Resonance Structures Formal Charge Calculator
Calculate formal charges on each atom of a Lewis structure to compare resonance structures.
Enter element, bonds, and lone pairs to get the formal charge.
Many molecules cannot be represented by a single Lewis structure. Carbonate ion CO₃²⁻ has three equivalent ways to draw the C=O double bond, with the other two oxygens carrying single bonds and the negative charges. Each is a resonance structure. The actual molecule is a hybrid — the bonding is averaged across all valid structures, with each C-O bond having a bond order of 4/3.
To pick which resonance structures contribute most to the hybrid, chemists compare formal charges. The structure with formal charges closest to zero, with negative charges on the most electronegative atoms, is the major contributor.
Formal charge formula:
FC = V − L − B
where V is the valence electron count for the atom (4 for C, 5 for N, 6 for O, etc.), L is the number of non-bonding (lone pair) electrons on the atom, and B is the number of bonds attached to it (count each double bond as 2, each triple bond as 3).
Worked example for carbonate CO₃²⁻ in the structure with one C=O and two C-O⁻:
The carbon has 0 lone pairs and 4 bonds (one double, two single). FC(C) = 4 − 0 − 4 = 0. The double-bonded oxygen has 2 lone pairs (4 electrons) and 2 bonds. FC = 6 − 4 − 2 = 0. Each single-bonded oxygen has 3 lone pairs (6 electrons) and 1 bond. FC = 6 − 6 − 1 = −1. Total = 0 + 0 + (−1) + (−1) = −2 ✓ (matches the ion’s charge)
Rules for evaluating resonance structures:
- The sum of all formal charges equals the molecule’s overall charge.
- Smaller absolute formal charges are better.
- Negative formal charges should sit on more electronegative atoms (F > O > N > C).
- Adjacent atoms with formal charges of the same sign are unfavorable.
- Resonance is most important when multiple equivalent (or near-equivalent) structures exist.
The major contributor in CO₃²⁻ has all three oxygen positions equivalent (each spends 1/3 of the time as the double-bonded one), explaining the equal C-O bond lengths observed experimentally.
Use this calculator one atom at a time. To check a complete structure, calculate formal charge for each atom and verify they sum to the overall charge of the molecule or ion.
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