Ionization energy (IE) is the energy required to remove the most loosely bound electron from a neutral atom in the gas phase.
Here’s how it changes in the periodic table:
1. Across a Period (→ left to right)
- Ionization energy increases.
- Reason:
- Nuclear charge (number of protons) increases.
- Electrons are added to the same shell (no extra shielding).
- Stronger attraction between nucleus and valence electrons makes them harder to remove.
Example: Sodium (Na) has a much lower IE than Chlorine (Cl) in Period 3.
2. Down a Group (↓ top to bottom)
- Ionization energy decreases.
- Reason:
- New electron shells are added as you move down.
- Valence electrons are farther from the nucleus and experience more shielding.
- Weaker attraction means less energy is required to remove an electron.
Example: Lithium (Li) has a higher IE than Cesium (Cs) in Group 1.
Summary of Ionization Energy Trend:
- Across a period → increases.
- Down a group → decreases.
Note: There are small exceptions (like between oxygen and nitrogen, or between boron and beryllium) due to electron repulsion in orbitals.