Ionization Energy Rises Across a Period, But Not Smoothly
Moving across a period, the nuclear charge increases as protons are added to the nucleus.
Electrons are added to the same principal energy level, so shielding stays roughly constant.
The stronger pull between the nucleus and the outer electrons raises the first ionization energy.
This steady rise is broken at two points in each period, and those dips are the clue to sublevels.
The Two Dips That Break the Trend
Group 2 to Group 13 (for example, Be to B)
Beryllium has a full 2s² sublevel, while boron's outer electron begins the 2p sublevel.
The 2p electron is higher in energy and is shielded by the 2s electrons, so less energy is needed to remove it.
Boron therefore has a lower first ionization energy than beryllium, against the general trend.
Group 15 to Group 16 (for example, N to O)
Nitrogen has a half-filled 2p³ sublevel, with one electron in each p orbital and none paired.
Oxygen's fourth 2p electron must pair up in an orbital that already holds an electron.
The repulsion between this pair makes one electron easier to remove, so oxygen's ionization energy dips below nitrogen's.
Tip
Explain a dip using a specific cause: the start of a new sublevel (group 2 to 13) or the first electron pairing in a p orbital (group 15 to 16).
Why the Dips Are Evidence for Sublevels
If electrons filled a single, featureless energy level, ionization energy would rise smoothly across the whole period.
Instead, the dips appear at the same points in every period, showing that electrons occupy distinct s and p sublevels with slightly different energies.
Across the d-block, first ionization energy changes only gradually, which fits electrons filling an inner 3d sublevel close in energy to the 4s.
Successive Ionization Energies: Evidence for Main Levels
Removing electrons one after another from the same atom shows a related but separate pattern.
A very large jump appears when the next electron must come from a lower main energy level, which is evidence for shells rather than sublevels.
Example
Successive ionization energies of magnesium (kJ mol⁻¹) show where each electron comes from:
1st: 738 (a 3s electron)
2nd: 1451 (the second 3s electron)
3rd: 7733 (a 2p electron)
The huge rise between the 2nd and 3rd values marks the move from the $n=3$ shell to the $n=2$ shell, evidence for main energy levels; the smaller dips across a period are what reveal sublevels within a shell.
Common Mistake
Ionization energy does not increase smoothly across a period.
Explain a discontinuity with the new sublevel or the first p-orbital pairing, not a vague appeal to stability.
Active recall
Why is the first ionization energy of boron lower than that of beryllium?
Why does oxygen have a lower first ionization energy than nitrogen?
What do the discontinuities across a period reveal about the arrangement of electrons?
How does the large jump between magnesium's 2nd and 3rd ionization energies provide evidence for main energy levels?
