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Explain why Aluminium and Sulfur do not fit the expected trends of first ionisation energies of period 3?

Firstly, we should define the first ionisation energy

The first Ionisation Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form a mole of gaseous ions, each with an 1+ charge.

The expected trend as we go across the period is that ionisation energy will increase as the nuclear charge increases and the number of shielding shells of electrons does not change. So the extra amount of protons means the nucleus holds the outer electrons more strongly so it requires more energy to remove an electron.

Aluminium has a lower ionisation energy than Magnesium. This is unexpected as Al has more protons. This can be explained by electron configurations.

Magnesium's electron config: 1s2s2 2p3s

Aluminium's electron config: 1s2s2 2p3s3p

Aluminium's outer electron is in a p orbital. This p sub-level is of higher energy than the s sub-level and so less energy is required to remove this electron.

Sulfur has a lower ionisation energy than phosphorous. This again is explained by their electron configurations

Phosphorous: 1s2s2 2p3s3p3

Sulphur: 1s2s2 2p3s3p4

A p sub-level has 3 sub-shells which can hold 2 electrons in each sub-shell. In phosphorous, the electrons are unpaired with one electron in each sub-shell with parallel spins. However, sulphur has 4 electrons so one of them must pair in a sub shell. As electrons are both negative particles, the paired electrons repel each other and so it is easier to remove the unpaired electron in phosphorous- so less energy is required.

Jake G. GCSE Chemistry tutor, A Level Chemistry tutor

9 months ago

Answered by Jake, a GCSE Chemistry tutor with MyTutor


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