Predicting the Dominant Lewis Structure
This is Higher Level (HL) content.
Formal Charge (FC) Calculation
\[\text{FC} = V - L - \tfrac{B}{2}\]
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⁻)
Octet Rule Conforming Structure
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
Expanded Octet Structure (Preferred)
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
Electronegativity vs Negative Formal Charge
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.
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