$word_{count}$: 230

EA (electron gain enthalpy) = enthalpy change when gaseous atom gains an electron. More negative = more favourable.

General trends: EA becomes more negative across a period, less negative down a group.

Key exceptions:

1. EA(F) less negative than EA(Cl): F is too small — electron repulsion in compact 2p shell. Cl(-349) > F(-328) kJ/mol.
2. EA of N ≈ 0: half-filled 2$p^3$ resists addition. Similarly Be, Mg (filled $s^2$) have near-zero EA.
3. Noble gases: EA ≈ 0 (filled valence shell).
4. Second EA is always positive: adding electron to an already negative ion ($O^-$ → $O^{2-}$) costs energy (+780 kJ/mol).

Electronegativity (Pauling scale): tendency to attract shared electrons in a bond. Increases across period, decreases down group. F = 4.0 (highest), Cs = 0.7 (lowest among main group). EN determines bond polarity: delta EN > 1.7 → ionic; 0.4-1.7 → polar covalent; < 0.4 → non-polar.

EN depends on hybridisation: C(sp) > C(sp2) > C(sp3). Higher s-character → more electronegative. Noble gases have no Pauling EN (don't form bonds normally).