Understanding Ionisation Energy Basics

Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. Think of it as how tightly an atom holds onto its electrons - the higher the ionisation energy, the harder it is to remove an electron.

The process is always endothermic $energy-absorbing$, which makes perfect sense when you consider you're pulling negatively charged electrons away from positively charged nuclei. For oxygen, the first ionisation looks like: O₍ₘ₎ → O⁺₍ₘ₎ + e⁻

Once you've removed one electron, removing a second one (the second ionisation energy) becomes even harder because you're now pulling an electron from a positively charged ion. The pattern continues - each successive ionisation energy gets larger.

Quick Tip: Always remember that ionisation equations must show gaseous atoms and ions - that's why you'll see (g) subscripts everywhere!

Periodic Trends and Key Exceptions

Down a group, ionisation energy decreases because atoms get larger with more electron shells. Those outer electrons are further from the nucleus and better shielded by inner electrons, making them easier to remove.

Across a period, ionisation energy generally increases as nuclear charge grows stronger, pulling electrons closer. However, there are two crucial exceptions that examiners love to test.

The Group 2 to 3 exception happens because Group 3 elements have their outer electron in a p-orbital rather than an s-orbital. P-orbitals are higher energy and slightly further from the nucleus, making that electron easier to remove.

The Group 5 to 6 exception occurs because Group 6 elements have two electrons sharing a p-orbital. The repulsion between these paired electrons makes one easier to remove than from Group 5's singly-occupied orbitals.

Exam Alert: These exceptions are 6-mark question favourites - make sure you can explain both the orbital types and electron repulsion concepts clearly!