Definition of first ionization energy
The first ionization energy (often written as \(IE_1\)) is the energy required to remove the most weakly held electron from a gaseous neutral atom:
A larger \(IE_1\) means the valence electron is held more strongly by the nucleus.
Periodic table trend: across a period and down a group
Ionization energy on the periodic table is governed mainly by the competition between nuclear attraction and shielding by inner electrons (often summarized by effective nuclear charge \(Z_{\mathrm{eff}}\)).
| Direction on periodic table | Typical change in \(IE_1\) | Primary reasons |
|---|---|---|
| Left → right across a period | Increases | \(Z\) increases while shielding changes modestly within the same principal shell, so \(Z_{\mathrm{eff}}\) rises; atomic radius contracts and valence electrons are held more tightly. |
| Top → bottom down a group | Decreases | Valence electrons occupy higher principal energy levels (larger \(n\)), farther from the nucleus; increased shielding reduces attraction, making removal easier. |
Important exceptions (why “generally” is necessary)
Two recurring exceptions appear in many periods: (1) starting a new p subshell, and (2) pairing electrons within a p subshell.
| Exception pattern | Typical comparison | Chemical explanation |
|---|---|---|
| New subshell begins (s → p) | Group 2 vs Group 13 (e.g., Mg vs Al) | A p electron is, on average, higher in energy and less penetrating than an s electron in the same shell, so it is easier to remove; \(IE_1\) can dip at the first p element. |
| Electron pairing in p orbitals | Group 15 vs Group 16 (e.g., P vs S) | When one p orbital becomes paired, electron–electron repulsion increases; removing one paired electron can require slightly less energy, creating a small dip. |
Apply the trend: ordering first ionization energies in Period 3
Consider Period 3 elements: Na, Mg, Al, Si, P, S, Cl, Ar. The overall expectation is an increase from Na → Ar, with two dips at Al (start of p subshell) and S (paired p electron).
Correct lowest-to-highest order: \[ IE_1(\mathrm{Na}) < IE_1(\mathrm{Al}) < IE_1(\mathrm{Mg}) < IE_1(\mathrm{Si}) < IE_1(\mathrm{S}) < IE_1(\mathrm{P}) < IE_1(\mathrm{Cl}) < IE_1(\mathrm{Ar}) \]
Justification using electron configurations (valence level \(n=3\))
The valence-shell patterns explain both the general rise (increasing \(Z_{\mathrm{eff}}\)) and the two anomalies.
| Element | Valence configuration | Ionization-energy reasoning |
|---|---|---|
| Na | 3s1 | Single 3s electron is relatively easy to remove → very low \(IE_1\). |
| Mg | 3s2 | Higher \(Z_{\mathrm{eff}}\) than Na and a filled 3s subshell → higher \(IE_1\). |
| Al | 3s2 3p1 | First p electron is less penetrating and easier to remove than Mg’s 3s electron → dip: \(IE_1(\mathrm{Al}) < IE_1(\mathrm{Mg})\). |
| Si | 3s2 3p2 | Increasing \(Z_{\mathrm{eff}}\) across the period strengthens attraction → \(IE_1\) rises from Al. |
| P | 3s2 3p3 | Half-filled p subshell is relatively stable → higher \(IE_1\) than Si. |
| S | 3s2 3p4 | One p orbital is paired, increasing repulsion → slight dip: \(IE_1(\mathrm{S}) < IE_1(\mathrm{P})\). |
| Cl | 3s2 3p5 | Strong \(Z_{\mathrm{eff}}\) and small radius → \(IE_1\) rises again. |
| Ar | 3s2 3p6 | Closed-shell noble gas configuration → highest \(IE_1\) in the period. |
Visualization: qualitative first ionization energy trend across Period 3
Final takeaway for ionization energy on the periodic table
Ionization energy periodic table trends follow increasing \(Z_{\mathrm{eff}}\) across periods and increasing distance and shielding down groups. Across Period 3, these principles yield the order \( \mathrm{Na} < \mathrm{Al} < \mathrm{Mg} < \mathrm{Si} < \mathrm{S} < \mathrm{P} < \mathrm{Cl} < \mathrm{Ar} \), with Al and S as the key exceptions.