Periodicity shows a fairly smooth increasing trend across a period for ionisation energy. However, between groups 2 & 3 and groups 5 & 6, the trend doesn't appear to be followed. Using your knowledge of chemistry, explain why the trend isn't followed here
For elements in group 2, their valence electrons consist of "ns2" where n is the principle quantum number/shell number. However, for elements in group 3, their valence electrons consist of "ns2np1". As the p-orbitals are higher in energy than the s-orbitals, this makes the p electron easier to remove and hence required less energy input - resulting a slightly lower 1st ionisation energy than expected. For elements in group 5, they all have singly filled p-orbitals ("np3), but those in group 6 have one p-orbital with a pair of electrons occupying it ("np4"). This is known as an orbital pair with anti-spin. The pairing of electrons within an orbital causes repulsion between those occupying it - this highers the electrons energy. As their energy is now higher, it is easier to remove the electron as a result, which lowers the 1st ionisation energy.
