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.