Ionisation energy

Ionisation energy (ionization energy in the US) is the energy needed to remove electrons from atoms. The first ionisation energy is the energy it takes to turn one mole of gaseous atoms into one mole of gaseous 1+ ions, by taking a single electron from every atom. It generally increases across a period and decreases down a group, and the pattern of values is evidence for how electrons are arranged in shells and sub-shells.

First ionisation energy definition and equation

AQA's A-level specification asks you to define first ionisation energy and write equations for first and successive ionisation energies. Always show gaseous atoms and ions with state symbols:

Na(g) → Na⁺(g) + e⁻ first ionisation energy of sodium

The value is positive, because energy has to be put in to pull an electron away from the attraction of the nucleus. It is measured in kJ mol⁻¹.

What affects ionisation energy?

Three factors decide how strongly the outer electron is held:

• Nuclear charge: more protons attract the electron more strongly, raising the ionisation energy.
• Distance from the nucleus: an electron in a shell further out is held less strongly.
• Shielding: inner electrons repel the outer electron and reduce the attraction it feels from the nucleus.

Ionisation energy across period 3

Across period 3, the first ionisation energy generally rises. The nuclear charge increases by one each time, but the electrons being removed are in the same shell, with similar shielding, so they are held more strongly. AQA asks you to explain how first ionisation energies in period 3 (sodium to argon) give evidence for electron configuration in sub-shells. There are two dips:

The dip from magnesium to aluminium

Magnesium's outer electron is in the 3s sub-shell (1s² 2s² 2p⁶ 3s²); aluminium's is in 3p (… 3s² 3p¹). A 3p electron has slightly higher energy and is a little further from the nucleus, and it is partly shielded by the 3s electrons, so it is easier to remove even though aluminium has one more proton. OpenStax describes this as a small deviation from the trend each time a new sub-shell begins.

The dip from phosphorus to sulfur

Phosphorus has three 3p electrons, one in each 3p orbital. Sulfur has four, so two of them share an orbital. The paired electrons repel each other, which makes one of them easier to remove, so sulfur's first ionisation energy is lower than phosphorus's.

The same pattern appears in period 2, with dips at boron and oxygen.

Ionisation energy down group 2

Approximate first ionisation energies from OpenStax:

Going down the group, each new element adds a shell, so the electron being removed sits further out with more inner shells screening it. The extra protons don't make up for that, so the ionisation energy falls. This is part of why the group 2 metals become more reactive down the group: their outer electrons are lost more easily.

Successive ionisation energies

The second ionisation energy removes an electron from a 1+ ion, the third from a 2+ ion, and so on:

Ca⁺(g) → Ca²⁺(g) + e⁻ second ionisation energy of calcium

Each successive ionisation energy is larger than the one before, because the electron is being removed from an increasingly positive ion. But the increase isn't smooth. A big jump appears when the next electron has to come from a shell closer to the nucleus. OpenStax explains that this jump corresponds to the removal of core electrons, which are harder to remove than the valence electrons.

Example: identifying the group

Calcium's first three ionisation energies, from OpenStax, are 589.8, 1145.4 and 4912.4 kJ mol⁻¹. The second is about double the first, but the third is over four times the second (4912.4 ÷ 1145.4 = 4.3). The big jump comes after two electrons, so calcium has two outer electrons and is in group 2.

For potassium the jump comes straight after the first: 418.8 then 3051.8 kJ mol⁻¹, a factor of about 7.3, so potassium has one outer electron, in group 1.

Common mistakes

• Leaving out state symbols. The atoms and ions must be gaseous: Mg(g) → Mg⁺(g) + e⁻.
• Writing the second ionisation as removing two electrons. It removes one electron from the 1+ ion: Mg⁺(g) → Mg²⁺(g) + e⁻.
• Explaining the Mg to Al dip with pairing. That dip is about the new 3p sub-shell; pairing explains the P to S dip.
• Saying shielding increases across a period. Across a period the electrons go into the same shell, so shielding stays roughly the same while nuclear charge rises.
• Forgetting "one mole" in the definition.

Electron arrangements are built up in the Atomic Structure course, and the electron configuration guide covers sub-shells in detail. For practice in a short timed paper, try today's A-level Chemistry: Year 1 paper.

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