Plots of first ionization energy across a period do not rise perfectly smoothly.

First ionization energy plotted against atomic number across many elements, showing the repeating sawtooth pattern across periods. The local dips within each period are the key visual evidence that the overall increase is interrupted by electronic-structure effects (e.g., entry into a higher-energy $p$ subshell and electron pairing repulsion). Source

These small irregularities are important because they reveal the existence of subshells and show changes in electron arrangement between successive elements.

Why first ionization energy data is useful

When comparing successive elements in the same period, there is an overall increase in first ionization energy because nuclear charge increases. However, the increase is not perfectly regular. Some elements have a slightly lower first ionization energy than the element before them.

These small drops are valuable evidence because they cannot be explained by proton number alone. Instead, they show that:

• electrons occupy different subshells

• subshells have different energies

• electron arrangement within a subshell can change the ease of electron removal

This means first ionization energy data supports a more detailed picture of atomic structure than a simple model with only shells.

Evidence for subshells in Period 2

The drop from beryllium to boron

The first clear evidence for subshells appears between beryllium and boron.

• Be: $1s^2\ 2s^2$

• B: $1s^2\ 2s^2\ 2p^1$

From Be to B, the extra electron goes into the 2p subshell, not the 2s subshell. Although boron has a greater nuclear charge, its first ionization energy is lower than beryllium’s.

This shows that the electron removed from boron is in a different type of subshell. A 2p electron is easier to remove than a 2s electron because:

• the 2p subshell is at higher energy

• the electron is slightly less strongly attracted to the nucleus

• it experiences slightly more shielding than a 2s electron in the same shell

So, the drop from Be to B is evidence that the second shell is split into at least two subshells with different energies.

The drop from nitrogen to oxygen

Another important pattern appears between nitrogen and oxygen.

• N: $1s^2\ 2s^2\ 2p^3$

• O: $1s^2\ 2s^2\ 2p^4$

Nitrogen has three electrons in the 2p subshell. These occupy separate p orbitals before pairing. Oxygen has one more 2p electron, so one orbital now contains a pair of electrons.

Orbital (box) diagrams for second-period $2p$ filling illustrate Hund’s rule and where electron pairing begins. Comparing nitrogen ($2p^3$) with oxygen ($2p^4$) shows that oxygen must contain one paired set in a single $p$ orbital, increasing repulsion and lowering the first ionization energy relative to nitrogen. Source

The first ionization energy of oxygen is lower than that of nitrogen, even though oxygen has a higher nuclear charge. This happens because:

• paired electrons in the same orbital repel each other

• this repulsion makes one of the paired electrons easier to remove

This drop does not provide evidence for a new subshell. Instead, it provides evidence for a change in electronic structure within the same subshell.

Evidence for subshells in Period 3

The same pattern appears again in Period 3.

First ionisation energy across Period 3 (Na→Ar) with the characteristic dips at Mg→Al and P→S highlighted by the shape of the curve. The Mg→Al dip corresponds to removal of a higher-energy $3p$ electron rather than a $3s$ electron, while the P→S dip reflects increased electron–electron repulsion when pairing starts in the $3p$ subshell. Source

Repetition across periods strengthens the evidence.

The drop from magnesium to aluminum

• Mg: $[Ne]\ 3s^2$

• Al: $[Ne]\ 3s^2\ 3p^1$

Aluminum has a lower first ionization energy than magnesium because the electron removed from aluminum is in the 3p subshell, which is higher in energy than the 3s subshell.

This is strong evidence that the third shell also contains separate subshells with different energies.

The drop from phosphorus to sulfur

• P: $[Ne]\ 3s^2\ 3p^3$

• S: $[Ne]\ 3s^2\ 3p^4$

Phosphorus has three unpaired 3p electrons. Sulfur has one paired 3p orbital. The repulsion between paired electrons makes sulfur’s outer electron easier to remove, so sulfur has a lower first ionization energy than phosphorus.

Again, this supports the idea that first ionization energy depends not only on shell and subshell, but also on the arrangement of electrons within orbitals.

What the data shows about electronic structure

From these patterns, chemists can infer several key ideas:

• s and p subshells exist within the same main shell

• p subshells are higher in energy than s subshells in the same shell

• electrons do not all have exactly the same energy within one shell

• pairing of electrons can reduce first ionization energy because of repulsion

The repeated drops from:

• Be to B

• N to O

• Mg to Al

• P to S

are especially important. They show that first ionization energy data gives evidence for both subshell structure and changes in electron arrangement as atomic number increases.

How to interpret first ionization energy graphs

When using first ionization energy data, do not focus only on the overall rise across a period. The most useful evidence often comes from the exceptions.

A strong explanation should:

• identify the pair of successive elements being compared

• state the relevant electron configurations

• identify whether the electron removed is from an s or p subshell

• explain whether the drop is caused by higher subshell energy or electron repulsion in a pair

A weak explanation usually says only that the trend “goes down” without linking it to electronic structure. In this topic, the value of the data is that it reveals where the electronic structure changes and what kind of change has occurred.