It also shows the number of Aluminium and Oxygen atoms involved.
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