Why Does First Ionization Energy Generally Increase Across a Period? (HL)

First ionization energy generally increases across a period because nuclear charge increases while shielding remains approximately constant. The outer electron is therefore attracted more strongly to the nucleus, so more energy is required to remove it.

First ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms:

$$\text{X}(\text{g}) \rightarrow \text{X}^{+}(\text{g}) + \text{e}^{-}$$

Across a period, each successive element has one more proton. This increases the positive nuclear charge. The additional electrons enter the same principal energy level, so they do not introduce a substantial new inner shell of electrons.

Consequently, electron shielding by inner shells remains approximately constant. The outer electrons experience a greater effective nuclear charge, often represented qualitatively as $Z_{\mathrm{eff}} = Z - S$, where $Z$ is nuclear charge and $S$ represents shielding.

The stronger nucleus-electron attraction pulls the outer energy level closer to the nucleus, decreasing atomic radius. More energy must then be supplied to overcome this attraction and remove an electron.

The increase is not perfectly smooth because electron configuration affects the electron being removed:

A common misconception is that shielding increases greatly because more electrons are added. Across a period, the added electrons occupy the same principal energy level, so the increase in nuclear charge is the dominant effect.

Exam technique: For an HL explain question, link increased nuclear charge, approximately constant shielding, increased effective nuclear charge, reduced atomic radius, and stronger attraction. If given a graph, use electron configurations to explain the group 2-to-13 and group 15-to-16 discontinuities.