🟣 This is Higher Level (HL) content.
Formal Charge Formula
\[\text{FC} = V - L - \tfrac{B}{2}\]
V = valence electrons | L = lone pair electrons | B = bonding electrons
Shortcut: FC = (valence e⁻ from periodic table) − (dots) − (sticks)
Rules for Choosing the Best Structure
- Minimise magnitude – formal charges should be as close to zero as possible
- Charge conservation – sum of all FCs must equal the overall charge
- Electronegativity – any negative FC should be on the most electronegative atom
- No adjacent like charges – avoid placing same-sign FCs on neighbouring atoms
Worked Example – Phosphate Ion (PO₄³⁻)
Structure A: Octet
P single-bonded to 4 O atoms
FC on P: 5 − 0 − 4 = +1 ❌
FC on each O: 6 − 6 − 1 = −1
Total: +1 + 4(−1) = −3 ✓ but high charge separation
Structure B: Expanded Octet ✓
P double-bonded to 1 O, single to 3 O
FC on P: 5 − 0 − 5 = 0 ✅
FC on =O: 6 − 4 − 2 = 0
FC on –O: 6 − 6 − 1 = −1
Total: 0 + 0 + 3(−1) = −3 ✓ minimal charge separation
⚠️ Examiner Tip
A common arithmetic error: miscounting bonding vs lone pair electrons. Always draw the full Lewis structure first, then apply FC = V − dots − sticks to each atom systematically.