Electron Affinity Explained Simply

Electron affinity therefore reveals how much an atom “wants” an extra electron.

Factors Affecting Electron Affinity

Electron affinity depends on three main ideas:

1. Nuclear Charge

More protons → stronger attraction for incoming electrons → more exothermic electron affinity.

This is why halogens have large electron affinities.

2. Atomic Radius

Large atoms → electron farther from nucleus → weaker attraction → less exothermic electron affinity.

Smaller atoms attract electrons more strongly.

3. Electron–Electron Repulsion

If the added electron must enter an already crowded subshell, repulsion increases and electron affinity becomes less favorable.

Example:
Adding an electron to oxygen’s 2p subshell introduces repulsion.

Periodic Trends in Electron Affinity

Across a Period (left to right): becomes more exothermic

Why?

• Nuclear charge increases
• Atomic radius decreases
• Attraction for added electron increases

Result: Elements toward the right side (especially non-metals) gain electrons more readily.

Down a Group: becomes less exothermic

Why?

• Atomic radius increases
• Shielding increases
• Added electron is farther from the nucleus

Halogens still strongly attract electrons, but the effect weakens down the group.

Special Case: Halogens

Halogens (Group 17) have:

• High nuclear charge
• One electron away from noble-gas configuration
• Small atomic radii (top of group)

Thus, they have the most exothermic electron affinities in the periodic table.

Example:
Cl has a more exothermic electron affinity than F due to reduced electron–electron repulsion in the larger 3p orbital.

This is a classic IB exception.

First vs Second Electron Affinity

First Electron Affinity

Usually exothermic.
X(g) + e⁻ → X⁻(g)

Second Electron Affinity

Always endothermic because you are adding an electron to a negative ion.
X⁻(g) + e⁻ → X²⁻(g)

Example:
O(g) + e⁻ → O⁻(g) (exothermic)
O⁻(g) + e⁻ → O²⁻(g) (endothermic)

The strong repulsion explains why second electron affinities are positive.

Why Electron Affinity Matters in IB Chemistry

Electron affinity helps explain:

• Why halogens are so reactive
• Why noble gases rarely form ions
• Why metals tend to lose electrons rather than gain them
• Trends in non-metal reactivity
• Formation of anions in ionic bonding

It also reinforces atomic structure concepts like nuclear charge and shielding.

Common IB Misunderstandings

“Electron affinity is the same as electronegativity.”

No—electronegativity is attraction in a bond, while electron affinity is about adding an electron to a gaseous atom.

“Electron affinity always releases energy.”

First electron affinity usually does, second electron affinity does not.

“Smaller atoms always have higher electron affinity.”

Usually—but oxygen is an exception because of repulsion in the 2p subshell.

“All non-metals have high electron affinity.”

Only some do—sulfur and phosphorus have moderate values.

FAQs

Why do noble gases have almost no electron affinity?

Their outer shells are full, so adding an electron is energetically unfavorable.

Why is chlorine’s electron affinity more exothermic than fluorine’s?

Fluorine is very small, causing strong repulsion in the crowded 2p orbital.

Which elements have the highest electron affinities?

Halogens, especially chlorine and bromine.

Conclusion

Electron affinity is the energy change that occurs when a gaseous atom gains an electron. It depends on nuclear charge, atomic radius, and electron repulsion. Across a period, electron affinity becomes more exothermic; down a group, it becomes less exothermic. Understanding electron affinity helps IB Chemistry students explain reactivity, periodic trends, and the formation of negative ions.

Learn why the mole allows chemists to connect atomic-scale particles to measurable quantities used in real experiments.