Ionization energy increases across a period and decreases down a group.
First: what is ionization energy ? It’s the minimum amount of energy required to remove an electron (to infinity) from the atom or molecule in the gaseous state
These are the factors that affect ionization energy:
- Atomic radius
- Nuclear charge
- Orbital penetration
- Electron pairing within a sub-shell
- Shielding or screening effect of the inner orbitals
Atomic radius:
- When an atomic radius decreases, ionization energy increases.
- Across a period, atomic radius decreases. Down a group, atomic radius increases. Therefore, across a period ionization energy increases down a group it decreases.
Nuclear charge:
- If a nucleus has a positive charge, there is a stronger attraction for the electrons. This results in a high amount of ionization energy.
- Across a period, attraction of electrons from the nucleus increases, while down a group it decreases. Therefore, across a period ionization energy increases down a group it decreases.
Orbital penetration:
- Since the s orbitals penetrate towards the nucleus more closely than the p orbitals, it’s easier to remove electrons from p orbitals than from s orbitals.
- Therefore, across a period ionization energy increases down a group it decreases.
Electron pairing within a sub-shell:
- Since repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals, within a sub-shell, paired electrons are easier to remove than unpaired ones.
- Across a period ionization energy increases, down a group it decreases.
Shielding or screening effect of the inner orbitals:
- Electrons in the outermost orbitals feel lesser attraction from the nucleus than expected, due to presence of electrons in inner orbitals.
- Therefore, across a period ionization energy increases, down a group it decreases.
Overall: across a period ionization energy increases and down a group ionization energy decreases.
