First ionisation energy measures how tightly an atom holds its outermost electron. Nuclear charge, distance, shielding and one subtle repulsion between paired electrons explain every rise and both of the dips across a period.
First ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions: X(g) → X+(g) + e−. It is larger when the nuclear charge is greater, the outer electron is closer, and shielding is lower. Across Period 3 it generally rises, with dips at aluminium (3p sits above 3s) and sulfur (spin-pair repulsion in 3p4). Down a group it falls because the outer electron is further away and more shielded.
Fig. 1: first ionisation energies across Period 3 (kJ per mole)
Section 01
The definition and the three factors.
Every ionisation energy question is answered with the same three ideas. Learn them as a set and always name the ones that apply.
Nuclear charge: more protons, stronger attraction for the outer electron, higher ionisation energy.
Distance from the nucleus (atomic radius): attraction falls off quickly with distance, so an outer electron in a higher shell is easier to remove.
Shielding: inner-shell electrons repel the outer electron and cancel part of the nuclear pull. More inner shells, more shielding, lower ionisation energy.
Spin-pair repulsion: two electrons sharing one orbital repel each other, so one of a pair is slightly easier to remove than a single electron in an orbital of the same energy.
The equation must show gaseous atoms and gaseous ions: Mg(g) → Mg+(g) + e−. Second ionisation energy starts from the 1+ ion: Mg+(g) → Mg2+(g) + e−.
Section 02
Across Period 3.
Step along the period in the simulator. Predict the direction of each change before you reveal it, and stop at aluminium and sulfur to see the two exceptions.
General increase, Na to Ar: nuclear charge rises by one each time, the outer electron stays in n = 3 with similar shielding, so it is held more tightly.
Dip at Al (578 < 738 for Mg): aluminium's outer electron is in 3p, which is higher in energy than magnesium's 3s and partly shielded by the 3s electrons, so it is easier to remove despite the extra proton.
Dip at S (1000 < 1012 for P): phosphorus has three unpaired 3p electrons; sulfur's fourth 3p electron pairs up. Spin-pair repulsion makes one of the paired electrons easier to remove.
Period 2 repeats the pattern: dips at boron (2p above 2s) and oxygen (pairing in 2p4).
Stage 1 · Learn
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First ionisation energy is the energy required to:
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Which change, on its own, increases the first ionisation energy of an atom?
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Magnesium has a higher first ionisation energy than sodium because magnesium has:
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First ionisation energy decreases down Group 1 because the outer electron is:
Section 03
Down a group.
Group 1 runs Li 520, Na 496, K 419, Rb 403 kJ per mole. Group 2 runs Be 900, Mg 738, Ca 590, Sr 550, Ba 503.
Outer electron further from the nucleus: a new shell is added with each period.
More shielding: each extra inner shell cancels more of the nuclear charge.
Nuclear charge increases too, but distance and shielding dominate, so the ionisation energy falls.
Examiner trap
Three things lose marks in a definition: forgetting gaseous, forgetting one mole, and writing a 2+ ion for a first ionisation. In explanations, say shielding or outer electron further from the nucleus; "more electrons" on its own is not an explanation.
Stage 2 · Exam
Exam-style questions
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Why is the first ionisation energy of aluminium lower than that of magnesium?
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Why is the first ionisation energy of sulfur lower than that of phosphorus?
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Which equation represents the first ionisation energy of magnesium?
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Which element has the highest first ionisation energy?
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Ionisation energy trends, mastered.
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